JNC TN8400 99-011 JP9955327 42
Thermodynamic Data for the Speciation and Solubility of Pd, Pb, Sn, Sb, Nb and Bi in Aqueous Solution
January, 1999
33002936.
JAPAN NUCLEAR CYCLE DEVELOPMENT INSTITUTE
3 1-09 319-1194
Inquiries about copyright and reproduction should be addressed to : Techinical Information Section Administration Division Tokai Works Japan Nuclear Cycle Development Institute 4-33 Muramatsu, Naka-gun, Ibaraki 319-1194, Japan. JNC TN8400 99-011 January,1999
Thermodynamic Data for the Speciation and Solubility of Pd, Pb, Sn, Sb, Nb and Bi in Aqueous Solution
Barbara Lothenbach**, Michael Ochs**, Hans Wanner*** and Mikazu Yui*
Abstract
This report provides thermodynamic data for predicting concentrations of palladium Pd, lead Pb, tin Sn, antimony Sb, niobium Nb and bismuth Bi in geologic environments, and contributes to an integration of the JNC chemical thermodynamic database, JNC-TDB (previously PNC-TDB), for the performance analysis of geological isolation system of high-level radioactive wastes. Besides treating hydrolysis in detail, this report focuses on the formation of complexes or compounds with chloride, fluoride, carbonate, nitrate, sulfate and phosphate. Other important inorganic ligands (sulfide for lead and antimony, ammonia for palladium) are also included. In this study, the specific ion interaction theory (SIT) approach is used to extrapolate thermodynamic constants to zero ionic strength at 25 °C.
*: Waste Isolation Research Division, Tokai works, Japan Nuclear Cycle Development Institute (JNC) **: BMG ENGINEERING Ltd., Switzerland **: HSK, Switzerland JNC TN8400 99-011 1999 £fU£
Pd, Pb,Sn,Sb, Bi
Barbara Lothenbach**, Michael Ochs**, Hans Wanner***, M#HfP*
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Preface
The PNC-TDB project was initiated by PNC in 1996 with the aim to establish, by March 31, 1998, an internationally acknowledged equilibrium database for 20 elements that were defined as potentially important by PNC for the safety of a nuclear waste repository. I accepted to coordinate the project at an international level, and Dr. Dhanpat Rai (Battelle PNNL, Richland) as well as Dr. Gregory Choppin (Florida State University, Tallahassee) accepted to participate as experts. Their task was the setting-up of an equilibrium database on the actinide elements. The tedious job of administrative project management, including the contracting business with all participants was assumed by Dr. Michael Ochs (BMG Engineering, Schlieren, Zurich). In the 2nd year of this project, he and his colleague Dr. Barbara Lothenbach also were put in charge to develop the chemical thermodynamics of some non-actinide elements.
The kick-off meeting was held at Tokai-mura, May 27-30, 1996, in the presence of the Japanese experts: Dr. O. Tochiyama, Dr. H. Moriyama, Dr. S. Nakayama and Dr. T. Yamaguchi. Nevertheless, the various contracts were not in place until October 1996. A second meeting involving only PNC, BMG and myself was held at BMG, Zurich, November 18 - 20, a third meeting was held at Richland, Washington, May 27 - 29, 1997, without participation of the Japanese experts, and a fourth one on October 29, 1997, on the occasion of the Migration '97 conference at Sendai, again with the presence of the Japanese experts.
It was clear from the beginning, and stated explicitly at the kick-off meeting, that the timescale of the project (1.5 years) was extremely ambitious. Obviously, it was impossible to achieve, for the final product, a quality level that would be comparable to that of the NEA-TDB. We nevertheless decided to choose a procedure that resembles the one of NEA-TDB at least in the basic principles. This will allow immediate improvements of the pre-selected data in a possible follow-up project.
PNC decided from the start to accept different procedures for the establishment of the actinide database on one hand, and of the non-actinide databases on the other hand. For the actinide compounds and complexes of the +III and +IV oxidation states PNC preferred the establishment of a complete Pitzer model to the adoption of the NEA database. Less weight was put on the development of databases for penta- and hexavalent actinide species. The motivation behind this decision was PNC's conviction that only the +III and +IV oxidation states of the actinides will be relevant in their performance assessment analyses due to their choice of near- field components and geological sites.
For the non-actinide elements, PNC preferred initially to have its own staff perform a review of the available literature and establish selected data sets, while the members of the expert team should review these data sets and make constructive comments in such a way that the PNC staff could then implement the improvements and thus accomplish the various data sets according to international standards. The state-of-the-art analysis of the project during the second meeting in November 1996 revealed that this procedure was unlikely to result in satisfactory data sets by the final deadline of March 31, 1998. It was then decided to split the non-actinide elements up into two sets, one to be treated by BMG according to the guidelines of the NEA-TDB, the other to be treated by PNC staff according to their own procedures.
in JNC TN8400 99-011
The tasks that had to be accomplished by the PNC-TDB project team were the following:
D. Rai/PNNL Database on actinide compounds and complexes in the +III and +IV oxidation states G. Choppin/FSU Database on actinide compounds and complexes in the +V and +VI oxidation states, plus the actinide redox potentials M. Ochs/BMG Database on the elements Bi, Nb, Pb, Pd, Sb and Sn
As I have indicated above, the data sets presented in this report cannot be considered to correspond to the quality level of the NEA-TDB due to the reasons outlined below. However, I consider them as an excellent, partly high-quality basis that will require some quality-checking with independent experimental literature in the case of the actinide models, and a more detailed review of some of the reported literature in the other cases.
My personal assessment of the quality of the database presented in this report and the reports by D. Rai, G. Choppin and co-workers, and of the improvements and refinements that may be required to obtain ,,final", high-quality data sets, is briefly summarized below:
Actinide elements: The actinide +III model is based on the extensive experimental experience of D. Rai with the chemical behavior of lanthanides and actinides. The evidence for analogous treatment of the +III oxidation states is compelling and convincing. A complete Pitzer parameter set is provided. It is important that the Pitzer parameters be used in modeling studies, because in some cases (e.g., the sulfate complexes) complex formation is expressed by Pitzer parameters rather than equilibrium constants. This data set has, to my knowledge, not yet been checked against independent experimental studies from other laboratories. This task, which was not part of the present project but would improve the credibility of the data set, is strongly recommended for the near future.
For the +IV oxidation state of the actinides, the analogy concept is less evident than for the +III oxidation state. The recommendation to cross-check the data against independent experiments is also valid here.
The reason why the +V and +VI oxidation states of the actinides have been treated in a less sophisticated way is due to the low weight PNC has assigned to them. Without expressing any criticism of the selected values themselves, I feel that the review procedure here is not sufficiently transparent to cope with international standards. These data are based on expert judgment rather than on detailed analysis of the available experimental papers. The same comment applies to the redox potentials. However, due to the low weight PNC had assigned to this part of the project, the corresponding budget was quite limited, and a detailed analysis of these chemical systems could not be carried out.
Non-actinide elements (Bi, Nb, Pb, Pd, Sb and Sn): For these elements all the experimental data have been compiled from the available literature. Detailed reviews of the experimental papers have in general not been performed due to time and
IV JNC TN8400 99-011 budget constraints, an exception being the hydrolysis papers of the palladium review. The procedure chosen is consistent with the guidelines of the NEA-TDB: Experimental data are used to perform SIT plots (SIT = Specific Ion Interaction Theory, cf. NEA-TDB). In cases where the experimental data from different sources show satisfactory agreement on the SIT plot, an extrapolation to zero ionic strength may provide confidence in the resulting thermodynamic constants. It will nevertheless be necessary to review the experimental papers in detail in order to achieve international quality standards for the selected data sets. In addition, important gaps have been identified in many cases, and experimental programs should be performed to provide the information required for credible predictive modeling.
As a conclusion, I believe that the objectives of the PNC-TDB project, as defined at the kick-off meeting and re-defined at the second meeting in June 1997, have been reached. It is somewhat unfortunate that part of the database comes with complete Pitzer parameter sets, others with SIT parameter sets, and yet others with no ion interaction parameters at all. However, as I have mentioned in the beginning of this Preface, the procedure that unavoidably had to lead to this difference in activity factor treatment was consciously chosen by PNC. In my view, the database presented here provides a good basis for further refinement. It could therefore carry the attribute ,,Phase I". I strongly recommend to follow the above suggestions for independent verification on one hand, and for more detailed review on the other hand, in addition to the possible definition of experimental programs to fill the gaps that are considered critical for performance assessment purposes.
Wiirenlingen, May 18, 1998 Hans Wanner Scientific Project Coordinator JNC TN8400 99-011
Executive Summary
PNC is planning to submit a new Performance Assessment Report by March 2000. Besides informations on geology, repository design etc., basic data on radionuclide behavior are required for the safety analysis. These basic data include chemical thermodynamic data in aqueous media, requiring the establishment of a chemical thermodynamic database, PNC-TDB, for the interaction of key radionuclides with significant ligands under relevant conditions.
The subject of the present report are the key elements:
• tin • antimony • lead • bismuth • niobium • palladium
For these elements, element-specific datasets have been developed based exclusively on experimental studies published in the literature, rather than relying on existing compilations. This report focuses on the formation of complexes or compounds with:
• hydroxide • chloride • fluoride • carbonate • nitrate • sulfate • phosphate
Other important inorganic ligands (sulfide for lead and antimony, ammonia in the case of palladium) are also included. The number of experimental studies varies from one element to the other, and the data set selected here cannot in general be considered 'complete' for geochemical modeling applications. Where data for different ionic strengths are available, the specific ion interaction equation (SIT) is used to extrapolate formation constants to zero ionic strength. The selected formation constants are listed in Section 3 of this report.
Sections 2 and 3 of this report contain
• the data selection criteria • an outline of the extrapolation procedure to I = 0 with the specific ion interaction equation (SIT) • the selected data for the key elements Sn, Sb, Pb, Bi, Nb, Pd
In the following Sections 4-9, the data selection for the individual key elements is discussed in detail. For each element, tables containing the complete data compilation as well as all calculations used in the data evaluation for the individual elements, are made available to PNC.
VI JNC TN8400 99-011
Table of contents
1 Introduction 1
1.1 Background 1
1.2 Organization of the JNC-TDB project 1
2 Standards and Conventions 3
2.1 Symbols, units, notations and conversion factors 3 2.1.1 Symbols and notation 3 2.1.2 Compilation of thermodynamic data 4 2.1.3 Phase designators 5 2.1.4 Physical constants 5 2.1.5 Equilibrium constants 5 2.1.6 Redox reactions 7
2.2 Data selection criteria 7
2.3 Ionic strength corrections 7
2.4 Auxiliary Data 8 2.4.1 Selected thermodynamic data for auxiliary species 9 2.4.2 Conversion of AfG°values to equilibrium constants 9
3 Selected data : 12
3.1 Selected data for Sn, Sb, Pb, Bi, Nb andPd 12
3.2 Gaps and uncertainties 22 3.2.1 Uncertainties 22
3.2.2 Gaps 22
4 Tin 27
4.1 Hydrolysis of tin(TV) 27 4.1.1 Hydrolysis of tin(IV) under acidic conditions: Sn4+, SnOH3+, 2+ + Sn(OH)2 and Sn(OH)3 28 4.1.2 Hydrolysis of tin(IV) under neutral and alkaline conditions: 2 Sn(OH)5-and Sn(OH)6 " 29 4.2 Solid tin(IV) oxides/hydroxides 32
4.2.1 Freshly precipitated Sn(OH)4(am) 33 4.2.2 SnO2(precip) 33
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4.2.3 SnO2(cassiterite) 34 4.2.4 Additional data compiled for the formation of tin(IV) hydroxide/oxide compounds 35 4.3 Other tin(TV) complexes and compounds 36
4.4 Redox reactions 40 4.4.1 Sn2+/Sn(cr) 40 4.4.2 Sn2+/Sn4+ 41
4.4.3 Sn(OH)4° 42 4.4.4 Additional data compiled for the tin redox system 43 4.5 Hydrolysis of tin(II) 45 4.5.1 SnOH+ 47
4.5.2 Sn(OH)2° 48 4.5.3 Sn(OH)3" 49 2 4.5.4 Sn3(OH)4 + 50 2+ 4.5.5 Sn2(OH)2 50 4.5.6 MSn(OH)3+ 50 4.5.7 Additional equilibrium data compiled for Sn(II) hydrolysis 50 4.6 Solid tin(H) oxide/hydroxide 54 4.6.1 SnO(cr) and Sn(OH)2(precip) 55 4.6.2 Additional equilibrium data compiled for tin(II) hydroxide/oxide compounds 56 4.7 Tin(II) chloride system 58 + 2 4.7.1 SnCl , SnCl2°, SnCl3-, and SnCl4 - 60 4.7.2 SnOHCl0 : 62 4.7.3 SnCl2(s) 63 4.7.4 SnOHCl(s) 63 4.7.5 Additional equilibrium data compiled for the tin(II) chloride system 64 4.8 Tin(II) fluoride system 68 + 4.8.1 SnF , SnF2°, andSnF3- 69 4.8.2 SnF2(s) 71 4.8.3 Additional equilibrium data compiled for the tin(II) fluoride system 71 4.9 Tin(II) carbonate system 73
4.10 Tin(II) nitrate system 74 + 2 4.10.1 SnNO3 , Sn(NO3)2°, Sn(NO3)3-and Sn(NO3)4 - 75 4.10.2 Tin(II) nitrate compounds 77 4.10.3 Additional equilibrium data compiled for tin(II) nitrate system 77 4.11 Tin(II) phosphate system 79
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4.12 Tin(II) sulfate system 80 4.12.1 Tin(n) sulfate complexes 80
4.12.2 SnSO4(s) 81 4.13 Tin(II) sulfide system 82 4.13.1 SnS(herzenbergite) 82
4.13.2 Sn2S3(s) and Sn3S4(s) 82 4.14 Comments on selected references 83
5 Antimony 87
5.1 Hydrolysis of antimony (III) 87 5.1.1 Sb3+ 89 5.1.2 SbOH2+ 90 + 5.1.3 Sb(OH)2 90 5.1.4 Sb(OH)4- 92 2+ 5.1.5 Sb2(OH)4 93 5.1.6 Sb2(OH)6° 93 5.1.7 Additional equilibrium data compiled for the hydrolysis of antimony(m) 93 5.2 Solid antimony(m)-oxide/hydroxide 95
5.2.1 ct-Sb2O3 (valentinite) 96 5.2.2 £-Sb2O3 (senarmontite) 97 5.3 Antimony(in) chloride system 99 3+ 2+ 2 5.3.1 Reactions of Sb : SbCl ,.SbCl2+, SbCl3°, SbCLr, SbCl5 - 3 andSbCl6 - 100 5.3.2 SbCl4~ and SbOHCl3- 101 5.3.3 SbCl3(s) and SbOCl(s) (or Sb4O5Cl2(s)) 102 5.3.4 Additional equilibrium data compiled for the antimony (III) chloride system 102 5.4 Antimony(ITf) fluoride system 104 3+ 2+ 5.4.1 Reactions of Sb :SbF , SbF2\ SbF3o, and SbF4- 104 5.4.2 Reactions of SbF3°:SbF4-, and SbF3OH- 105 0 5.4.3 SbOF or Sb(OH)2F° 105 5.4.4 SbF3(s) 105 5.5 Sb(III) sulfate system 107
5.5.1 SbOSO4- 107 5.5.2 Sb2(SO4)3(s) 107 5.6 Antimony(III) sulfide system 108 2 0 5.6.1 Sb2S4 -, HSb2S4-, and H2Sb2S4 108 5.6.2 Sb2S3(stibnite) 109
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5.7 Redox reactions 110
5.7.1 Sb(cr)/Sb(OH)3° Ill 5.7.2 Sb(in)/Sb(V) .....112
5.7.3 Sb(cr)/Sb(OH)5° 113 5.7.4 Additional data compiled for the redox potential of antimony 113
5.8 Hydrolysis of antimony(V) 115
5.9 Sb2O5(precip) 118
5.10 Sb2O4(cr)andSb6O13(cr) 120
5.11 Other antimony(V) complexes and compounds 122
5.12 Comments on selected references 123
6 Lead 126
6.1 Hydrolysis of lead 126 6.1.1 PbOH+ 130
6.1.2 Pb(OH)2° 131 6.1.3 Pb(OH)3- 132 2 6.1.4 Pb(OH)4 - 132 3+ 6.1.5 Pb2OH 133 4+ 6.1.6 Pb4(OH)4 134 2+ 6.1.7 Pb3(OH)4 135 3+ 6.1.8 Pb3(OH)3 and Pb3(OH)5+ 136 4+ 6.1.9 Pb6O(OH)6 , 137 6.1.10 Additional equilibrium data compiled for the lead hydroxide system 138 6.2 Solid lead-oxide/hydroxide phases 144 6.2.1 PbO(litharge) and PbO(massicot) 144 6.2.2 Precipitated lead hydroxide 145
6.2.3 Pb(OH)2(s) 145 6.2.4 Additional data for lead hydroxide/oxide compounds 145 6.3 Lead chloride system 149 6.3.1 Lead chloride complexes 149
6.3.2 PbCl2(s) 156 6.3.3 Mendipite: Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr) 156 6.3.4 Laurionite and paralaurionite 157 6.3.5 Additional equilibrium data compiled for the lead chloride system 158 6.4 Lead fluoride system 163
6.4.1 PbF+andPbF2° 164 6.4.2 PbF2(s) 165 JNC TN8400 99-011
6.5 Mixed lead fluoride chloride system 168 6.5.1 PbFCl0 169 6.5.2 Matlockite (PbClF(cr)) 169 6.6 Lead carbonate system 171
6.6.1 PbCO3° 173 2 6.6.2 Pb(CO3)2 - 174 4 6.6.3 Pb(CO3)3 -andPb(CO3)46- 175 6.6.4 PbCO3OH- 175 + 6.6.5 PbHCO3 , Pb(HCO3)2° and Pb(HCO3)3- 175 6.6.6 PbCO3(cerrusite) 176 6.6.7 Pb3(CO3)2(OH)2(hydrocerrusite) 177 6.6.8 Plumbonacrite 178
6.6.9 PbCO3PbO(s) and PbCO3(PbO)2(s) 178 6.6.10 Phosgenite 178 6.6.11 Additional data for the lead carbonate system 179 6.7 Lead nitrate system 184 6.7.1 Lead nitrate complexes 186
6.7.2 Pb(NO3)2(s) and PbOHNO3(cr) 188 6.7.3 Additional data compiled for the lead nitrate system 188 6.8 Lead phosphate system 191 + 6.8.1 PbH2PO4 and PbHPO4° 192 6.8.2 PbHPO4(cr) and Pb3(PO4)2(s) 192 6.8.3 Pb(H2PO4)2(s), Pb4(PO4)2O(s) and plumbogummite (PbAl3(PO4)2(OH)5(cr)) 193 6.8.4 Pyromorphites: Pb5(PO4)3Cl(s), Pb5(PO4)3F(s), and Pb5(PO4)3OH(s) 193 6.8.5 Pb10(PO4)6(OH)2(hydroxylapatite) 193 6.8.6 Additional data compiled for the lead phosphate system 194 6.9 Lead sulfate system 197 6.9.1 Lead sulfate complexes 198
6.9.2 PbSO4(anglesite) 199 6.9.3 Hinsdalite: PbAl3PO4SO4(OH)6(cr) 200 6.10 Lead sulfide system 204 6.10.1 Lead sulfide complexes 205 6.10.2 Galena (PbS) 205 6.10.3 Additional data compiled for the lead sulfide system 206 6.11 Redox equilibria 208 6.11.1 Pb2+/Pb(cr) 208 6.11.2 Pb2+/Pb4+ 208
6.11.3 PbO2andPb3O4 208 6.11.4 Data compiled for the lead redox system 209
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6.12 Lead(IV) 211
6.13 Comments on selected references 213
7 Bismuth 219
7.1 Hydrolysis of bismuth 219 7.1.1 BiOH2+ 222 + 7.1.2 Bi(OH)2 223 7.1.3 Bi(OH)3° 224 7.1.4 Bi(OH)4- 225 6+ 7.1.5 Bi6(OH)i2 226 7+ 6+ 5+ 7.1.6 Bi9(OH)2o , Bi9(OH)2i , and Bi9(OH)22 227 5 7.1.7 Bi3(OH)4 + 227 3 7.1.8 Bi6(OH)15 + 227 7.1.9 Additional equilibrium data compiled for the bismuth hydroxide system 228 7.2 Solid bismuth-oxide/hydroxide 232
7.2.1 a-Bi2O3(cr) 232 7.2.2 Precipitated Bi(OH)3(s) 233 7.2.3 Additional data compiled for solid bismuth oxides/hydroxides 233 7.3 Bismuth chloride system 236 7.3.1 Bismuth chloride complexes 236
7.3.2 BiOCl(s) and Bi(OH)2Cl(s) 242 7.3.3 BiCl3(s) 243 7.3.4 Additional equilibrium data compiled for the bismuth chloride system 243 7.4 Bismuth perchlorate 247 2+ 3+ 7.4.1 BiC104 or Bi(H2O)6 C1O4- 247 7.4.2 BiOClO4(precip) 247 7.5 Bismuth fluoride system 248
7.6 Bismuth carbonate system 250
7.7 Bismuth nitrate system 251 7.7.1 Bismuth nitrate complexes 251
7.7.2 BiONO3(s) 255 7.7.3 Additional equilibrium constants compiled for the bismuth(III) nitrate system 256 7.8 Mixed bismuth nitrate and chloride system 259 7.8.1 Bismuth chloride nitrate complexes 259
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7.8.2 Additional equilibrium data compiled for the bismuth(III) chloride nitrate system 263 7.9 Bismuth phosphate system 264
7.10 Bismuth sulfate system 265 7.10.1 Bismuth sulfate complexes 265
7.10.2 Bi2(SO4)3(s) 265 7.10.3 Equilibrium data compiled for the bismuth sulfate system 265 7.11 BP+/Bi(cr) 267
7.12 Comments on selected references 269
8 Niobium 275
8.1 Hydrolysis of niobium(V) 275 8.1.1 Monomeric Nb(V) species 275 8.1.2 Polymeric Nb(V) species 277 8 7 8.1.2.1 Protonation at pH > 8: N^Ch 9 -, Nb6Oi 9H -, 6 5 Nb6O19H2 -, and Nb6Oi 9H3 - 277 8 12 8.1.2.2 Very alkaline solutions: Nb^ 2(OH)4 -, Nb4O[6 - and NbO2(OH)43- 280 6 8.1.2.3 Neutral and acidic solutions: H^Nb] 2C>36 ~, 7 8 9 H5Nb12036 -, H4Nb12O36 -, andH3Nb12O36 - 280 8.1.3 Additional equilibrium data compiled for the niobium(V) hydroxide system ' 280 8.2 Solubility of solid niobium pentoxide ..- 283
8.2.1 Additional equilibrium data compiled for Nb2C>5(s) 285 8.3 Solid niobium phases: redox equilibria 286
8.4 Other niobium(V) complexes and compounds 287
8.5 Comments on selected references 289
9 Palladium 290
9.1 The redox pair Pd27Pd(s) 290
9.2 Hydrolysis of palladium(n) 292 9.2.1 Hydrolysis of palladium(II) 292 9.2.2 Additional equilibrium data compiled for the hydrolysis of palladium(II) 293 9.3 Solid palladium(II)-oxide/hydroxide 296
9.3.1 Pd(OH)2 (precip) 296
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9.3.2 PdO(cr) 296 9.4 Chloride complexes of palladium(II) 298 9.4.1 Chloride complexes 298 9.4.2 Additional equilibrium data compiled for palladium chloride complexes 299
9.5 Amino complexes of palladium(II) 302
9.6 Conclusions 302
9.7 Comments on selected references: 303
10 References 309
XIV JNC TN8400 99-011
1 Introduction
1.1 Background
JNC is planning to submit a new Performance Assessment Report by March 2000. Besides informations on geology, repository design etc., basic data on radionuclide behavior are required for the safety analysis. These basic data include chemical thermodynamic data in aqueous media, requiring the establishment of a chemical thermodynamic database, PNC-TDB, for the interaction of key radionuclides with significant ligands under relevant conditions.
JNC has defined two groups of key elements for the next Performance Assessment Report. They are termed '1st priority elements' and '2nd priority elements', respectively, and are comprised of the following elements:
• 1st priority elements: Pu, U, Np, Th, Am, Ra, Sn, Zr, Ni, Pd, Tc, Se, Pa, Cm and Cs. • 2nd priority elements: Sm, Ac, Po, Pb, Nb, Bi and Sb.
Ligands considered by JNC include hydroxide, carbonate, chloride, fluoride, sulfate and phosphate. In addition, solid-solution modeling may be considered for Ra, and possibly other elements. Since the aqueous chemistry of Cs is very simple, this element is not treated within the JNC-TDB development.
1.2 Organization of the JNC-TDB project
The JNC-TDB development project involves JNC staff members, an external expert group, and a Japanese expert advisory group. The expert group consists of Dr. H. Wanner, HSK; Dr. D. Rai, Batelle PNNL; Prof. G. Choppin, Florida State University; and Dr. M. Ochs, BMG. The roles and responsibilities of each member of the expert group are defined in Table 1.1. The Japanese expert advisory group comprises Prof. Tochiyama, Tohoku University, Prof. Moriyama, Kyoto University, and Dr. Nakayama, JAERI. The treatment of elements by the different groups is detailed in Table 1.1. JNC TN8400 99-011
Table 1.1: Assigned responsibilities for the JNC-TDB project for the complete 2-year period of FY96-97:
Elements Responsibility Timeframe
Pu, Ac, Th, u, Np, Pa, Am, Cm PNNL i FY96-97 Pd, Pb, Sb, Bi, Nb, Sn BMG2 FY97 Sm: , Se JNC (review by PNNL 1) FY96-97 Ra, Zr, Tc, Po JNC (review by BMG2) FY96-97 1 This includes Dr. D. Rai and his co-workers at Batelle PNNL, as well as Dr. G. Choppin and his co-workers at Florida State University 2 This includes Dr. M. Ochs, Dr. B. Lothenbach, and other personnel at BMG Engineering Ltd, with Dr. H. Wanner functioning as external advisor.
The key elements Sn, Sb, Pb, Bi, Nb, Pd are the subject of the present report. For these elements, element-specific datasets are developed in this report based exclusively on experimental studies published in the literature, rather than relying on existing compilations.
Seperate reports on actinide elements and Ac are being prepared by Dr. D. Rai and co-workers. The elements Sm, Ni, Se, Ra, Zr, Tc, and Po are treated by JNC. JNC TN8400 99-011
2 Standards and Conventions
2.1 Symbols, units, notations and conversion factors
2.1.1 Symbols and notation
The symbols for physical and chemical quantities used in this report are summarized in Table 2.1.
Table 2.1: Symbols and notation
AfG° the standard molar Gibbs energy of formation [kJ mol"] ] 1 ArG° the molar Gibbs energy of a reaction [kJ mol" ] AfH° the standard molar enthalpy of formation [kJ mol"1] S° the standard molar entropy [J K"1 moF] T absolute temperature [K] R molar gas constant (8.3145 J K"1 moF] n number of electrons involved in a redox reaction CB concentration of a solute B in [mol/L] me concentration of a solute B in [mol/kg solvent] YB activity coefficient of a substance B I ionic strength [mol/L]
Im ionic strength [mol/kg solvent] p Factor for the conversion of molarity, CB, to molality, ms, of substance B [dm3 solution per kg H2O] VB stoichiometric coefficient of a substance B (negative for reactants, positive for products) JNC TN8400 99 -Oil
2.1.2 Compilation of thermodynamic data
In the Sections 4 -9 of this report, thermodynamic data are compiled in tables. The following conventions have been used throughout:
'Reference': The references are ordered chronologically and alphabetically by the first two authors within each year. A more detailed description is given in Chapter II. 1.8 in [1992GRE/FUG]. The references are given in Section 3 of this report.
'Comments': T indicates the temperature (K) to which the given constant refers. Many comments were imported directly from the NEA database. I indicates the conditions under which the constant was determined, e.g., I = 0.1-1.
T: I indicates the ionic strength to which the given constant refers. This value is often, but not always, identical with the ionic strength given under 'Comments'.
'Medium': indicates, where available, the electrolyte in which the given constant is measured or refers to.
'Method': Abbreviations for the method of measurement are listed in Table 2.2.
Table 2.2: Abbreviations for experimental methods
cat = cation exchange col = colorimetric analysis con = conductivity measurements el = electrophoresis emf = electromotive force measurements at high temperatures extr = extraction fe = fluoride selective electrode kin = kinetic measurements n/a = method not known to the reviewers NMR= "FNMR pol = polarography pot = potentiometry se = sulfide electrode
SO2 = pSO2 measurements at high temperatures sol = solubility measurements sp = spectrophotometry, NMR tit = titration (evaluation of equilibrium through pH only) JNC TN8400 99-011
2.1.3 Phase designators
Chemical formulae may refer to different chemical species and are often required to be specified more clearly in order to avoid ambiguities. For example, PbCC>3 occurs as a solid and as aqueous complex. The distinction between the different phases is made by phase designators that immediately follow the chemical formula and appear in parentheses. The only formulae that are not provided with a phase designator in this report are aqueous ions. The use of the phase designators is described below:
The designator (1) is used for pure liquid substances, e.g.,
The designators ('name'), (cr), (precip) and (s) is used for solid substances. When the solid has a common name, ('name') is used, e.g. PbCO3(cerrusite). (cr) is used when it is known that the solid is crystalline, e.g. PbOHCl(cr). (precip) is used when it is known that the solid was precipitated from solution. Otherwise, where no such information is available, (s) is used.
For aqueous species no designators are used, e.g. PbOH+ or PbCO3°.
2.1.4 Physical constants
The fundamental physical constants are taken from [1992GRE/FUG] and are listed in Table 2.3.
Table 2.3: Fundamental physical constants. These values were taken from [1992GRE/FUG].
R molar gas constant 8.3145 J R-1 mol"1 F Faraday constant 96 485 C moW
2.1.5 Equilibrium constants
The IUPAC has not explicitly defined the symbols and terminology for equilibrium constants of reactions in aqueous solutions. In this report the conventions (based on the work of Baes and Mesmer, [1976BAE/MES]) given below have been used throughout.
Formation of an aqueous complex: JNC TN8400 99-011
• Formation of a solid:
Conventionally, equilibrium constants involving a solid are denoted as 'solubility constants' rather than as formation constants of the solid. An index 's' to the equilibrium constant indicates that the constant refers to a solubility process, as shown below:
m n MmLn(s)^mM + nL KSo=[M} [L]
Kso is the conventional solubility product and the subscript '0' indicates that the equilibrium reaction involves only uncomplexed aqueous species.
In this report, the formation of solid is treated analogously to the formation of aqueous species. I.e., the inverse solubility product is used, denoted with an asterisk:
mM + nL<^> MmLn(s) K*So =
Further notations used in this report are compiled in Table 2.4.
Table 2.4: Reactions
P measured cumulative formation constant for a reaction1 (in molarity units)
(3m measured cumulative formation constant for a reaction corrected from molarity to molaliry units (3° cumulative equilibrium constant valid at I = 0 Pb measured cumulative equilibrium constant for a reaction involving OH~ instead of H+ l. K consecutive (stepwise) equilibrium constant for a reaction Kso Solubility product of a solid 2. 2+ 2+ l\Pb ] Example: PbO(cr) + 2H+ <=» Pb + H2O; Kso = -^
K*so Formation constant of a solid 2.
2+ Example: Pb + H2O <=> PbO(cr) + 2H+; K so = }—=S [Pb2+] Kbso Solubility product of a solid involving OH~ instead of H+ '. K*bso Formation constant of a solid involving OH~ instead of H+ ]. ' Hydrolysis reactions are formulated in this report using H+ as component and not OH". 2 In this report, reactions involving a solid are normally formulated as formation reactions. JNC TN8400 99-011
2.1.6 Redox reactions
Redox reactions are usually quantified in terms of their electrode (half cell) potential E. E is identical to the electromotive force (emf) of a redox reaction that involves the standard hydrogen electrode as an electron donor or acceptor. In this review, electrode potentials are given as reduction potentials relative to the standard hydrogen electrode which acts as an electron donor.
The standard redox potential, E°, is related to the change of the Gibbs energy ArG° and the equilibrium constant K as outlined below:
ArG (at 298.15 K) nF nF
2.2 Data selection criteria
To assure a traceable data selection a list of selection criteria is given below. All studies selected in this review should fulfill the following criteria.
• Experimental study • I = constant • No complex formation with electrolyte (or corresponding correction possible) • All relevant species included (e.g. polymers) • Experimental details reported • T-298K
For more detailed information of the individual papers and additional criteria, see discussion of the individual species in Section 4 - 9 of this report.
2.3 Ionic strength corrections
Thermodynamic data always refer to a selected standard state. The standard state for a solute B in a solution is a hypothetical solution at the standard state pressure (0.1 MPa) and the standard temperature (298.15 K) in which me = mo = 1 mol/kg, and in which the activity coefficient JQ is unity. However, for many reactions, measurements cannot be made accurately (or at all) in dilute solutions from which the extrapolation to the standard state would be simple. In this report, thermodynamic data were extrapolated to the standard state (1=0) using the specific ion interaction equation (SIT) as described in [1992GRE/FUG]. An extensive description of this method and its use can be found in Appendix B of [1992GRE/FUG] and [1995SIL/BID]. Shortly, the correction consists of an extended Debye-Hiickel expression, in which the activity coefficients of the reactants and products depend only on the ionic charge of the reactants and JNC TN8400 99 - Oil the ionic strength of the solution, but it accounts for the medium specific properties by introducing ion pairing between the medium ions and the species involved in the equilibrium reactions. The correction of the measured data to 1=0 is made with the following equation:
Debye - Hueckel term
1 + 1.5X ~\lJ/Am.
Im = /xp where I ionic strength [mol/L]
Im ionic strength [mol/kg solvent] p Factor for the conversion of molarity, CB, to molality, me, of substance B [dm3 solution per kg H2O] P measured cumulative formation constant for a reaction expressed in molarity units
|3m measured cumulative formation constant for a reaction corrected from molarity to molality units Xv sum of the stoichiometric coefficients of the reaction.
The factors for the conversion of molarity, CB, to molality, mg, of a substance B for the different electrolytes at 298.15 K were taken from Table II.5 in [1995SIL/BID] and from Table II-l in [1976BAE/MES].
2.4 Auxiliary Data
In this section the thermodynamic data for auxiliary compounds and complexes are compiled. Many of the these auxiliary species are used in the evaluation of the recommended data given in Section 2. It is therefore essential to always use these auxiliary data in conjunction with the selected data. The use of other auxiliary data can lead to inconsistencies and erroneous results. Table 2.5 contains the selected thermodynamic data of the auxiliary species considered in this review (all taken from the recent NEA publication of [1995SIL/BID]). JNC TN8400 99-011
2.4.1 Selected thermodynamic data for auxiliary species
Table 2.5: Selected thermodynamic data for auxiliary species taken from [1995SIL/BID] complex AfG° [kJ/mol] Reference ci- -131.217 [1995SIL/BID] F- -281.523 [1995SEL/BID] HF° -299.675 [1995SIL/BID] 2 SO4 - -744.004 [1995SEL/BID] HS- 12.243 [1995SIL/BID] H2S° -27.648 [1995SIL/BID] H2S(g) -33.443 [1995SIL/BID] NO3- -110.794 [1995SIL/BID] CO32- -527.899 [1995SIL/BID] HCO3- -586.845 [1995SIL/BID] CO2° -385.970 [1995SIL/BID] CO2(g) -394.373 [1995SIL/BID] PO43- -1025.491 [1995SIL/BID] HPO42- -1095.985 [1995SDL/BID] [1995SIL/BID] H2PCV -1137.152 H2O -237.140 [1995SIL/BID] Ca2+ -552.806 [1995SIL/BID] Na+ -261.953 [1995SEL/BID] K+ -282.510 [1995SIL/BID]
2.4.2 Conversion of AjG ° values to equilibrium constants
Experimental papers measure and indicate in most cases log (3 values. In a few cases the results of experimental studies are given as AfG° values. To be able to compare these values with the log B or log K values selected in this report, these AfG° values were converted to B or K values according to
Arc° using the AfG° values of the auxiliary species (Table 2.5) and the AfG° values of the master 2+ 2+ 3+ species Sn , Sn(OH)4°, Sb(OH)3°, Sb(OH)5°, Pb , Bi selected in this report (Table 2.6). In many compilations, which were only used for comparison in this report, AfG° values are compiled instead of log p or log K values (e.g. in [1982WAG/EVA]). In general, these AfG° values were converted to log P or log K values using the AfG° values of the auxiliary species 2+ (Tables 2.5) and the AfG° values of the master species Sn , Sn(OH)4°, Sb(OH)3°, Sb(OH)5°, JNC TN8400 99-011
Pb2+, Bi3+, Pd2+ (Table 2.6) selected in this report. However, some compilations also report AfG° values for the master species (Table 2.7). In that case the AfG° values given in the respective compilation is used.
2+ Table 2.6: Thermodynamic data for the master species Sn , Sn(OH)4°, Sb(OH)3°, 2+ 3+ 2+ Sb(OH)5°, Pb , Bi , Pd selected in this report. For Nb(OH)50 no thermodynamic data was selected.
complex AfG° [kJ/mol] Reference
Sn2+ -26.42 this report Sn(OH)4° -944.16 this report Sb(OH)3° -643.0 this report Sb(OH)5° -992.62 this report Pb2+ -24.24 [1989COX/WAG] Bi3+ 95.55 this report Pd2+ 187.6 * this report 1 tentative value
Table 2.7: ThermodynamiThermodvnamic data for the master species Sn2+, Sn(OH)4°, Sb(OH)3°, Sb(OH)5°, Pb2+, Bi3+, Pd2+, given in previous compilations
complex AfG° [kJ/mol] Reference
Sn2+ -26.29 [1952LAT] -27.89 [1978COD] -29.89 [1980BEN/TEA] -27.2 [1982WAG/EVA] -27.14 [1985BAB/MAT] -27.2 [1985GAL] -27.23 [1988PHI/HAL] -27.62 [1989COX/WAG]
Sn(OH)4° -934.8 [1984KEL/HOU] -950.6 [1988PHI/HAL]
Sb(OH)3° -647.16 [1985BAB/MAT] -644.8 [1985PAS] -644.66 [1986ITA/NIS] -643.63 [1990SHI/ZOT] -643.9 [1994AKI/ZOT]
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Table 2.7: continued
Pb2+ -24.24 [1989COX/WAG] -24.34 [1952LAT] -24.44 [1969HEL] -24.34 [1977PAU] -24.00 [1978COD] -24.4 [1978ROB/HEM2] -24.43 [1980BEN/TEA] -24.06 [1981HEL/KIR] -24.39 [1981STU/MOR] -24.34 [1982PAU] -24.43 [1982WAG/EVA] -24.42 [1983LAN] -24.42 [1983SAN/BAR] -24.4 [1984 VIE/TAR] -24.42 [1985BAB/MAT] -24.34 [1985GAL] -24.23 [1985RAI/RYA] -24.00 [1988PHI/HAL]
Pb4+ 302.43 [1952LAT] -302.57 [1983LAN] 302.44 [1985GAL]
Bi3+ 82.79 [1968 ROB AVAL] 82.80 [1982WAG/EVA] 91.79 [1985BAB/MAT] 91.82 [1985LOV/MEK]
Nb(OH)5° -1448 [1985UDU/VEN] -1448 [1982WAG/EVA]
HNbO30 -985.3 [1985BAB/MAT]
Pd2+ 190.34 [1952LAT] 176.56 [1967IZA/EAT] 177.37 [1968GOL/HEP] 176.53 [1980BEN/TEA] 176.5 [1982WAG/EVA] 176.53 [1985BAB/MAT] 176.5 [1985COL] 176.47 [1988PHI/HAL]
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3 Selected data
This chapter presents the thermodynamic data set for bismuth, niobium, lead, palladium, antimony, and tin species selected in this review. The Tables 3.1 to 3.6 contain the recommended thermodynamic constants of chemical equilibrium reactions by which bismuth, niobium, lead, palladium, antimony, and tin compounds and complexes are formed. Besides treating hydrolysis in detail, this review focuses on the formation of complexes or compounds with chloride, fluoride, carbonate, nitrate, sulfate and phosphate. Other important inorganic ligands (sulfide for lead and antimony, ammonia in the case of palladium) are also included. The present report does not include any compounds or complexes containing organic ligands.
It should also be noted that the data set presented in this section may not be 'complete' for all conceivable systems and conditions. Gaps and uncertainties are pointed out in Section 2.2 and in various paragraphs in the Sections 4 -9.
3.1 Selected data for Sn, Sb, Pb, Bi, Nb and Pd
The Tables 3.1 to 3.6 contain the recommended thermodynamic data and should be read as follows:
Each element has at least one master species whose identity is given in the heading of each table. Redox sensitive elements may have more than one master species, allowing to separate the modeling of the different oxidation states, e. g. in case of kinetic inhibition. Each table contains the formation constants of all the complexes and solids proposed in this report.
Example for reading the contents of the database:
• Species SnOH+ (see Table 3.1) is composed of the following components: 1 x component Sn2+, and -1 x component H+. Addition of 1 x H2O1 gives the correct formation reaction for SnOH+ 2+ + + Sn + H2O(1) = SnOH + H with a formation constant log P of -3.75.
1 H2O has an activity of 1 in aqueous solutions according to the conventions used and is therefore not included in the tables.
12 JNC TN8400 99-011
Table 3.1: Selected thermodynamic data for reactions involving tin compounds and reactions as selected in Section 4 of this report. All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
2 + 2 components => Sn H+ ci- F- NO3- SO4 " e~ species logP
SnOH+ -3.75 1 -1
Sn(OH)2« -7.71 1 -2
Sn(OH)3" -17.54 1 -3 2+ Sn3(OH)4 -6.51 1 3 -4 SnCl+ 1.65 1 1
SnCl2° 2.31 1 2 SnCl3' 2.09 1 3 SnOHCl0 -2.27 1 -1 1 SnF+ 4.46 1 1 1
SnF2° 7.74 1 1 2 SnFf 9.61 i 1 3 + SnNO3 1.25 1 1
Sn(NO3)2» 1.74 1 2
Sn(NO3)3- 1.37 1 3 2 Sn(NO3)4 " 0.30 i 1 4 * SnSO4° 2.91 i 1 1 2 Sn(SO4)2 " 2.83 1 1 2
Sn(OH)2(precip) -2.82 1 -2 SnO(cr) -2.41 1 -2 SnOHCl(s) 2.42 1 -1 1
Sn(cr) -4.63 1 2 Sn4+ -51 1 -2
Sn(OH)4° -5.4 i 1 -4 -2 1 tentative values
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Table 3.1: continued components => Sn(OH)4° species logP
Sn4+ 0.4 i 1 4
Sn(OH)5" -7.97 1 -1 2 Sn(OH)6 " -18.40 1 -2
SnO2(precip) 7.46 1 0
SnO2(cassiterite) 8.0 1 0 tentative value (see Section 4.4.3); values for other protonated species are missing.
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Table 3.2: Selected thermodynamic data for reactions involving antimony compounds and reactions as selected in Section 5 of this report All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
+ components => Sb(OH)3° H ci- F- HS~ e~ species logP 11
Sb3+ -0.73 i 1 3 SbOH2+ 0.83 1 1 2 + Sb(OH)2 1.30 1 1
Sb(OH)4- -11.93 1 -1
Sb2(OH)6° 0.08 i 2 0 SbCl2+ 2.78] 1 3 1 + SbCl2 3.27' 1 3 2 SbF2+ 6.48' 1 3 1 + SbF2 12.65 i 1 3 2
SbF3° 18.36' 1 3 3 2 Sb2S4 " 42.53 2 2 4
HSb2S4~ 52.18 2 3 4
H2Sb2S4° 57.00 2 4 4
S b2 0, (valentinite) 8.72 2 0
Sb2S,(stibnite) 55.14 2 '3 3
Sb(cr) 11.99 1 3 3
Sb(OH)5 -21.84 1 -2 -2 1 tentative values
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Table 3.2: continued
+ components => Sb(OH)s° H species logP II
Sb(OH)6" -2.72 1 -1 4 Sb12(OH)64 " 20.34 12 -4 5 Sb12(OH)65 - 16.72 12 -5 6 Sb12(OH)66 - 11.89 12 -6 7 Sb12(OH)67 " 6.07 12 -7
Sb2O5(precip) 7.40 2 0
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Table 3.3: Selected thermodynamic data for reactions involving lead compounds and reactions as selected in Section 6 of this report. All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
p 2 + 2 3 2 components =$ b ci- F- CO3 - NO3- PO4 -SO4 " HS- species logp 11
PbOH+ -7.51 1 -1
Pb(OH)2° -16.95 1 -2
Pb(OH)3" -28.02 1 -3 3+ Pb2OH -7.18 2 -1 4+ Pb4(OH)4 -20.63 4 -4 2+ Pb3(OH)4 -22.48 3 -4 + Pb3(OH)5 -30.72 3 -5 4+ Pb6(OH)8 -42.68 6 -8 PbCl+ 1.55 1 1
PbCl2° 2.00 1 2
PbCl3" 2.01 1 3 2 PbCl4 " 1.35 1 4 PbF 2.27 1 1
PbF2° 3.01 1 2 PbFCl0 3.55 1 1 1
PbCO3° 7.30 1 1 2 Pb(CO3)2 - 10.13 1 2 + PbNO3 1.06 1 1
Pb(NO3)2° 1.48 1 2
Pb(NO3)3" 0.76 i 1 3
PbHPO4° 15.45 i 1 1 1 + PbH2PO4 21.05 i 1 2 1 PbSO4° 2.82 1 1 2 Pb(SO4)2 " 2.37 i 1 2 Pb(HS)2° 12.34 i 1 2 Pb(HS)3" 13.59 ' 1 3
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Table 3.3: continued
2+ 2 3 2 components => Pb H+ Cl~ F~ CO3 " NO3- PO4 -SO4 - HS~ e~ species log p
PbO(red, litharge) -12.68 1 -2 PbO(yellow, massicot) -12.96 1 -2
Pb(OH)2(precip) -13.05 1 -2 PbCl2(s) 4.81 1 2 PbOHCl(cr) -0.62 1 -1 1
PbF2(s) 7.52 1 2 PbFCl(matlockite) 8.82 1 1 1
PbCO3 (cerrusite) 13.23 1 1 Pb3(CO3)2(OH)2 17.64 1 3 -2 2 (hydrocerrusite)
Pb10(CO3)6(OH)6O 41.21 l 10 -8 6 (plumbonacrite)
PbOHNO3(cr) -2.94 1 1 -1 PbHPO4(s) 23.78 1 1 1 1 Pb3(PO4)2(s) 44.40 1 3 2 Pb4(PO4)2O(s) 37.09 1 4 -2 2 Pb(H2PO4)2(s) 48.94 1 1 4 2 Pb5(PO4)3OH 62.80 1 5 -1 3 (hydroxy pyromorphite)
Pb5(PO4)3Cl 84.40 1 5 1 (chloro pyromorphite)
Pb5(PO4)3F 71.60 1 5 1 (fluoro pyromorphite)
PbSO4(anglesite) 7.81 1 PbS (galena) 12.17 1 -1
Pb(cr) -4.25 1 2
] PbO2(s) -48.98 1 -4 -2 J Pb3O4(s) -70.98 3 -8 -2 1 tentative values JNC TN8400 99-011
Table 3.4: Selected thermodynamic data for reactions involving bismuth compounds and reactions as selected in Section 7 of this report. All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
3 + 2 components => Bi H+ ci- CO3 " NO3- e- species logP
BiOH2+ -0.92 1 -1 + Bi(OH)2 -2.56 1 -2
Bi(OH)3 -5.31 1 -3
Bi(OH)4- -18.71 1 -4 6+ Bi6(OH)12 1.34 6 -12 7+ Bi9(OH)20 -1.36 9 -20 6+ Bi9(OH)21 -3.25 9 -21 5+ Bi9(OH)22 -4.86 9 -22 5+ Bi3(OH)4 -0.80 1 3 -4 BiCl2+ 3.65 1 1 + BiCl2 5.85 1 2
BiCl3 7.62 1 3
BiCl4" 9.06 1 4 2 BiCl5 " 8.33 i 1 5 3 l BiCl6 " 7.64 1 6 2+ BiNO3 1.97 1 1 + Bi(NO3)2 2.95 1 2
Bi(NO3)3 3.62 1 3
Bi(NO3)4- 3.09 1 4 + BiClNO3 5.16 1 1 1
BiCl(NO3)2° 5.28 1 1 2
BiCl2NO3° 6.86 1 2 1
BiCI2(NO3)2- 5.75 1 2 2
BiCl3NO3" 8.09 1 3 1
a-Bi2O3(cr) -0.76 2 -6 BiOCl(s) 8.47 1 -2 1
(BiO)2CO3(cr) 14.27 1 2 -4 1
(BiO)4(OH)2CO3(cr) 8.68 i 4 -10 1
BiONO3(s) 2.75 1 -2 1
Bi(cr) 16.74 1 3 ! tentative values
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Table 3.5: Selected thermodynamic data for reactions involving niobium compounds and reactions as selected in Section 8 of this report. All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
+ components => Nb(OH)5° H species log P _J
Nb(OH)6- -6.6 1 -1
Nb2O5(s) 16.0 i 2 1 tentative value (see Section 8.2)
Nb(0H)5° and Nb(OH)6 are hypothetical species (see Section 8.2).
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Table 3.6: Selected thermodynamic data for reactions involving palladium compounds and reactions as selected in Section 9 of this report. All ionic species listed in this table are aqueous species. The data refer to the reference temperature of 298 K and to the standard state, i.e., a pressure of 0.1 MPa and, for aqueous species, infinite dilution (1=0).
2+ + components => Pd H cr NH3 e~ species log [3 pdcr 5.1 1 l
PdCl2° 8.3 1 2
PdCl3" 10.9 1 3 2 PdCl4 - 11.7 1 4 2 PdCl3OH ~ 2.5 1 -1 3 2 PdCl2(OH)2 - -7.0 i 1 -2 2 2+ Pd(NH3) 9.6 1 1 2+ Pd(NH3)2 18.5 1 2 2+ Pd(NH3)3 26.0 1 3 2+ Pd(NH3)4 32.8 1 4
Pd(cr) 32.9 ' 1 2 1 tentative values
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3.2 Gaps and uncertainties
3.2.1 Uncertainties
Some authors report the standard deviations of their experimentally determined formation constants, but these do not represent the quality of the reported values in absolute terms. It is important not to confuse the statistical standard deviation with the accuracy. The standard deviation is calculated from the dispersion of equally weighted points, while the accuracy reflects the reliability and reproducibility of an experimental value and also includes all kinds of systematic errors. The estimation of systematic errors of a method is difficult and can only be made by a person who is familiar with the experimental method. In many cases the determination of the standard deviations is not possible because either only one or two data points are available, or the authors did not report the individual values. In this report no attempt was made to calculate systematically the uncertainties connected to experimentally determined formation constants. For further indication of the reliability of different methods see discussions for the individual species in Sections 4-9.
3.2.2 Gaps
For the user it is important to consider that the selected data set presented in Section 3.1 may not be 'complete' with respect to all conceivable systems and conditions; there are gaps in the information, particularly concerning complex formation of tin(FV), antimony and niobium with inorganic ligands as well as the hydrolysis and the solubility of the oxides of niobium and palladium. For each individual key element discussed in this report, gaps are listed below. While some missing data may not be important from a practical point of view (e.g. the exact solubility of a very soluble solid, e.g. SnCl2(s); or the formation constants for complexes that will only be important in very concentrated solutions, e.g. the formation of tin(IV) nitrates), other missing data (e.g. the formation of solids with sulfide) may be more important. Gaps considered important by the authors of this review are printed in bold.
The gaps also are shortly discussed in the respective sections of the key elements in this report. This information may be used as a basis for the assignment of research priorities.
22 JNC TN8400 99-011
Tin(II): Gaps in the tin(II) thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
SnCl2(s) very soluble SnF2(s) soluble Sn(NC>3)2(s) very soluble SnSC>4(s) moderately soluble 2 SnCO3(s), SnCO3° and Sn(CO3)2 - Sn3(PO4)2(s) SnSO4(s) SnS(s)
Tin(IV): Gaps in the tin(IV) thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
- chloride complexes and solids ' - fluoride complexes and solids l - nitrate complexes and solids J l - sulfate complexes and solids 1 4+ 3+ 2+ + - Sn , SnOH , Sn(OH)2 , Sn(OH)3 - carbonate complexes and solids - phosphate complexes and solids - sulfide complexes and solids - redox equilibrium Based on the data listed in Tables 4.7 to 4.9, it is our feeling that the complex formation of Sn(IV) with these ligands is rather weak as compared to the hydrolysis of Sn(IV).
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Antimony(III): Gaps in the antimony(III) thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
sulfate complexes and solids weak complexes nitrate complexes and solids no information available carbonate complexes and solids no information available phosphate complexes and solids no information available
Antimony(V): Gaps in the antimony(V) thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
5+ 4+ 3+ 2+ + - Sb , SbOH , Sb(OH)2 , Sb(OH)3 , or Sb(OH)4 stable at pH < 1 - chloride complexes and solids weak complexes - fluoride complexes and solids ? - nitrate complexes and solids ? - sulfate complexes and solids ? - carbonate complexes and solids ? - phosphate complexes and solids * ? - sulfide complexes and solids ?
24 JNC TN8400 99-011
Lead: Gaps in the lead thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
- PbC>2 (s), Pb3O4(s) very oxidizing conditions necessary - Pb(IV) hydrolysis very oxidizing conditions necessary
more information needed: - hydrocerrusite - phosphate complexes and solids
Bismuth: Gaps in the bismuth thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
B1CI3 (s) very soluble fluoride and sulfate complexes J BiF3 (s) ? BiPO4(s) Bi2(SO4)3(s) Bi2S3(s)
25 JNC TN8400 99-011
Niobium: Gaps in the niobium thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
chloride complexes and solids fluoride complexes and solids nitrate complexes and solids sulfate complexes and solids carbonate complexes and solids phosphate complexes and solids sulfide complexes and solids niobium hydrolysis in the neutral and acidic pH range, formation of polynuclear complexes under alkaline conditions ?
Palladium: Gaps in the palladium thermodynamic database. Gaps considered important by the authors of this review are printed in bold.
complex or solid Comments
- fluoride complexes and solids weak complexes - nitrate complexes and solids weak complexes - sulfate complexes and solids ? - carbonate complexes and solids probably not important in aqueous solutions - phosphate complexes and solids - sulfide complexes and solids - hydrolysis at very low Pd(II) cone (> lO10 M) to prevent formation of polymers
- Pd(OH)2 (s)
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4 Tin
Tin exists in several oxidation states from -IV to +IV in its compounds, but only the +11 and +IV states are important in its aqueous chemistry. The +IV state is particularly stable in natural environments, and cassiterite SnC>2(cr), is the major source of this element in nature [1976BAE/MES, 1985GAL, 1995WIB]. Little quantitative information exists on the solubility of SnC>2(cr), but it is rather low (< 10"8 M) in pH region 1 to 7. The stannous ion (Sn2+) is easily oxidized, and SnO(s) is much more soluble than SnC>2(s). The redox reactions of Sn(0) to Sn(II) and Sn(IV) are discussed in Section 4.4.
Reliable thermodynamic data concerning the complexes formation of Sn(OH)4° with inorganic ligands are quite limited. Data for the formation of complexes and solids of tin(IV) with the following inorganic ligands can be found in the literature: water, chloride, fluoride, phosphate, carbonate, sulfate and sulfide. Sufficient experimental data, however, to calculate equilibrium constants were only available for the hydrolysis of tin(IV) and for the solubility products of SnO2(precip) and SnO2(cr) (Sections 4.1 and 4.2).
For the formation of Sn2+ complexes or compounds with water, chloride, fluoride, nitrate, phosphate, sulfate and sulfide thermodynamic data are available. Equilibrium constants for the hydrolysis of tin(II) and the complex formation with chloride, fluoride, nitrate, and sulfate are calculated from experimental data and are given in the Sections 4.5 to 4.12. Also the solubility products of Sn(OH)2(precip), SnO(cr), and SnOHCl(s), are calculated from experimental data given in the literature (Sections 4.6 and 4.7).
4.1 Hydrolysis of tin(IV)
The knowledge on tin(IV) hydrolysis is limited. Amaya and coworkers [1997AMA/CHI,
1998ODA/AMA] measured the solubility of Sn(IV) in dilute NaC104 solutions between pH 2 and 13.5. They observed an increase of the solubility of Sn(IV) above pH = 7, while it remained constant at = 10"8 M between pH 2-7. This indicates the predominance of the uncharged Sn(0H)4° species. In strong acid or base tin(IV) solubility increases. As the hydrolysis of tin(IV) in acidic medium is not well defined, all log p values given in the following paragraphs refer to the Sn(OH)4° complex.
The data used for calculating the formation constants of tin(IV) hydroxide complexes are compiled in Table 4.1. Additional data for the tin(TV) hydroxide system data which were not chosen for the calculation of log (3° values in this report are compiled in Table 4.2 and 4.3.
27 JNC TN8400 99-011
Table 4.1: Experimentally determined equilibrium data compiled for the hydrolysis of n Sn(OH)4°, according to the equilibrium: mSn(OH)4° + nH2O «=> Snm(OH)4m+n- + nH*. These data were chosen for the evaluation of recommended values in the present report. Additional information see Section 4.14: 'Comments to selected references'. Method: sol = solubility measurements.
log 6m 4m+n Reference Comments I (M) Medium Method
log (3li5: Sn(OH)4° + H2O <=> Sn(OH)f + H+ -1.15 i [1998ODA/AMA] T= 298.15 K, 1=0.1 0.1 NaClO^ sol
2 log P1>6: Sn(OH)4° + 2H2O <=> Sn(OH)6 - + 2H+
-17.74 * [1998ODA/AMA] T= 298.15 K, 1=0.1 0.1 NaC104 sol 1 recalculated by [1998ODA/AMA] based on the experiments of [1997AMA/CHI]
4+ 3+ 2+ 4.1.1 Hydrolysis oftin(IV) under acidic conditions: Sn , SnOH , Sn(OH)2 and Sn(OH)3+
[1958JOH/KRA] and [1959JOH/KRA] have established, using ultracentrifugation, that Sn(IV) exists in the form of monomeric species in acidic and basic solutions. The only study carried out at 298 K is by [1971NAZ/ANT] (Table 4.2) who studied the hydrolysis of tin(IV) under acidic conditions in 1 M KNO3. [1971NAZ/ANT] used a spectrophotometric method in which the competition between the hydrolysis reaction and tomplexation with salicylfluorone was measured. The log (3 values obtained by [1971NAZ/ANT] for the protonation of Sn(OH)4° to 2+ 3+ 4+ Sn(OH)3+, Sn(OH)2 , Sn(OH) , and Sn are given in Table 4.2. The total Sn(IV) concentration in these experiments was 10-5 M. Considering the possible interaction of Sn(IV) with the nitrate ions of the electrolyte solution, these log 6 values can be considered as estimates, but not as exact values for the stability constants of Sn(IV) hydrolysis at I = 1. Nevertheless, based on the observation of [1971NAZ/ANT] and [1997AMA/CHI] it is clear that protonation of Sn(OH)4° will occur at pH < 2. [1934HUE/TAR] observed the predominance of the Sn4+ cation in solution containing > 0.5 M HC1. This indicates that the protonation of 4+ Sn(OH)4° to Sn takes place at slightly higher pH values than indicated by the data of [1971NAZ/ANT] (Table 4.2). The only other experimental values determined for Sn(OH)4° protonation is by [1970KUR/BAR] at 100 °C. In contrast to the observations at 25 °C, protonation of the Sn(OH)4° ion was observed by [1970KUR/BAR] already at a pH of 7.
From the different redox equilibria (cf. Section 4.4.3) a tentative log K° value of 0.40 can be 4+ calculated for the reaction Sn(OH)4° + 4H+ <=> Sn + 4H2O. For the other hydrolyzed Sn(IV) complexes it is not possible to propose any reliable thermodynamic values based on the experimental data available.
28 JNC TN8400 99-011
4.1.2 Hydrolysis oftin(IV) under neutral and alkaline conditions: Sn(OH)f and Sn(OH)(,2-
Hydrolysis of tin(IV) at 25 °C was investigated by [1970BAR/KLI] and Amaya and co-workers [1997AMA/CHI, 1998ODA/AMA]. In contrast to [1997AMA/CHI], who proposed the formation of Sn(OH)5~ and Sn(OH)g2" under alkaline conditions, [1970BAR/KLI] proposed only the formation of Sn(OH)5-. However, [1970BAR/KLI] calculated pH from the amount of NaOH added initially. These calculated pH values are probably significantly larger than the real pH in solution, making their results difficult to interpret.
Based on solubility measurements, Amaya and co-workers [1997AMA/CHI, 1998ODA/AMA] proposed the formation of Sn(OH)5~ and Sn(OH)62" under alkaline conditions. The log P15 and
Pi56 values given by [1998ODA/AMA] for I = 0.1 (these values were based on the experiments of [1997AMA/CHI], later recalculated by [1998ODA/AMA]) are given in Table 4.1. Extrapolation to 1=0 was made using SIT (see Section 2). Based on the interaction coefficients given in Appendix B.3 of [1992GRE/FUG] an As near or equal 0 can be assumed for both reactions.
Sn(OH)4° + H2O ^ Sn(OH)5- + H+ log pli5° = -7.97, As = 0 2 Sn(OH)4° + 2H2O ^ Sn(OH)6 " + 2H+ log Pi'6° =-18.40, Ae = 0
Some compilations [1984HOU/KEL, 1985GAL, 1992PEA/BER] list thermodynamic data for the species SnO32~. This is only a different notation for Sn(OH)62\
Table 4.2: Additional, experimentally determined equilibrium data compiled for the hydrolysis of Sn(OH)4°, n + according to the equilibrium: mSn(OH)4° + nH2O <=> Snm(OH)4m+n- + nH . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given at the end of the table or in Section 4.14. Method: sol = solubility measurements, sp = spectrophotfometry, and pot = potentiometry. log (3m,4m+n Reference Comments I(M) Medium Method
+ 4+ log P!X,.- Sn(OH)4° + 4H <=* Sn + 4H2O 0.87 2 [1971NAZ/ANT1 T= 298.15 K, 1=1 1 KNO, sp
+ + 3 log Pi,,,: Sn(OH)4° + 3H <=> SnOH + 3H2O 21.82 ] [1970KUR/BAR] T=373 K, I=dil NaOH sol 1.44 2 [1971NAZ/ANT1 T= 298.15 K, 1=1 1 KNO, sp
+ 2+ log Pi,:,.• Sn(OH)4° + 2H <=> Sn(OH)2 + 2H2O 14.69 1 [1970KUR/BAR] T=373 K, I=dil NaOH sol 1.55 2 [1971NAZ/ANTJ T= 298.15 K, 1=1 1 KNO, sp
+ + log Sn(OH)4° + H Sn(OH)3 + H2O 7.54 ' [1970KUR/BAR] T=373 K, I=dil NaOH sol 1.22 2 [1971NAZ/ANT] T= 298.15 K, 1=1 1 KNO3 sp 2.50 • [1981DAD/SOR] T= 573 K, I=n/a self medium sol 2.06 3 [1981DAD/SOR1 T= 298 K, I=n/a self medium sol
29 JNC TN8400 99-011
Table 4.2: continued
+ log Pi,5- Sn(OH)4° + H2O <=> Sn(OH)5- + H -12.4 4 [1970BAR/KLI] T=298.15K, 1=0.2-2.5 NaOH sol -9.06 > [1970KUR/BAR] T=373 K, I=dil NaOH sol 5 -7.65 [1997AMA/CHI] T= 298.15 K, 1=0.1 0 NaC104 sol -7.97 6 fl998ODA/AMAl T= 298.15 K, 1=0.1 0
+ log Pi,6: Sn(OH)4° + 2H2O <=> Sn( Table 4.3: Thermodynamic data for the hydrolysis of Sn(OH)4° taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log (3m,4m+n Reference Comments I (M) Medium + 4+ log Pi,0- Sn(OH)4° + 4H 4* Sn + 4H2O 0.31 ' [1952LAT] T= 298.15 K, 1=0 0 -1.02 [1955DEL/ZOU] T= 298.15 K, 1=0 0 2.41 2 [197OKUR/BAR] T=373 K, I=dil NaOH 2.61 [1973KLI/BAR] T= 298 K, 1=0 0 1.27 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 1.53 [1979VAS/GLA] T= 298.15 K, 1=3 3 HC1O4 1.73 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 1.94 [1984HOU/KEL] T= 298.15 K( 1=0 0 0.86 3 [1984HOU/KEL] T= 298.15 K,I=0 0 1.26 ] [1985BAB/MAT] T=298.15K, 1=0 0 0.31 1 [1985GAL] T= 298.15 K, 1=0 0 -0.32 [1987BROAVAN] T= 298.15 K, 1=0 0 -0.64 fl988PHI/HAL] T= 298.15 K, 1=0 0 + 3+ log Pi,;: Sn(OH)4° + 3H <=? SnOH + 3H2O 3.10 [1973KLI/BAR] T= 298 K, 1=0 0 1.79 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 2.07 [ 1979V AS/GLA] T= 298.15 K, 1=3 3 HC1O4 2.33 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 1.41 [1984HOU/KEL] T= 298.15 K, 1=0 0 1.43 3 [1984HOU/KEL] T=298.15K, 1=0 0 1.59 [1987BRO/WAN] T= 298.15 K, 1=0 0 -0.18 [1988PHI/HAL] T= 298.15 K, 1=0 0 30 JNC TN8400 99-011 Table 4.3: continued + 2+ log Pi,-24: Sn(OH)/ + 2H tt Sn(OH)2 + 2H20 2.91 [1973KLI/BAR] T= 298 K, 1=0 0 1.88 [1979V AS/GLA] T= 298.15 K, 1=2 2 HC1O4 2.18 [1979VAS/GLA] T= 298.15 K, 1=3 3 HCIO4 2.48 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 1.38 [1984HOU/KEL] T= 298.15 K,I=0 0 1.49 3 [1984HOU/KEL] T= 298.15 K,I=0 0 2.21 [1987BRO/WAN] T= 298.15 K, 1=0 0 -0.95 [1988PHI/HAL] T= 298.15 K, 1=0 0 + log Pu: S?i(OH)4° + H< <=> Sn(OH)3 + H2O 2.03 [1973KLI/BAR] T=298K, 1=0 0 1.31 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 1.22 [1979VAS/GLA] T= 298.15 K, 1=3 3 HC1O4 1.22 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 0.73 l [1980BEN/TEA] T= 298.15 K, 1=0 0 0.72 l [1982WAG/EVA] T= 298.15 K, 1=0 0 -0.52 [1984HOU/KEL] T= 298.15 K, 1=0 0 1.22 3 [1984HOU/KEL] T= 298.15 K, 1=0 0 1.66 [1987BRO/WAN] T= 298.15 K, 1=0 0 -0.24 H988PHI/HAL1 T= 298.15 K, 1=0 0 + log PliS: Sn(OH)4° + H2O <=> Sn(OH)s- + H -2.77 [1987BRO/WAN] 298. 15 K, 1=0 0 -11.18 [1988PHI/HAL] T= 298. 15 K, 1=0 0 -8.85 4 [1992PEA/BER] rp 298. 15 K, 1=0 0 2 log PL6: Sn(OH)4° + 2H2O <=> Sn(OH)6 ' + 2H -20.95 • [1952LAT] T= 298.15 K, 1=0 0 -24.5 [1955DEL/ZOU] T=298.15K, 1=0 0 -21.60 4 [1975KRA] T=298.15K, I=n/a -19.11 [1984HOU/KEL] T= 298.15 K, 1=0 0 -20.96 ' [1985GAL] T= 298.15 K, 1=0 0 -6.57 [1987BRO/WAN] T= 298.15 K, 1=0 0 -22.19 [1988PHI/HAL] T= 298.15 K, 1=0 0 -20.86 4 [1992PEA7BER] T= 298.15 K, 1=0 0 2 log Pl6: Sn(OH)4° <=> SnO3 - + 2H+ + H2O -21.42 [1984HOU/KEL] T= 298.15 K, 1=0 0 -23.15 ' [1985GAL] T= 298.15 K, 1=0 0 -23.05 4 f!992PEA/BERl T= 298.15 K, 1=0 0 1 O calculated with a A,G of -944 kJ/mol for Sn(OH)4° (Section 4.4.3) 2 calculated by [1970KUR/BAR] from thermodynamic data and own measurements. 3 values given in the text; they do not correspond to the values calculated from the A,G0 values given in Table 1 of [1984HOU/KEL], 4 calculated with a log (310 of 0.40 (value from Section 4.4.3). 31 JNC TN8400 99-011 4.2 Solid tin(FV) oxides/hydroxides Fresh precipitates of tin(IV) salt from solutions are amorphous at room temperature. After aging X-ray patterns show broadened reflexes of poorly crystalline SnOj [1963FEI/SCH, 1997AMA/CHI]. [1970KUR/BAR] found from X-ray diffraction measurements that the solid phase precipitated from Sn(IV) solution varied from cassiterite at 300 *C to a structure similar to varlamoffite (Sn(>2 • XH2O) at 100 CC. Their starting material was freshly precipitated Sn(OH)4(s) and equilibrium was attained in 24 hours only at temperatures above 100 *C Both, the crystalline SnO2(cassiterite), and precipitated SnO2(precip) are quite insoluble, while a badly defined, freshly precipitated Sn(OH)4(am) is quite soluble (Table 4.5). This solid, however, transforms to SnC>2(precip) within a month. Only a few experimental measurements of Sn(IV) solubility can be found in the literature. The data used for the calculations of the formation constants of tin(IV) hydroxide/oxide compounds are compiled in Table 4.4. Additional data are given in Table 4.5 and 4.6. Table 4.4: Experimentally determined equilibrium data compiled for the formation of tin(IV) hydroxide/oxide compounds. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 4.14: 'Comments to selected references'. Method: sol = solubility measurements. log K*so Reference Comments I (M) Medium Method log K*so: Sn(OH)40 t=> SnO2(precipitated) + 2H2O 7.46 ' [1997AMA/CHT] T= 298.15 K, 1=0.1 0.1 NaClO. sol log K*so: Sn(OH)40 <=> SnO2(crystalline) + 2H2O 2 8.0 [1997AMA/CHTJ T= 298.15 K, 1=0.01 0.01 NaClO4 sol 1 In [1997 AM A/CHI], this value is reported as log K*°so with the assumption that the activity coefficient of Sn(OH)4° is unity. 2 calculated from experimental data of [1997 AM A/CHI] in this report 32 JNC TN8400 99-011 4.2.1 Freshly precipitated Sn(OH)4(am) The compilation of [1969FEI/SCH] proposes for freshly precipitated Sn(OH)4(am) a log K*So of 2.0 for the reaction Sn(OH)4° <=> Sn(OH)4(freshly precipitated). A similar solubility product is determined by [1970KUR/BAR] at 100 °C. However, Sn(OH)4(am) will turn into a poorly crystalline SnO2(precip) with a constant solubility within a month of aging [1963FEI/SCH, 1997AMA/CHI]. No solubility constant for Sn(OH)4(freshly precipitated) is recommended in the present report. 4.2.2 SnO2(precip) Only a few experimental measurements of the solubility of precipitated, amorphous SnO2 can be found in the literature. Besides the data given in the compilation of [1969FEI/SCH] which lack experimental detail, the solubility of precipitated SnC>2 has been determined recently by [1989BAY/EWA] in cement equilibrated water and by [1997AMA/CHI] and [1998ODA/AMA] 8 in 0.1 M NaC104. [1989BAY/EWA] measured at a pH of 9 a Sn(IV) concentration of 6xlO" M. [1997AMA7CHI] measured in the pH range 2 to 7 a constant solubility of about 3xlO"8 M. [1997AMA/CHI] calculated a log K*So = 7.46. The additional experiments of [1998ODA/AMA] in 0.1 MNaC104 confirmed the solubility data determined by [1997AMA/CHI] (Table 4.4). [1998ODA/AMA] also showed that the presence of chloride and sulfate influences the SnO2 solubility (Table 4.5). The reaction Sn(OH)4° <=» SnO2(precip) + 2H2O is expected to have a Ae ~ 0, because no charged species are involved in the reaction. Thus, the value determined by [1997AMA/CHI] (Table 4.5) may be used to calculate the solubility of precipitated SnO2(precip): Sn(OH)4° o SnO2(precip) + 2H2O log K*°so= 7.46 33 JNC TN8400 99-011 4.2.3 SnO2(cassiterite) Cassiterite crystallizes at elevated temperatures (« 300 °C), while at lower temperature precipitates are less crystalline [1970KUR/BAR]. From the early studies of [1929GRU/LIN] it can be concluded that the solubility of cassiterite in water is smaller than 106 mol Sn(IV) L1. [1970BAR/KLI] measured the solubility of synthetic cassiterite in water, dilute HNO3 and dilute NaOH. They observed a constant, pH-independent (between pH 2 - 11) solubility of 4 x 10 7 mol Sn(IV) L"1. No detection limit for the Sn(FV) is indicated which makes it difficult to decide whether the measured concentrations below pH 11 correspond to the detection limit or are real concentrations. Thus these data were not used to calculate cassiterite solubility More recently, [1997AMA/CHI] determined cassiterite solubility in 0.01 M NaC104. They observed a constant, pH-independent (between pH 2 - 8) solubility of about 9 x 10-9 mol Sn(IV) L"1. From the average of the experimental values, log K*so = 8.0 can be calculated. Extrapolation to 1=0 of the data given by [1997AMA/CHI] assuming a Ae of 0, as no charged species are involved, gives: Sn(OH)4° <=> SnO2(cassiterite) + 2H2O logK*°S0=8.0 This value is in fair agreement with the log K*°So of 8.50 given by [1963FEI/SCH]. At higher temperature, cassiterite solubility was determined in different references (Table 4.5). Extrapolation of the measurements of [1981DAEVSOR] to 298 K gives a log K*So of = 7.5 for cassiterite, a value which is in fair agreement with the log K*so of 8.05 as determined by [1997AMA/CHI] for cassiterite (SnO2(cr)). [1973KLJ/BAR], however, determined in the temperature range 473 - 673 K a higher solubility of cassiterite than [1981DAD/SOR] and [1988BAR/SHA] (Table 4.5). Also at 298 K, the solubility measured by [1973KLI/BAR] is much higher than observed in other references [1926GRU/LIN, 1963FEI/SCH, 1997AMA/CHI] (Table 4.5). For further comments see Section 4.14. 34 JNC TN8400 99-011 4.2.4 Additional data compiled for the formation of tin(FV) hydroxide/oxide compounds Table 4.5: Additional experimentally determined equilibrium data compiled for the formation of tin(TV) hydroxide/oxide compounds. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given at the end of the table, in Section 4.2 and in Section 4.14: 'Comments to selected references'. Method: sol = solubility measurements. log K*so Reference Comments I (M) Medium Method logK*S0:Sn(OH)4° <=> Sn(OH)4(freshly precipitated) 2.00 l [1963FEI/SCH] T= 293 K, I=dil 0 n/a sol 2.21 [1970KUR/BAR] T=373 K, I=dil 0 NaOH sol log K*so: Sn(OH)4° <=> SnO2(precip) + 2H2O 5.00 ' [1963FEI/SCH] T= 293 K, I=dil 0 n/a sol 2 7.46 [1997AMA/CHI] T= 298.15 K, 1=0.1 0 NaC104 sol 2 7.46 [1998ODA/AMA] T= 298.15 K, 1=0.1 0 NaClO4 sol 7.65 3 [1998ODA/AMA] T= 298.15 K, 1=0.53; contains 0.4 M NaCl 0.53 NaCl sol 6.94 3 [1998ODA/AMA] T= 298.15 K, 1=0.57 (0.04 M NaCl) 0.57 NaClCVNaCl sol 6.98 3 [1998ODA/AMA] T= 298.15 K, 1=0.54 (0.004 M NaCl) 0.54 NaClOyNaCl sol 3 6.88 [1998ODA/AMA] T= 298.15 K, 1=0.54 (0.4 M NaC104) 0.53 NaClO4 sol 3 7.04 [1998ODA/AMA] T= 298.15 K, 1=0.43 (0.1 M Na2SO4) 0.43 Na2SO4 sol 3 7.29 [1998ODA/AMA] T= 298.15 K, 1=0.56 (0.01 M Na2SO4 0.56 NaC10VNa2S04 sol 3 7.35 [1998ODA/AMA] T= 298.15 K, 1=0.53 (0.001 M Na,SO4 0.53 NaClOVNajSO, sol logK*S0:Sn(OH)4° <=> SnO2(cassiterite) + 2H2O > 6 [1926GRU/LIN] T= 298 K, I=n/a n/a sol 8.50 ' [1963FEI/SCH] T= 298 K, 1=0 0 n/a sol 6.40 4 [1970BAR/KLI] T= 298.15 K, I=dil 0 no sol 6.44 5 [1973KLI/BAR] T= 298.15 K, 1=0 0 no sol 6.04 [1973KLI/BAR] T= 373 K, 1=0 0 no sol 5.49 [1973KLI/BAR] T= 473 K, 1=0 0 no sol 5.25 [1973KLI/BAR] T=573K,I=O 0 no sol 5.07 [1973KLI/BAR] T= 673 K, 1=0 0.1 NaC104 sol 6.23 [1981DAD/SOR] T=473K, 1=0 0 no sol 5.86 [1981DAD/SOR] T=573K, 1=0 0 no sol 5.60 [1981DAD/SOR] T= 673 K, 1=0 0 no sol 7.49 6 [1981DAD/SOR] T= 298 K, 1=0 0 no sol 8.07 [1988BAR/SHA] T= 573 K, I=dil 0 no sol from unpublished experimental results [1957EGG]. In this report, dissolved species is assumed to be Sn(OH)4°, see also comment in Section 4.14 extrapolated to 1=0 with Davies equation by [1997 AM A/CHI]; [1998ODA/AMA] gives the same value chloride and sulfate form complexes with tin(IV) linear extrapolated from different I, no detection limit indicated see comments in Section 4.14 calculated from values determined at 200-400 °C 35 JNC TN8400 99-011 Table 4.6: Thennodynamic data for the formation of tin(IV) hydroxide/oxide compounds taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K*so Reference Comments I (M) log K*s0: Sn(OH)4° <=> Sn(OH)4(freshly precipitated) 1.33 > [1952LAT] T= 298.15 K, 1=0 0 -1.80 ] [1968SUS/KHO] T= 298 K, I=n/a 3.00 [1975KRA] T= 298.15 K, I=n/a 3.10 [1984HOU/KEL] T= 298.15 K, 1=0 0 1.33 ' [1985BAB/MAT] T= 298.15 K, 1=0 0 1.33 ' [1985GAL1 T= 298.15 K, 1=0 0 log K*S0: Sn(OH)4° <=> SnO2(cassiterite) + 2H2O 8.72 ] [1952LAT] T= 298.15 K, I=n/a 8.75 » [1954COU] T= 298.15 K, I=n/a 6.66 [1955DEL/ZOU] T= 298.15 K, 1=0 0 8.79 l [1963WIC/BLO] T= 298.15 K,I=n/a 8.72 > [1968SUS/KHO] T= 298 K, I=n/a 8.76 ' [1971NAU/RYZ] T= 298.15 K,I=n/a 8.77 ' [1978COD] T= 298.15 K, 1=0 0 8.76 l [1978ROB/HEM2] T= 298.15 K,I=n/a 8.77 l [1979KUB/ALC] T= 298.15 K, I=n/a 8.76 ] [1980BEN/TEA] T=298.15K, I=n/a 8.78 l [1982PAN] T= 298.15 K,I=n/a 8.71 ' [1982WAG/EVA] T= 298.15 K, I=n/a 9.68 [1984HOU/KEL] T= 298.15 K, 1=0 0 8.72 ' [1985BAB/MAT] T= 298.15 K,I=n/a 8.71 ' [1985GAL] T= 298.15 K, I=n/a 14.85 [1987BRO/WAN] T= 298.15 K, 1=0 0 7.58 [1988PHI/HAL] T= 298.15 K,I=n/a 5.40 2 [1992PEA/BER1 T= 298.15 K, 1=0 0 1 calculated with a AfG° of -944 kJ/mol for Sn(OH)4° (Section 4.4.3) 2 calculated in this report with a log P, 0 of 0.40 (value from Section 4.4.3) 4.3 Other tin(IV) complexes and compounds For other tin(IV) species and compounds the amount of experimentally determined data available in the literature is rather limited. From the available data it can be expected that Sn(IV) forms stable complexes and compounds with chloride, fluoride, carbonate, sulfate and sulfide. In Table 4.7 to 4.9 formation constants for tin(IV) species and compounds are compiled. No data are recommended in this report as not sufficient data are available to judge the reliability of these data and to extrapolate formation constants to 1=0. The following tables serve only to give a general idea of the strength of complex formation with tin(IV). 36 JNC TN8400 99-011 Table 4.7: Experimentally determined equilibrium data compiled for the tin(IV) system. These data were not chosen in the present report for the evaluation of recommended stability values because not sufficient data are available to extrapolate the formation constants to 1=0. Method: pot = potentiometry, sol = solubility measurements, se = sulfide selective electrode. logp Reference Comments I (M) Medium Method 4+ 3+ log Pu: Sn + Ci <=> SnCl 3.71 [1978FAT/ROU1 T= 298.15 K, 1=5 HC1O, pot 4+ log Pu: Sn + 2CI- < 6.46 [1978FAT/ROU1 T= 298.15 K, 1=5 HC1O. pot 4+ + log Pu: S?i + 3CI <=> SnCl3 8.78 ri978FAT/ROUl T= 298.15 K, 1=5 HC10d pot 4+ log Ph4: Sn + 4Ci <=? SnCl4° 9.48 ri978FAT/ROU] T= 298.15 K, 1=5 HC1O_ pot 4+ log P!5: Sn + 5CI <=> SnCl5- 11.23 ri978FAT/ROU1 T= 298.15 K, 1=5 HC1O4 pot 4+ 2 log Ph6: Sn + 6CI- <=> SnCl6 - 12.40 ri978FAT/ROUl T= 298.15 K, 1=5 HC1O, pot 2 + log PUJ: Sn(0H)4° + CO3 - + H <=> Sn(OH)3CO3- + H2O 7.71 ' [1971KUE7BAR] T= 298.15 K, 1=0.1-0.5 NaHCO, sol 2+ 2 log Pi,oX- SnO2(aq) + 2H2SO4 <=* Sn(SO4) + SO4 - + 2H2O -1.30 [1955BRU] T= 303 K, I=n/a H2SO4 sol -1.55 ri955BRUl T=291 K, I=n/a sol 2 2 log Pi.o.i-' SnS2(s) + S - <=> SnS3 - 5.31 [1968HSE/REC] T= 298 K, 1=0.1 0.1 NaNO, se log Ks0: Sn(OH)4° + 2HS- + 2 SnS2(s) + 4H2O 36.45 2 [1984KOC/TOP] T= 298.15 K, 1=1-4 HCIO4 sol 1 extrapolated to 1=0 with Debye-HUckel equation by [1971KUR/BAR] 2 calculated with a AfG° of -944 kJ/mol for Sn(OH)4° (Section 4.4.3) 37 JNC TN8400 99-011 Table 4.8: ThermodynamJc data for tin(IV) system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K so Reference Comments I(M) 4+ 0 log Pl,0,4-Sn + 4Ci <=> SnCl4 - 11.52 [1984HOU/KEL] T= 298.15 K, 1=0 0 4+ log Pi,o,6-Sn + 6CI <=> SnClt 0.21 [1984HOU/KEL] T= 298.15 K, 1=0 0 11.23 [1984KOC/TOP] T= 298.15 K, I=n/a + 0 log PIJJ- Sn(OH)4°+ F~+ H 2 log Pi,o,6--Sn(OH)4° + 6F- + 4H"" <=> SnF6 - + 4H2O 12.64 ' [1952LAT] T= 298.15 K, 1=0 0 12.64 ' [1985GAL] T= 298.15 K, 1=0 0 : 2 + 2+ log Pi,o,i Sn(OH)4° + SO4 ' + 4H <=> SnSO4 + 4H2O -2.69 ! [1980BEN/TEA] T= 298.15 K, 1=0 0 -2.70 ' [1982WAG/EVA] T= 298.15 K, I=n/a 6.25 [1988PHI/HAL1 T= 298.15 K, I=n/a + log Pi.0,2- Sn(OH)4° + 2SO/- + 4H <^> Sn(SO4)2° + 4H2O -0.32 l [1982WAG/EVA] T= 298.15 K, I=n/a -1.44 ri988PHI/HALl T= 298.15 K, I=n/a 2 + log Kso: Sn(OH)4° + 2SO4 - + 4H <=> Sn(SO4)2(s) + 4H2O -5.73 ' [1952LAT] T= 298.15 K, I=n/a -15.55 ' [1977BAR/KNA] T= 298.15 K, I=n/a -15.55 ' [1979KUB/ALC] T= 298.15 K, I=n/a -5.73 ' [1985BAB/MAT] T= 298.15 K, I=n/a -5.73 [ [1985GAL] T= 298.15 K, I=n/a -6.84 ri988PHI/HAL] T= 298.15 K, I=n/a + log Kso: Sn(OH)4° + 2HS- + 2H <=> SnS2(s) + 4H2O 30.51 ! [1974MIL] T= 298.15 K, I=n/a 30.51 ' [1977BAR/KNA] T= 298.15 K, I=n/a 30.51 ' [1979KUB/ALC] T= 298.15 K, I=n/a 36.53 ' [1985GAL] T= 298.15 K, I=n/a 35.40 fl988PHI/HALj T= 298.15 K, I=n/a 0 calculated with a Afi of -944 kJ/mol for Sn(OH)4° (Section 4.4.3) 38 JNC TN8400 99-011 Table 4.9: Thermodynamic data for Sn4+ predicted by [1984BRO/WAN]. This table serves only for comparison. log Pu logPii2 log Pi,3 log Pi,4 log Pii5 logPi,6 References Comments 4+ 4 log Pmn: mSn + nCt & SnmCln -" 3.88 7.12 9.84 12.10 13.93 15.34 [1987BROAVAN1 T= 298.15 K, 1=0 4+ 4 n log pmin: mSn + nF <* Snm¥n ' 8.89 17.29 25.34 33.07 40.53 47.72 [1987BROAVAN1 T= 298.15 K, 1=0 4+ 2 4 2 log pmn: mSn + nCO3 - « Snm(CO3)n - " 11.72 21.84 30.62 38.14 44.47 49.65 [1987BROAVAN1 T= 298.15 K, 1=0 4+ 4n log pm_n: mSn + nHCOf <=> Snm(HCO3)n 6.85 12.18 16.22 19.09 20.84 21.50 [1987BROAVAN] T= 298.15 K, 1=0 4+ 4 log pmin: mSn + nNOy <=> Snm(NO3)n -" 1.83 2.16 1.25 -0.81 [1987BROAVAN1 T= 298.15 K, 1=0 4 2 log /3min: mSn * + nHPO4 - « S 12.88 24.86 36.21 47.02 57.35 67.24 [1987BROAVAN] T= 298.15 K, 1=0 4+ 4 log £„,,„: mSn + nH2PO/ <=> SnJH2PO4)n -" 4.75 8.54 11.62 14.10 16.04 17.46 [1987BROAVAN] T= 298.15 K, 1=0 4+ 4 2 log Pm,n: mSn + nSOf <=> Snm(SO4)n - " 4.95 8^61 11.21 12.88 13.65 13.57 [1987BRO/WAN] T= 298.15 K, 1=0 39 JNC TN8400 99-011 4.4 Redox reactions Tin exists in several oxidation states from -IV to +IV in its compounds, but only the +11 and +IV states are important in its aqueous chemistry. The +IV state is particularly stable in natural environments, and cassiterite is the major source of this element in nature [1976BAE/MES, 1985GAL, 1995WIB]. The stannous ion (Sn2+) is easily oxidized. The data used for the calculations of the redox reactions of tin are compiled in Table 4.10. Additional data for the tin redox system are compiled in Table 4.11 and 4.12. These data were not chosen for the calculation of log K° values in this report. Table 4.10: Experimentally determined equilibrium data compiled for redox reactions of tin. Calculated in this report from the respective E° values. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 4.14: 'Comments to selected references'. Method: pol = polarography, pot = potentiometry. logK Reference Comments KM) Medium Method log K: Sn2+ + 2e~ <=> Sn(cr) -5.06 i [1928PRY] T= 298.15 K, 1=0.08 0.08 HC1O4 pot -5.11 ! [1928PRY] T= 298.15 K, 1=0.2 0.2 HC1O4 pot -5.11 1 [1928PRY] T= 298.15 K, 1=0.35 0.35 HC1O4 pot -5.24 2 [1949RIC/POP] T= 298.15 K, 1=1 1 KC1 pot 3 -5.13 [1970BON/TAY] T= 298.15 K, 1=1 1 NaC104 pol 1 calculated in this report from the E° (0.4532, 0.434 and 0.4252 V, respectively) and the respective Sn(II) concentrations as given by [1928PRY] using E°(HgCl/Hg) = 0.2444 V [1996STU/MOR; saturated KC1] 2 calculated by [1949RIC/POP] from data measured with a Calomel electrode in KC1 and KC1O4 medium. Corrected by [1949RIC/POP] for the interaction with Cl, extrapolated to 1=1 with Debye-Huckel approximation. 3 calculated in this report from the E° value (0.374 V) given by [1970BON/TAY] using E°(AgCl/Ag) = 0.2223 V[1996STU/MOR;I=1] 4.4.1 Sn2+/Sn(cr) The redox potential of the reaction Sn2+ + 2e" <=> Sn(cr) has been measured by several authors in acidic media. [1928PRY] and [1949RIC/POP] determined the potential against a calomel electrode in perchlorate solutions and potassium chloride solutions, respectively. [1970BON/TAY] measured in 1 M HCIO4 the potential by polarography against a silver- silver chloride reference electrode. Extrapolation of these values given in Table 4.10 to I = 0 is shown in Figure 4.1 and results in: 40 JNC TN8400 99-011 Sn2+ Sn(cr) logK° = -4.63 E° = -O.137V Sn2+ + 2e~ <=> Sn(cr) 0.5 1 1.5 L. molal 2+ t Figure 4.1: Plot of log K + 4 D vs. Im for the reaction : Sn + 2e- <=> Sn(cr) at 25 C. The straight line shows the result of the linear regression: Ae = -0.25; log K°= -4.63. 4.4.2 Sn2+ISn4+ Sn2+ oxidizes readily to Sn4+ under oxic conditions [1934HUE/TAR, 1979VAS/GLA]. [1934HUE/TAR] measured with a standard hydrogen electrode the redox potential of this reaction in HCl (Table 4.11). Down to a HCl concentration of 0.5 M, [1934HUE/TAR] observed no hydrolysis of the Sn4+ cation present in concentrated acid, while in 0.1 and 0.2 M HCl [1934HUE/TAR] found a strong shift in their potentiometric measurements indicating the hydrolysis of Sn4+. Unfortunately, [1934HUE/TAR, 1979VAS/GLA] had chloride present in their measurements. Sn(II) and probably also Sn(IV) (cf. Table 4.9) has a strong tendency to form chloride complexes. As no data are available for the complex formation between Sn(IV) and chloride, it is not possible to calculate a log K° for the reaction Sn4+ + 2er <=> Sn2+ from these data. [1934HUE/TAR and 1979VAS/GLA] extrapolated from their measurements a log K° ~ 5 for the redox equilibrium Sn4+ + 2e~ <=> Sn2+ (Table 4.11). For the lack of any better data, a tentative log K° value of 5 corresponding to E° = 0.148 V may be used. This value, however, is debatable as the tin chloride complexes are not, or only partly, considered and as the activity corrections made by [1934HUE/TAR and 1979VAS/GLA] are somewhat doubtful. 41 JNC TN8400 99-011 Combining log K° value of 5 with the redox potential of Sn2+ as defined in Section 4.4.1 gives a tentative log K° of 0.37 tor the reaction Sn4+ + 2 H2(g) <=> Sn(cr) + 4H+, corresponding to an E° of 0.005 V. 4.4.3 Sn(OH)4° No measurements of the redox reaction between aqueous Sn(II) and Sn(IV) at neutral or alkaline conditions exist. Reliable data for the protonation of Sn(OH)4° to Sn4+ would be needed for the calculation of the redox equilibria between the species Sn(OH)4° and Sn2+ used as master species in this report. However, such data are not available at present (cf. Section 4.1.1). However, another possibility is to calculate the standard molar Gibbs energy of formation of dissolved species from solubility measurements and the standard molar Gibbs energy of formation of the respective solid. From the AfH° and S° data given in the review of [1989COX/WAG] a AfG° of -515.826 kJ/mol for cassiterite SnO2(cr) is calculated by [1995SEL/BID]. This value is accepted in this report as the standard molar Gibbs energy of formation for cassiterite. Using this value, the solubility of cassiterite as defined in Section 4.2.3, and a AfG° of -237.14 kJ/mol for H2O, a tentative standard molar Gibbs energy of formation for Sn(OH)4° is calculated to be -944.156 kJ/mol resulting in (also these values are tentative): 0 Sn(OH)4° + 2 H2(g) <=> Sn(cr) + 4H2O logK =0.77 E° =0.011 V 2+ or Sn(OH)4°+ 2H++ H2(g) « Sn + 4H2O log K° = 5.40 E° = 0.160 V Only two values for this reaction are reported in the literature (Table 4.12). The value given by [1984HOU/KEL] is based on the hydrolysis data of [1971NAZ/ANT]. Also for the reaction from Sn(OH)4° to Sn4+ a tentative constant can be calculated from the above values: 4+ Sn(OH)4° + 4H+ <=» Sn + 4H2O log K° = 0.40 42 JNC TN8400 99-011 4.4.4 Additional data compiled for the tin redox system Table 4.11: Experimentally determined equilibrium data compiled for the tin redox system. These data were not chosen in the present report for the evaluation of recommended stability values.. Method: pot = potentiometry. logK Reference Comments I (M) Medium Method K: Sn2+ + 2e- <=> Sn(cr) l -4.59 [1928PRY] T= 298.15 K, 1=0.08-0.35 0 HC1O4 pot -4.94 2 [1949RIC/POP] T= 298.15 K, 1=0.1 1 KC1 pot -5.17 2 [1949RIC/POP] T= 298.15 K, 1=1 1 KC1 pot -5.10 2 [1949RIC/POP] T= 298.15 K, 1=1.8 1 KC1 pot log K: Sn2+ <=> Sn4+ + 2e~ -5.21 3 [1934HUE/TAR] T=298.15,1=0.5-2 0 HC1 pot -2.37 4'5 [1934HUE/TAR] T=298.15,1=0.1 0.1 HC1 pot -4.23 4'5 [1934HUE/TAR] T=298.15,1=0.2 0.2 HC1 pot -4.88 4 [1934HUE/TAR] T=298.15,1=0.5 0.5 HC1 pot -4.75 4 [1934HUE/TAR] T=298.15,1=0.8 0.8 HC1 pot -4.67 4 [1934HUEyTAR] T=298.15,1=1.1 1.1 HC1 pot -4.48 4 [1934HUE/TAR] T=298.15,I=2 2 HC1 pot -5.16 6 [1934HUE/TAR] T=298.15,1=0.1 0.1 HC1 pot -5.41 6 [1934HUE/TAR] T=298.15,1=0.2 0.2 HC1 pot -5.67 6 [1934HUE/TAR] T=298.15,1=0.5 0.5 HC1 pot -5,09 6 [1934HUE/TAR] T=298.15,1=0.8 0.8 HC1 pot -4,69 6 [1934HUE/TAR] T=298.15,1=1.1 1.1 HC1 pot -3.85 6 [1934HUE/TAR] T=298.15,1=2 . 2 HC1 pot 7 -6.78 [1979VAS/GLA] T= 298.15 K, 1=2.1 2.116 HC1O4 pot 7 -6.58 [1979V AS/GLA] T= 298.15 K, 1=3.1 3.104 HC1O4 pot 7 -6.58 [1979 V AS/GLA] T= 298.15 K, 1=3.1 3.142 HC1O4 pot 7 -6.38 [1979V AS/GLA] T= 298.15 K, 1=4.1 4.086 HC1O4 pot -5.15 7.8 [1979 V AS/GLA] T= 298.15 K, 1=2-4 0 HC104 pot 1 extrapolated to 1=0 with Debye-Hiickel term by [1928PRY]. Calculated from E°=0.1359 V. 2 calculated from [1949RIC/POP] from data measured with a Calomel electrode. Extrapolated to 1=1 with Debye-Hiickel approximation. 3 extrapolated to 1=0 after correction for activity by [1934HUE/TAR], Calculated in this report from E°=0.154 V. 4 corrected for activity of H+ by [1934HUE/TAR]. Calculated in this report from corrected E° vbalues. 5 the observed potential is a function of tin concentration, indicating hydrolysis of tin. 6 calculated in this report from the E° values and the respective Sn(IV)/Sn(II) concentrations as given by [1934HUE/TAR], Mean of 5 measurements. 1 paper in Russian, no experimental details available. 8 extrapolated to 1=0 with Debye-Huckel term by [1979VAS/GLA]. Calculated in this report from E°=0.1522 V. 43 JNC TN8400 99-011 Table 4.12: Thermodynamic data for tin redox system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK Reference Comments I (M) Medium log K: Sn2+ + 2e- <=> Sn(cr) 4.60 [1941GAR/HEI] T=298.15K, I=n/a 4.61 [1952LAT] T= 298.15,1=0 0 4.60 ' [1955DEL/ZOU] T= 298.15 K, 1=0 0 4.60 l [1975KRA] T= 298.15 K, I=n/a 4.88 [1976VAS/KOK] T= 298.15 K, 1=0 0 HClOa 4.89 [1978COD] T= 298.15,1=0 0 4.88 [1979VAS/GLA] T= 298.15 K, 1=2-4 0 HC1CL 5.24 [1980BEN/TEA] T= 298.15,1=0 0 4.77 [1982WAG/EVA] T= 298.15,1=3 3 4.61 [1984HOU/KEL] T= 298.15 K, 1=0 0 4.76 [1985BAB/MAT] T= 298.15,1=0 0 All [1985GAL] T= 298.15,1=0 0 4.77 [1988PHI/HAL] T= 298.15,1=0 0 4.84 [1989COXAVAG] T= 298.15,1=0 0 4.77 [1992PEA/BER] T= 298.15 K, 1=0 0 log K: Sn2+ <=> Sn4+ + 2e -5.07 [1952LAT] T= 298.15,1=0 0 -5.10 :1 [1955DEL/ZOU] T= 298.15 K, 1=0 0 -5.21 '1 [1975KRA] T= 298.15 K,I=n/a -5.08 [1984HOU/KEL] T= 298.15 K, 1=0 0 -4.27 [1985BAB/MAT] T= 298.15,1=0 0 -5.25 [1985GAL] T= 298.15,1=0 0 -5.05 [1988PHI/HAL] T= 298.15,1=0 0 -5.10 [1992PEA/BER1 T= 298.15 K, 1=0 0 + 2+ log K: Sn(OH)4° + 2H + H2(g) ±=> Sn + 4H2O 7.02 [1984HOU/KEL] T= 298.15 K, 1=0 0 4.41 [1988PHI/HAL] T= 298.15,1=0 0 calculated in this report from E° values. 44 JNC TN8400 99-011 4.5 Hydrolysis of tin(II) Tin(II) is easily oxidized by air to tin(TV) and hydrolyzes readily. In solutions containing more 5 2+ 2+ than 10" M Sn(II), the formation of polymeric Sn(II) species (Sn3(OH)4 and Sn2(OH)2 ) can be observed under acidic conditions [1958TOB, 1976BAE/MES]. Between pH = 4 - 8 the solubility of SnO(s) does not depend on pH, indicating the predominance of the uncharged Sn(OH)2° species. In strong acid or in base tin(II) solubility increases [1976BAE/MES]. The data used for the calculations of the formation constants of tin(II) hydroxide complexes are compiled in Table 4.13. Additional data for the tin(II) hydroxide system which were not chosen for the calculation of log (3° values in this report are compiled in Table 4.14 and 4.15. Table 4.13: Experimentally determined equilibrium data compiled for the hydrolysis of Sn2+, 2+ 2m n + according to the equilibrium: mSn + nH2O <=> Snm(OH)n - + nH . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 4.14: 'Comments to selected references'. Method: sol = solubility measurements. Reference Comments I (M) Medium Method log (3m,n 2+ log fiu: Sn + H20 <=> SnOH+ + H+ -3.92 [1958TOB] T= 298.15 K, 1=3 3 NaC104 pot -3.70 [1976GOB] T= 298.15 K, 1=3 3 NaC104 pot -4.10 [1981PET/MIL] T= 298.15 K, 1=0.1 0.1 NaNO3 pot -3.80 [1981PET/MEL] T= 298.15 K, 1=0.5 0.5 NaNO3 pot -4.10 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNO3 pot -3.77 [1997SAL/FER] T= 298.15 K, 1=3 3 NaCIO, pot 2+ log pli2: Sn + 2H2O o Sn(OH)2° + 2H+ -1.90 [1981PET/MIL] T= 298. 15 K, 1=0.1 0.1 NaNO3 pot -7.90 [1981PET/MIL] T= 298. 15 K, 1=0.5 0.5 NaNO3 pot -7.80 [1981PET/MIL] T= 298. 15 K, 1=1 1 NaNO, pot 45 JNC TN8400 99-011 Table 4.13: continued 2+ + log Pu: Sn + 3H2O ^> Sn(OH)3- + 3H -17.96 [1977MAR] T= 298.15 K, 1=3 3 NaC104 pot -17.5 [1981PET/MIL] T= 298.15 K, 1=0.1 0.1 NaNO3 pot -17.7 [1981PET/MIL] T= 298.15 K, 1=0.5 0.5 NaNO3 pot -17.6 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNO, pot 2+ 2+ log P3 3Sn + 4H2O <=> Sn3(OH)4 + 4H+ -6.11 [1958TOB] T= 298.15 K, 1=3 3 NaClO4 pot -6.81 [1976GOB] T= 298.15 K, 1=3 3 NaClO4 pot -6.87 [1997SAL/FER] T= 298.15 K, 1=3 3 NaCIO, pot [1973JOH/OHT] have concluded from X-ray investigations in 3 M Sn(II) perchlorate solutions that hydrolysis leads to the formation of Sn3(OH)42+ in concentrated Sn(II) solutions: their observation confirm the potentiometric data of [1958TOB], [1958TOB] studied in his careful 2+ + work the hydrolysis of Sn(II) in 3 M NaC104. He determined both the Sn and the H concentration. Titration experiments were conducted in 2.5 - 40 mM of Sn(II) and down to a pH of 1. He interpreted his results with four species: Sn2+, SnOH+, Sn3(OH)42+ and Sn2(OH)22+. His results were recalculated by [1964LIN/TU] with least-squares analysis. [1964LIN/TU] doubted the existence of SnOH+. In place of this species [1964LIN/TU] proposed the existence of Sn2(OH)3+. [1976GOB] and [1997SAL/FER] confirmed later in their work the existence of Sn3(OH)42+ and SnOH+ and rejected the formation of the species Sn2(OH)3+ proposed by [1964LIN/TU]. Many of the earlier measurements of Sn(II) hydrolysis [1939GOR, 1941GAR/HEI, 1952VAN/RHO] neglected the formation of polynuclear tin(II) complexes in their solutions. Thus, their data (given in Table 4.14) were not chosen for the calculation of log (3° values in this report. 46 JNC TN8400 99-011 4.5.1 SnOH+ Under acidic conditions and in solutions containing more than 10"5 M Sn(II), the formation of 2+ 2+ 2+ polymeric Sn(II) species (Sn3(OH)4 or Sn2(OH)2 ), can be expected besides Sn and SnOH+. [1981PET/ML] studied with anodic stripping voltammetry the hydrolysis of tin(II) in 7 0.1 - 1 M NaNO3, NaCl and artificial seawater with a Sn concentration of 10" M and observed only the formation of mononuclear Sn(H) complexes. Further data for the hydrolysis of Sn2+ to SnOH+ are given by [1958TOB] and [1976GOB]. From these experimental results (given in Table 4.13), the following formation constant can be calculated (see Figure 4.2) 2+ Sn + H2O SnOH+ + H+ 1,1 = - 3.75 2+ + + Sn + H2O <^> SnOH + H lm, molal 2+ Figure 4.2: Plot of log (3U + 2 D vs. Im for the reaction : Sn + H2O o SnOH+ + H+ at 25 °C. The straight line shows the result of the linear regression: As = - 0.13; log P°ii = - 3.75. 47 JNC TN8400 99-011 4.5.2 Sn(OH)2° 2+ Formation constants for Sn(OH)2° from Sn have been determined by [1941GAR/HEI] and [1981PET/MIL]. [1941GAR/HEI] measured Sn(II) solubility in 0 - 0.4 M HC1 and NaOH 2+ + assuming only the presence of Sn , SnOH , Sn(OH)2° and Sn(OH)3- at variable I. [1941GAR/HEI] neglected the formation of polymeric species. Under acidic conditions also the formation of tin(II) chloride complexes cannot be excluded. An additional difficulty of the work of [1941GAR/HEI] is that the free acidity was not measured but was calculated indirectly from assumed equilibria. [1981PET/MIL] studied with anodic stripping voltammetry the hydrolysis 7 of tin(II) with an Sn concentration of 10" M. From the data of [1981PETVMIL] in NaNO3 (given in Table 4.13), the following formation constant can be calculated (see Figure 4.3) 2+ Sn + 2H2O <=> Sn(OH)2°+2H+ log =-7.71 2+ + Sn + 2H2O <=> Sn(OH)2° + 2H lmi molal 2+ Figure 4.3: Plot of log (31,2 + 2 D vs. Im for the reaction : Sn + 2H2O <=> Sn(OH)2° + 2H+ at 25 °C. The straight line shows the result of the linear regression: As = - 0.31; log J3°i,2 = -7.71. 48 JNC TN8400 99-011 4.5.3 Sn(OH)3- 2+ Formation constants for Sn(OH)3- from Sn have been determined by [1941GAR/HEI, 1973GAB/SRI, 1977MAR, 1978DIC/LOT1 and 1981PET/MIL]. The value of [1941GAR/HEI] was not chosen for extrapolation to I = 0 (cf. Section 4.5.2: Sn(OH)2°). Extrapolation of the data of [1977MAR] in 3 NaC104 and of [1981PET7MIL] in NaNO3 (given in Table 4.13) gives (see Figure 4.4): 2+ Sn + 3H2O ^ Sn(OH)3-+3H+ log(3°u =-17.54 2+ + Sn + 3H2O <=> Sn(OH)3- + 3H -15 -15.5 -- -16 = -0.12x-17.54 -18 -18.5 -19 -19.5 • -20 L, molal 2+ Figure 4.4: Plot of log pu + 0D vs. Im for the reaction : Sn + 3H2O <=> Sn(OH)3" + 3H+ at 25 °C. The straight line shows the result of the linear regression: Ae = 0.12; log P\3 = - 17.54. [1973GAB/SRI] and [1978DIC/LOT1] determined the redox potential of the reactions Sn(cr) + 3OH" <=> Sn(OH)3- + 2e- as 0.90 and 0.88 V, respectively. From these data log Pi,3 values for 2+ + reactions Sn + 3H2O <^=> Sn(OH)3" + 3H can be calculated (Table 4.14). These values are rather estimates than exact measurements due to possible errors introduced by assuming log P values for the Sn2+/Sn(cr) redox reaction and a log Kw- The selected hydrolysis constants describe the experimental data well. A closer look at the stepwise formation constants shows that they are unequally distributed (Sn2+ to SnOH+, Sn(OH)2°, and Sn(OH)3": 3.75, 3.96 and 9.83). The third stepwise hydrolysis constants is 6 order of magnitude higher than the first two. This unequal distribution of the stepwise hydrolysis constants could be due to a number of reasons, e.g., not all (i.e. polynuclear) relevant species have been considered or due to experimental artefacts (detection limit). 49 JNC TN8400 99-011 2 4.5.4 Sn3(OH)4 + In solution containing more than 10"5 M Sn(II), the formation of polymeric Sn(II) species, can be expected besides Sn2+ and SnOH+ under acidic conditions. [1958TOB, 1976GOB and 2+ 1997SAL/FER] determined formation constants for Sn3(OH)4 in 3 M NaC104 (Table 4.13). Extrapolation of the mean log p3>4 of-6.82 to 1=0 gives (assuming an As of -0.16 as observed 2+ for Pb3(OH)4 , Section 6.1.7): ' 2+ 2+ 3Sn + 4H2O <=> Sn3(OH)4 + 4H+ log (3\4 =-6.51, Ae =-0.16 The value at 1=0 is rather tentative as no data at different I are available. 2+ 4.5.5 Sn2(OH)2 2+ [1928PRY2] and [1958TOB] propose the formation of Sn2(OH)2 in concentrated Sn(II) solutions. More recently, [1976GOB] and [1997SAL/FER] have presented potentiometric data + 2+ which convincingly indicate the formation of SnOH and Sn3(OH)4 , but seem to exclude the 2+ formation of Sn2(OH)2 . 4.5.6 MSn(0H)3+ 2+ [1981 PET/MIL] determined an equilibrium constant for the ion pair formation Sn + 3H2O + + 2+ 2+ 2+ W & MSn(0H)3+ + 3H+, where MP+ is Mg , Ca , or Sr (Table 4.14). As this is the only study of this kind, no log pij 3 is calculated for 1=0. 4.5. 7 Additional equilibrium data compiled for Sn(II) hydrolysis Table 4.14: Additional experimentally determined equilibrium data compiled for the hydrolysis of Sn2+, 2+ 2m n according to the equilibrium: mSn + nH2O *=> Snm(OH)n " + nH*. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text in Section 4.5 and Section 4.14. Method: sol = solubility measurements, sp = spectrophotometry, and pot = potentiometry. log fWn Reference Comments KM) Medium Method 2+ + + log Pij: •!?n + H2O <=> SnOH + H -1.70 ' [1939GOR] T= 298.15 K, 1=0.01-0.5 0 NH4NO3/HC1O4 tit -2.06 > [1941GAR/HEI] T= 298.15 K, 1=0-0.04 var HCl/NaOH sol -1.70 > [1952VAN/RHO] T= 298.15 K, 1=3 3 NaC104 pot <-2.57 [1965MES/IRA] T= 298.15 K, 1=1 1 NaC104 tit -3.70 [1974GOB] T= 298.15 K, 1=3 3 NaC104 pot -3.10 2 [1981PET/MIL] T= 298.15 K, 1=0.5 0.5 NaCl pot -3.10 2 [1981PET/MIL] T= 298.15 K, 1=0.7 0.7 ASW2 pot -2.18 2 [1995DJU/JEL1 T=298.15K, 1=3 3 NaCl pot 50 JNC TN8400 99-011 Table 4.14. continued 2+ log P,,2: Sn + 2H2O <=> Sn(OH)2° + 2W -7.06 ' [1941GAR/HEI] T= 298.15 K, 1=0-0.04 var HCl/NaOH sol -8.20 2 [1981PET/MIL] T= 298.15 K, 1=0.5 0.5 NaCl pot -8.20 2 [1981PET/MIL] T= 298.15 K, 1=0.7 0.7 ASW2 pot 2+ + log P,,.,.- Sn + 3H2O <=> Sn(OH)f + 3H -16.61 1 [1941GAR/HEI] T= 298.15 K, 1=0-0.4 var HCl/NaOH sol -16.59 1 [1941GAR/HEI] T= 298.15 K, 1=0.005 0.005 NaOH sol -15.63 3 [1973GAB/SRI] T= 298.15 K, 1=0.25 0.25 NaOH pot -16.88 4 [1978DIC/LOT1] T= 298.15 K, 1=0.5-5 NaOH pot -17.8 2 [1981PET/MEL] T= 298.15 K, 1=0.5 0.5 NaCl pot -17.2 2 [1981PET/MIL1 T= 298.15 K, 1=0.7 0.7 ASW2 pot 2+ 2+ + log P3,<,: 3Sn + 4H2O <=> Sn3(OH)4 + 4H -6.81 [1974GOB] T= 298.15 K, 1=3 3 NaClO4 pot -2.70 » [1995DJU/JEL1 T= 298.15 K, 1=3 3 NaCl pot 2+ 2+ + log 2Sn + 2H2O <=> Sn2(OH)2 + 2H 5 -2.74 [1928PRY2] T=298.15K,I=dil Sn(C104)2 • tit 5 -2.70 [1928PRY2] T= 298.15 K, I=dil SnCl2 tit -4.10 5-e'< [1928PRY2] T= 298.15 K, 1=0.5 0.5 KC1 tit 5 -4.45 [1958TOB1 T= 298.15 K, 1=3 3 NaC104 pot 2+ 2+ + + log PiiU,: Sn + 3H2O + M <=> MSn(0H)3 + 3H 7 -16.20 [1981PET/MIL] T= 298.15 K, 1=0.5 0.5 NaCl pot 1 polynuclear complexes not considered 2 formation of tin chloride complexes probable; ASW = Artificial seawater 3 estimated from potentiometric data assuming for 1=0.25 a log Kw of -13.76, log P (Sn2+/Sn(cr)) of 5.11 (see Section 4.4: Redox reactions) 4 extrapolated to 1=0 using Pitzer's approach by [1978DIC/LOT1]. Estimated in this report from potentiometric data assuming a log Kw of -14.0, log K (Sn2+/Sn(cr)) of 4.63 (see Section 4.4: Redox reactions) 5 this species does probably not exist, cf. Section 4.5.5 6 formation of Sn(II) chloride complexes 7 M2+ = Mg2+, Ca2+, or Sr2+ Table 4.15: Thermodynamic data compiled for the hydrolysis of Sn2+ taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. Reference Comments I (M) Medium 2+ + + log Sn + H2O <=> SnOH + H -3.40 [1976BAE/MES] T= 298.15 K, 1=0 0 -3.78 [1976SMI/MAR] T= 298.15 K, 1=3 3 -3.40 fl981BAE/MESl T= 298.15 K, I=n/a 51 JNC TN8400 99-011 Table 4.15: continued -1.67 [1982WAG/EVA] T= 298.15 K, 1=0 0 -3.70 [1982SMI/MAR] T= 298.15 K, 1=3 3 -3.88 [1984HOU/KEL] T= 298.15 K, 1=0 0 -2.10 [1985BAB/MAT] T= 298.15 K, 1=0 0 -1.54 [1985GAL] T= 298.15 K, 1=0 0 -3.67 [1987BRO/WAN] T= 298.15 K, 1=0 0 -3.41 [1988PHI/HAL] T= 298.15 K, 1=0 0 -3.40 [1989SMI/MAR] T= 298.15 K, 1=0 0 -3.78 [1989SMI/MAR] T= 298.15 K, 1=3 3 -3.64 [1989SMI/MAR] T= 293.15 K, 1=0.5 0.5 -1.55 [1992PEA/BER1 T= 298.15 K, 1=0 0 2+ + log Pi.2: Sn + 2H2O <=> Sn(OH)2° + 2H -7.1 [1963FEI/SCH] T= 293 K, 1=0 0 -7.06 [1976BAE/MES] T= 298.15 K, 1=0 0 -7.00 [1984HOU/KEL] T= 298.15 K, 1=0 0 -6.64 [1985BAB/MAT] T= 298.15 K, 1=0 0 -8.64 [1987BROAVAN] T= 298.15 K, 1=0 0 -7.07 [1988PHI/HAL] T= 298.15 K, 1=0 0 -7.10 [1989SMI/MAR] T= 298.15 K, 1=0 0 -7.58 [1989SMI/MAR] T= 293.15 K, 1=0.5 0.5 -7.06 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ + log fiu: Sn + 3H2O <=> Sn(OH)y + 3H 15.88 [1952LAT] T= 298.15 K,I=n/a 15.88 [1955DEUZOU] T= 298.15 K, 1=0 0 16.7 [1963FEI/SCH] T= 293 K, 1=0 0 16.6 [1976BAE/MES] T= 298.15 K, 1=0 0 17.96 [1982SMI/MAR] T= 298.15 K, 1=3 3 18.71 [1984HOU/KEL] T= 298.15 K, 1=0 0 16.16 [1985BAB/MAT] T= 298.15 K, 1=0 0 16.03 [1985GAL] T= 298.15 K, 1=0 0 14.78 [1987BRO/WAN] T= 298.15 K, 1=0 0 16.64 [1988PHI/HAL] T= 298.15 K, 1=0 0 16.60 [1989SM1/MAR] T= 298.15 K, 1=0 0 17.94 [1989SMI/MAR] T= 298.15 K, 1=3 3 18.54 [1989SMI/MAR] ' T= 293.15 K, 1=0.5 0.5 16.61 [1992PEA/BER1 T= 298.15 K, 1=0 0 2+ + log Pu: Sn + 2H2O <=> SnOOH- + 3H -15.8 [1975KRA] T= 298.15 K, I=n/a -15.82 [1984HOU/KEL] T= 298.15 K, 1=0 0 -16.05 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ 2 + log p14; Sn + 4H2O <=> Sn(OH)4 - + 4H -22.05 [1987BROAVAN] T= 298.15 K, 1=0 0 -21.98 [1988PHI/HAL] T= 298.15 K, 1=0 0_ 52 JNC TN8400 99-011 Table 4.15: continued 2+ 2+ + log fh - 2Sn + 2H2O <=> Sn2(OH)2 + 2H -4.58 [1964LIN/TU] T= 298.15 K, 1=3 3 NaClO, -4.77 [1976BAE/MES] T= 298.15 K, 1=0 0 -4.46 [1976SMI/MAR] T= 298.15 K, 1=3 3 -4.46 [1982SMI/MAR] T= 298.15 K, 1=3 3 -4.39 [1984HOU/KEL] T= 298.15 K, 1=0 0 -4.79 [1988PHI/HAL] T= 298.15 K, 1=0 0 -4.80 [1989SMI/MAR] T= 298.15 K, 1=0 0 -4.46 [1989SMI/MAR] T= 298.15 K, 1=3 3 -4.77 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ + + log (32J: 2Sn + 3H2O Sn2(OH)3 + 3H -6.66 [1964LIN/TU] T= 298.15 K, 1=3 3 NaC104 -6.77 [1984HOU/KEL] T= 298.15 K, 1=0 0 2+ 2 + log P26: 2Sn + 3H2O <=> Sn2O3 - + 6H -30.36 fl984HOU/KEL1 T= 298.15 K, 1=0 2+ 2+ + log Ps,4- 3Sn + 4H2O <=> Sn3(OH)4 + 4H -6.85 [1964LIN/TU] T= 298.15 K, 1=3 3 NaC104 -6.88 [1976BAE/MES] T= 298.15 K, 1=0 0 -6.79 [1976SMI/MAR] T= 298.15 K, 1=3 3 -6.79 [1982SMI/MAR] T= 298.15 K, 1=3 3 -6.65 [1984HOU/KEL] T= 298.15 K, 1=0 0 -6.26 [1987BRO/WAN] T= 298.15 K, 1=0 0 -6.93 [1988PHI/HAL] T= 298.15 K, 1=0 0 -6.90 [1989SMI/MAR] T= 298.15 K, 1=0 0 -4.47 [1989SMI/MAR] T= 298.15 K, 1=3 3 -6.88 T1992PEA/BER1 T= 298.15 K, 1=0 0 2+ + + log P3,5: 3Sn + 5H2O <=> Sn3(OH)s + 5H -8.41 f!987BROAVANl T= 298.15 K, 1=0 2+ + + log $4,4: 4Sn + 4H2O <=* Sn4(OH)/ + 4H -3.07 [1987BROAVAN1 T= 298.15 K, 1=0 2+ 4+ + log (l6i8: 6Sn + 8H2O <=> Sn6(OH)8 + 8H -7.81 [1987BROAVAN1 T= 298.15 K, 1=0 53 JNC TN8400 99-011 4.6 Solid tin(II) oxide/hydroxide SnO(cr) is more soluble than SnO2(precip). Dissolved tin(II) species precipitate from solution as Sn(OH)2(precip). The hydroxide is unstable with respect to the oxide, SnO(cr) and may undergo progressive dehydration to SnO(cr) or oxidation to SnO2(precip) [1963SHA/DAV, 1984HOU/KEL and references therein]. [1906GOL/ECK] observed that the white Sn(OH)2(precip) precipitates changed to black SnO(cr) in aqueous solutions and that with this change the amount of dissolved Sn(II) decreased. In contrast, [1955DEL/ZOU] stated that dehydration of Sn(OH)2(precip) to SnO(cr) is complete only at T « 100 °C. [1955DEL/ZOU] give no experimental evidence for their statement. Both, the crystalline SnO(cr), and precipitated Sn(OH)2(precip) are quite soluble (Table 4.16). Only few experimental measurements of Sn(II) solubility can be found in the literature. The data used for the calculations of the formation constants of tin(II) hydroxide/oxide compounds are compiled in Table 4.16. Additional data are given in Tables 4.17 and 4.18. Table 4.16: Experimentally determined equilibrium data compiled for the formation of tin(II) hydroxide/oxide compounds. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 4.14: 'Comments to selected references'. Method: pot = potentiometry, sol = solubility measurements. log K so Reference Comments I (M) Medium Method 2+ log K*so: Sn + H2O <=> SnO(cr) + 2H+ -2.41 i [1941GAR/HEI] T= 298.15 K, I=dil 0 water sol -2.93 [1977MAR] T= 298.15 K, 1=3 3 NaClO. pot 2+ log K*so: Sn + 2H2O <=> Sn(OH)2(precip) + 2H+ -2.84 i [1906GOL/ECK] T= 298.15 K, I=dil 0 water sol -2.79 [1922PRY] T= 298.15 K, I=dil 0 water sol recalculated from solubility in water and log (3,2 = -7.71; see also text 54 JNC TN8400 99-011 4.6.1 SnO(cr) and Sn(OH)2(precip) In water [1941GAR/HEI] determined the solubility of SnO(cr) to be 5 x 10-6 M Sn(TI) and 5 reported for Sn(OH)2(precip) a solubility of 1.35 x 10" M (determined by [1906GOL/ECK]). 2+ From these values and a log (31>2 of -7.71 (see Section 4.5.2) for the reaction Sn + 2H2O <=> Sn(OH)2° + 2H+, the following solubility can be calculated (Figure 4.5) Sn2+ + H O 2 SnO(cr) + 2H+ log K*°so = -2.41 For Sn(OH)2(precip) the solubility may be best defined by the mean of the measurements of [1922PRY2] and [1906GOL/ECK] (Table 4.16): 2+ Sn + 2H2O o Sn(OH)2(precip) + 2H+ log K*°so = -2.82 The values calculated by [1941GAR/HEI] themselves are different (Table 4.15 and 4.16) as [1941GAR/HEI] used a different log pu value. 2+ + Sn + H2O *=> SnO(s) + 2H lm, molal 2+ + Figure 4.5: Plot of log K*so + 2 D vs. Im for the reaction : Sn + H2O <=> SnO(s) + H at 25 °C. The straight line shows the result of the 'linear regression': Ae = -0.01; log K*°so = -2.41. Calculated from data given in Table 4.16. 55 JNC TN8400 99-011 4.6.2 Additional equilibrium data compiled for tin(ll) hydroxide/oxide compounds Table 4.17: Additional experimentally determined equilibrium data compiled for the formation of tin(II) hydroxide/oxide compounds. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 4.14: 'Comments to selected references'. Method: sol = solubility measurements. log K*so Reference Comments I (M) Medium Method 2+ + log K*s0: Sn + H2O <=> SnO(s) + 2H -1.76 ' [1941GAR/HEI] T= 298.15 K, I=dil 0 water sol -2.30 H965MES/IRA1 T= 298.15 K, 1=1 1 NaC104 sol 2+ + log K*s0: Sn + 2H2O <=> Sn(OH)2(precipitated) + 2H -2.19 ' [1941GAR/HEI] T= 298.15 K,I=dil water 2 0.33 [1995DJU/JEL1 T= 298.15 K, 1=3 NaCl pot 1 formation of polymeric species neglected; 2 formation of tin chloride complexes probable Table 4.18: Thermodynamic data compiled for the formation of tin(II) hydroxide/oxide compounds taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K*so Reference Comments KM) 2+ + log K*so: Sn + H2O <=> SnO(s) •f 2H -1.12 [1937HOA] T= 298.15 K,I=n/a -1.08 [1952LAT] T=298.15 K, I=n/a -1.17 ' [1954COU] T=298.15K, I=n/a -1.07 [1955DEL7ZOU] T= 298.15 K, 1=0 0 -1.80 [1963FEI/SCH] T= 293 K, 1=0 0 -1.17 l [1963WIC/BLO] T= 298.15 K, I=n/a -1.10 ' [1964HIR] T=298.15K, I=n/a -1.17 * [1971NAU/RYZ] T= 298.15 K,I=n/a -1.76 [1976BAE/MES], T= 298.15 K, 1=0 0 -1.80 [1976SMI/MAR] T= 298.15 K, 1=0 0 -1.38 [1978COD] T=298.15K, I=n/a -101.4 ' [1979KUB/ALC] T= 298.15 K,I=n/a -1.76 [1981BAE/MES] T= 298.15 K, 1=0 -1.12 ' [1982PAN] T=298.15K, I=n/a -1.30 [1982WAG/EVA] T= 298.15 K,I=n/a -1.04 [1984HOU/KEL] T= 298.15 K, 1=0 0 -1.30 [1985BAB/MAT] T= 298.15 K,I=n/a -1.29 [1985GAL] T=298.15K, I=n/a -1.12 ' [1985JAC/HEL2] T=298.15K, I=n/a -2.17 [1986KOV/RYZ] T=298.15K, I=n/a -1.31 [1988PHL«AL1 T= 298.15 K, I=n/a 56 JNC TN8400 99-011 Table 4.17: continued -2.25 [1989COX/WAG] T= 298.15 K, I=n/a -1.76 ri992PEA/BER] T= 298.15 K, 1=0 2+ + log K*so- Sn + 2H2O <=> Sn(OH)2(precipitated) + 2H -1.51 [1952LAT] T=298.15 K, I=n/a -1.5 [1955DEL/ZOU] T= 298.15 K, 1=0 -2.21 [1980BEN/TEA] T= 298.15 K, I=n/a -1.43 [1984HOU/KEL] T= 298.15 K, 1=0 -1.73 [1985BAB/MAT] T= 298.15 K, I=n/a -1.67 [1985GAL] T=298.15 K, I=n/a -2.42 [1987BROAVAN] T= 298.15 K,I=0 -1.75 [1988PHI/HAL] T= 298.15 K, I=n7a calculated with a AfG° of -26.42 kJ/mol for Sn2+ (Section 4.4.1) 57 JNC TN8400 99-011 4.7 Tin(II) chloride system Tin(II) forms complexes and compounds with chloride. Also the existence of mixed tin(II) chloride hydroxide complexes and compounds is reported. Experimentally determined data are given in Table 4.19. Further thermodynamic data are compiled in Table 4.20 and 4.21. Table 4.19: Experimentally determined equilibrium data compiled for the tin(II) chloride system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 4.14: 'Comments to selected references'. Method: pot = potentiometry, sol = solubility measurements. ) Medium log Pm,n Reference Comments KM: Method log Pif- Sn2+ + Cl- & SnCl+ 1.11 [1951DUK/PIN] T= 298.15 K, 1=2 2 HC1O4 pot 1,06 1 [1961RAB/MOO] T= 298.15 K, 1=2 2 cio; pot 2 1.14 [1961RAB/MOO] T= 298.15 K, 1=3 3 NaC104 pot 1.45 [1962HAJ/ZOL] T= 298.15 K, 1=4 4 H2SO4/HC1 sol 1,09 [1975FED/BOL] T= 298.15 K, 1=0.5 0.5 NaC104 pot 1.02 [1975FED/BOL] T= 298.15 K, 1=1 1 NaC104 pot 1.18 [1975FED/BOL] T= 298.15 K, 1=3 3 NaClO4 pot 1.34 [1975FED/BOL] T= 298.15 K, 1=4 4 NaC104 pot 1.80 [1975FED/BOL] T= 298.15 K, 1=6 6 NaC104 pot 1.08 [1976SAM/LYA] T= 298.15 K, 1=1 1 NaC104 pot 0.73 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNO, pot 2+ log Pi,2: Sn + 2CI- <=> SnCl2° 1.72 1 [1961RAB/MOO] T= 298.15 K, 1=2 2 cio; pot 1.70 2 [1961RAB/MOO] T= 298.15 K, 1=3 3 NaC104 pot 2.35 [1962HAI/ZOL] T= 298.15 K, 1=4 4 H2SO4/HC1 sol 1.36 [1975FED/BOL] T= 298.15 K, 1=0.5 0.5 NaClO4 pot 1.13 [1975FED/BOL] T= 298.15 K, 1=1 1 NaC104 pot 1.78 [1975FED/BOL] T= 298.15 K, 1=3 3 NaClO4 pot 2.12 [1975FED/BOL] T= 298.15 K, 1=4 4 NaClO4 pot 3.04 [1975FED/BOL] T= 298.15 K, 1=6 6 NaC104 pot 1.85 [1976SAM/LYA] T= 298.15 K, 1=1 1 NaC104 pot 1.08 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNO, pot 58 JNC TN8400 99-011 Table 4.19: continued 2+ log PJi3: Sn + 3CI- a SnClf 1.50 ] [1961RAB/MOO] T= 298.15 K, 1=2 2 cio; pot 1.66 2 [1961RAB/MOO] T= 298.15 K, 1=3 3 NaC104 pot 2.46 [1962HAI/ZOL] T= 298.15 K, 1=4 4 H2SO4/HC1 sol 1.65 [1975FED/BOL] T= 298.15 K, 1=3 3 NaC104 pot 2.12 [1975FED/BOL] T= 298.15 K, 1=4 4 NaC104 pot 3.30 [1975FED/BOL] T= 298.15 K, 1=6 6 NaClO4 pot log Pi,4: Sn2+ + 4CI- <=> SnCUr7- 2.31 [1962HAI/ZOL] T= 298.15 K, 1=4 4 H7SO4/HC1 sol 2+ + log Pi,1,1: Sn + H2O + Cl- 4=* SnOHCl°+ H 3 -2.77 [1952VAN/RHO] T= 298.15 K, 1=3 3 NaC104 pot -2.90 [1981PET/MIL] T= 298.15 K, 1=0.5 0.5 NaCl pot 2+ log K*so: Sn + H2O + Cl- <=t> SnOHCl(s) + H+ 1.38 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.169 Cl- sol 1.37 4 [1930RAN/MUR] T= 298.15 K,I=n/a 0.17 Cl- sol 1.38 4 [1930RAN/MUR] T= 298.15 K,I=n/a 0.17 Cl- sol 1.38 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.179 Cl- sol 1.79 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.058 Cl- sol 1.70 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.068 Cl- sol 2.25 4 [1930RAN/MUR] T= 298.15 K,I=n/a 0.018 Cl- sol 2.24 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.015 Cl- sol 2.25 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.015 Cl- sol 2.25 4 [1930RAN/MUR] T= 298.15 K,I=n/a 0.014 Cl- sol 2.25 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.014 Cl- sol 2.24 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.014 Cl- sol 1.82 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.088 Cl- sol 1.67 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.082 Cl- sol 2.00 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.039 Cl- sol 1.47 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.13 Cl- sol 1.87 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.045 Cl- sol 1.80 4 [1930RAN/MUR] T= 298.15 K,I=n/a 0.048 Cl- sol 1.78 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.048 Cl- sol 1.77 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.053 Cl- sol 1.85 4 [1930RAN/MUR] T= 298.15 K, I=n/a 0.048 Cl- sol 1 recalculation of measurements from [1950DUK/COU] 2 recalculation of measurements from [1952VAN/RHO] 3 + 0 calculated in this report from a log KiiU of 1.04 for SnOH + CY <=> SnOHCl [1952VAN/RHO] and a log 2+ + (3U of -3.81 for Sn + H2O <=> SnOH+ + H at I = 3 (see Table 4.13) 4 recalculated in this report from measurements of [1930RAN7MUR]. Values have been corrected for the formation of Sn chloride complexes. Uncorrected data are given in Table 4.20 59 JNC TN8400 99-011 + 2 4.7.1 SnCl , SnCl2°, SnClf, and SnCl4 - A number of equilibrium constants for the formation of tin(II) chloride complexes can be found in the literature. As early as 1928, [1928PRY] determined with potentiometric measurements formation constants for tin(II) chloride complexes and corrected these values to 1=0. [1961RAB/MOO] recalculated the experiments of [1950DUK/COU] and [1952VANRHO] with a non-linear least square fit. From the data given in Table 4.19 the following formation constants for the tin chloride complexes can be derived with the SIT equation (see also Figure 4.6, 4.7 and 4.8): 2+ + Sn HhCl- <=» SnCl log P°U = 1.65 2+ Sn Hh2Cl- « SnCl2° log P°i,2 = 2.31 2+ Sn Hh3Cl" «=> SnCl3- log P°i,3 = 2.09 These constants extrapolated with the SIT equation are comparable with the constants determined by [1928PRY] for I = 0 (Table 4.20). No equilibrium constant is recommended in this report for the reaction Sn2+ + 4C1" <=> SnCU2" as this constant was only determined at I = 4 (Table 4.19). The Ae values calculated agree well with the Ae values calculated for Pb(II) chloride complexes (see Section 6.3). Sn2* + CI- ^» SnCI+ c 0 4.5 4 Q3.5 + 3 ci2-5 g) 2 1.5 0 1 y = 0.14x + 1.65 0.5 n 0 2 4 6 8 10 lm, molal 2+ + Figure 4.6: Plot of log Pu + 4 D vs. Im for the reaction : Sn + Cl' <=* SnCl at 25 °C. The straight line shows the result of the linear regression: Ae = -0.14; log (3\i = 1.65. Calculated from data given in Table 4.19. 60 JNC TN8400 99-011 2+ Sn + 2CI- «^> SnCI2° c 0 4.5 4 Q 3.5 CO + 3 O O^^o CM ,-T 2.5 CO. 2 _g 1.5 y = 0.26X + 2.31 1 0.5 • n • 1 1 H 1 4 6 8 10 lm> molal 2+ Figure 4.7: Plot of log (3]>2 + 6 D vs. Im for the reaction : Sn + 2C1- <=> SnCl2° at 25 °C. The straight line shows the result of the linear regression: Ae = -0.26; log (3°ii2 = 2.31. Calculated from data given in Table 4.19. Sn2+ + 3CI" y = 0.29x + 2.09 4 6 10 lm> molal 2+ Figure 4.8: Plot of log (3li3 + 6 D vs. Im for the reaction : Sn + 3C1" «=> SnCl3- at 25 °C. The straight line shows the result of the linear regression: AE = -0.29; log P\3 = 2.09. Calculated from data given in Table 4.19. 61 JNC TN8400 99 -Oil 4.7.2 SnOHCl0 A mixed tin hydroxide chloride complex can be formed. [1952VAN/RHO] gives for the reaction + 0 SnOH + Cl- « SnOHCl a log KU;i of 1.04 and [1981PET/MIL] gives a log Ku,i of 1.14 2+ 0 for the same reaction and a log J5ITI,I of -2.9 for the reaction Sn + H2O + Cl" <=> SnOHCl + H+. Extrapolation to I = with the SIT equation in Figure 4.9 gives 2+ 0 + Sn + H2O + Cl- SnOHCl + H (3°i,i,i = -2.27 2+ + Sn +H2O+CI" <=> SnOHCI°+H 0 - -0.5 - -1 Q -1.5 + -2 -——-—"""^ —&-——•"" -2.5 oa y = 0.12x -2.27 -3 -3.5 .4, -4.5 - -5 1 • 1 h 1 2 mola! 2+ 0 Figure 4.9: Plot of log pu>1 + 4 D vs. Im for the reaction : Sn + H2O + Cl- <=> SnOHCl + H+ at 25 °C. The straight line shows the result of the 'linear regression': As = - O 0.12; log (3 1)U = -2.27. Calculated from data given in Table 4.19. 62 JNC TN8400 99-011 4.7.3 SnCl2(s) Direct measurements of tin chloride solubility are not available. Based on the data compiled in Table 4.21 it can be concluded that SnCl2(s) is easily soluble. Under environmental conditions, SnOHCl(s) is thermodynamically more stable. 4.7.4 SnOHCl(s) [1930RAN/MUR] determined the solubility of SnOHCl(s) in water (pH range form 1.1 to 2.2). Their experimental results are given in Table 4.19. [1930RAN/MUR] extrapolated their results to 1=0 with the Debye-Hiickel term, resulting in a log K*°So of 2.75 (Table 4.20). In this report, the values reported by [1930RAN/MUR] are corrected for the formation of Sn(II) chloride complexes. Extrapolation with the SIT equation to 1=0 gives (Figure 4.10): 2+ + Sn + H2O SnOHCl(s) + H log K*°so = 2.42 2+ + Sn +H2O+CI- & SnOHCI(s)+H 3 -I 2.5 2 Q 1.5 + 1 y = -3.28x + 2.42 o 0) 0.5 0- lo g -0.5 -1 -1.5 C) 0.1 0.2 0.3 0.4 0 lm, molal 2+ Figure 4.10: Plot of log K*so + 4 D vs. Im for the reaction : Sn + H2O + Cl" <=> SnOHCl(s) + H+ at 25 °C. The straight line shows the result of the linear regression: Ae = 3.28; log K*°so = 2.42. Calculated from the data given in Table 4.19. 63 JNC TN8400 99-011 4. 7.5 Additional equilibrium data compiled for the tin(II) chloride system Table 4.20: Additional experimentally determined equilibrium data compiled for the tin(ll) chloride hydroxide system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given at the end of this table. Method: pot = potentiometry, sol = solubility measurements. Reference Comments I (M) Medium Method 2+ log Pu: Sn + Cl- <=> SnCl" 2 1.51 »- [1928PRY] T= 298.15 K, 1=0.08-0.35 0 HC1, KC1 pot 1.85 2 [1949RIC/POP] T= 298.15 K, 1=0.1-4 0 KC1 pot 3 1.05 [1950DUK/COU] T= 298.15 K, 1=2 2 NaC104/HC104 pot 3 1.15 [1952VAN/RHO] T= 298.15 K, 1=3 3 NaC104 pot 0.704 [1962HAI/ZOL] T= 298.15 K, 1=2-6 var HC1 sol 2 1.87 [1975EED/BOL1 T= 298.15 K, 1=0 0 NaC104 pot 2+ log P,t2- Sn + 2CI <=> SnCl2° 2.25 I- [1928PRY] T= 298.15 K, 1=0.08-0.35 0 HC1, KC1 pot 2.31 2 [1949RIC/POP] T= 298.15 K, 1=0.1-4 0 KC1 pot 3 1.76 [1950DUK/COU] T= 298.15 K, 1=2 2 NaC104/HC104 pot 3 1.70 [1952VAN/RHO] T= 298.15 K, 1=3 3 NaC104 pot 0.78 4 [1962HAI/ZOL] T= 298.15 K, 1=2-6 var HC1 sol 2 2.38 [1975FED/BOL] T= 298.15 K, 1=0 0 NaC104 pot 2+ log Pu: Sn + 3d <=P SnClf 1 2 2.03 - [1928PRY] T= 298.15 K, 1=0.08-0.35 0 HC1, KC1 pot 1.94 2 [1949RIC/POP] T= 298.15 K, 1=0.1-4 0 KC1 pot 3 1.14 [1950DUK/COU] T= 298.15 K, 1=2 2 NaC104/HC104 pot 3 1.68 [ 1952V AN/RHO] T= 298.15 K, 1=3 3 NaC104 pot 4 0.38 [1962HAI^OL] T= 298.15 K, 1=2-6 var HC1 sol 2 1.93 [1975FED/BOL1 T= 298.15 K, 1=0 0 NaCIO, pot 2+ 2 log Pl4: Sn + 4CI <=> SnCl4 - 1.50 >'2 [1928PRY] T= 298.15 K, 1=0.08-0.35 0 HC1, KC1 pot 2.00 2 [1949RIC/POP] T= 298.15 K, 1=0.1-4 0 KC1 pot 3 1.14 [1950DUK/COU] T= 298.15 K, 1=2 2 NaC104/HC104 pot 2.31 [1962HAI/ZOL] T= 298.15 K, 1=4 4 H2SO4/HC1 sol -0.24 4 [1962HAI/ZOL] T=298.15K, 1=2-6 var HC1 sol 2+ + log P,,u .• Sn + H2O + Cl- <=>SnOHCl°+ H 4 -0.66 [1952V AN/RHO] T=298.15K, 1=3 3 NaC104 pot -1.35 5 [1986KOV/RYZ] T= 298.15 K, 1=0 0 water sol 64 JNC TN8400 99-011 Table 4.20: continued /+ + log K'so: i5/2 + H2O + Ci <=> SnOHCl(s) + H 1.29 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.169 Cl- sol 1.27 7 [1930RAN/MUR] T= 298.15 K.I^n/a 0.17 Cl- sol 1.28 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.17 Cl- sol 1.28 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.179 Cl- sol 1.76 7 [1930RAN/MUR] T= 298.15 K, I=n/a 0.058 Cl- sol 1.66 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.068 Cl- sol 7 2.24 [1930RAN/MUR] T= 298.15 K,I=n/a 0.018 Cl- sol 2.23 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.015 Cl- sol 2.23 7 [1930RAN/MUR] T= 298.15 K, I=n/a 0.015 Cl- sol 2.24 7 [1930RAN/MUR] T= 298.15 K, I=n/a 0.014 Cl- sol 2.24 7 [193ORAN/MUR] T= 298.15 K, I=n/a 0.014 CI- sol 2.23 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.014 Cl- sol 1.777 [1930RAN/MUR] T= 298.15 K, I=n/a 0.088 Cl- sol 1.62 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.082 Cl- sol 1.97 7 [1930RAN/MUR] T=298.15K, I=n/a 0.039 Cl- sol 1.39 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.13 Cl- sol 7 1.84 [1930RAN/MUR] T= 298.15 K,I=n/a 0.045 Cl- sol 1.77 7 [1930RAN/MUR] T= 298.15 K, I=n/a 0.048 Cl- sol 1.75 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.048 Cl- sol 1.74 7 [193ORAN/MUR] T= 298.15 K, I=n/a 0.053 Cl- sol 1.82 7 [1930RAN/MUR] T= 298.15 K,I=n/a 0.048 Cl- sol 2.75 2 [1930RAN/MUR1 T= 298.15 K, I=n/a 0 Cl- sol 1 measured with Calomel electrode 2 extrapolated to 1=0 by the respective authors with Debye-Huckel equation 3 these values were recalculated by [1961RAB/MOO], see Table 4.19 4 I not constant 5 calculated by [1952VAN/RHO] without considering polymeric tin(II) hydroxide species: Corrected value in Table 4.19 6 extrapolated from measurement at 500 °C 7 formation of Sn chloride complexes not considered. Constants corrected for the formation of Sn chloride complexes are given in Table 4.18 Table 4.21: Thermodynamic data compiled for the tin(II) chloride hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. Reference Comments I (M) Medium 2+ + log 0U: Sn + Ci SnCl 1.14 [1967AHR] T= 298.15 K, 1=3 3 1.51 [1976SMI/MAR] T= 298.15 K, 1=0 0 1.08 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.17 [1976SMI/MAR] T= 298.15 K, 1=3 3 1.45 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.64 [1976V AS/KOK] T=298.15K, 1=0 0 HC1O4 1.00 [1976V AS/KOK] T= 298.15 K, 1=0.5 0.5 HC1O4 1.01 [1976V AS/KOK] T= 298.15 K, 1=1 1 HC1O4 1.06 [1976V AS/KOK] T= 298.15 K, 1=2 2 HC1O4 1.16 [1976 V AS/KOK] T= 298.15 K, 1=3 3 HC1O4 1.06 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 65 JNC TN8400 99-011 Table 4.21: continued 1.16 [ 1979V AS/GLA] T= 298.15 K, 1=3 3 HC1O4 1.28 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 1.64 [1982SMI/MAR] T= 298.15 K, 1=0 0 1.05 [1982SMI/MAR] T= 298.15 K, 1=0.5 0.5 1.05 [1982SMI/MAR] T= 298.15 K, 1=1 1 1.08 [1982SMI/MAR] T= 298.15 K, 1=2 2 1.17 [1982SMI/MAR] T= 298.15 K, 1=3 3 1.40 [1982SMI/MAR] T= 298.15 K, 1=4 4 1.80 [1982SMI/MAR] T=298.15K, 1=6 6 1.12 [1982WAG/EVA] T= 298.15 K, 1=3 3 1.29 [1984HOU/KEL] T= 298.15 K, 1=0 0 1.50 [1985BAB/MAT] T= 298.15 K, 1=0 0 1.80 1 [1985GAL] T= 298.15 K, 1=0 0 2.08 [1987BRO/WAN], T= 298.15 K, 1=0 0 1.76 [1988PHI/HAL] T= 298.15 K, 1=0 0 0.62 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ log /?/>2- Sn + 2CI- & SnCl2° 2.25 [1976SMI/MAR] T= 298.15 K, 1=0 0 1.72 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.72 [1976SMI/MAR] T= 298.15 K, 1=3 3 2.35 [1976SMI/MAR] T= 298.15 K, 1=4 4 2.43 [1976VAS/KOK] T= 298.15 K, 1=0 0 HC1O4 1.47 [1976V AS/KOK] T= 298.15 K, 1=0.5 0.5 HC1O4 1.51 [1976VAS/KOK] T=298.15K, 1=1 1 HC1O4 1.61 [1976V AS/KOK] T= 298.15 K, 1=2 2 HC1O4 1.79 [1976 V AS/KOK] T= 298.15 K, 1=3 3 HC1O4 1.57 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 1.72 [1979VAS/GLA] T= 298.15 K, 1=3 3 HC1O4 1.90 [1979VAS/GLA] T= 298.15 K, 1=4 4 HC1O4 2.43 [1982SMI/MAR] T= 298.15 K, 1=0 0 1.42 [1982SMI/MAR] T= 298.15 K, 1=0.5 0.5 1.50 [1982SMI/MAR] T= 298.15 K, 1=1 1 1.68 [1982SMI/MAR] T= 298.15 K, 1=2 2 1.75 [1982SMI/MAR] T= 298.15 K, 1=3 3 2.24 [1982SMI/MAR] T= 298.15 K, 1=4 4 3.04 [1982SMI/MAR] 1- 298.15 K, 1=6 6 1.73 [1982 WAG/EVA] T=298.15K, 1=3 3 2.06 [1984HOU/KEL] T= 298.15 K, 1=0 0 2.26 [1985BAB/MAT] T= 298.15 K, 1=0 0 2.25 [1986KOV/RYZ] T= 298.15 K, 1=0 0 3.46 [1987BRO/WAN], T= 298.15 K, 1=0 0 1.72 [1988PHI/HAL] T= 298.15 K, 1=0 0 1.43 [1992PEA/BER] T=298.15K, 1=0 0 2+ log Pu: Sn + 3Cl- <=> SnClf 2.00 [1976SMI/MAR] T= 298.15 K, 1=0 0 1.50 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.70 [1976SMIMAR] T= 298.15 K, 1=3 3 2.50 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.40 [1976V AS/KOK] T= 298.15 K, 1=0 0 HC1O4 0.44 [1976V AS/KOK] T= 298.15 K, 1=0.5 0.5 HC1O4 66 JNC TN8400 99-011 Table 4.21: continued 0.78 [1976VAS/KOK] T= 298.15 K, 1=1 1 HC1O4 1.18 [ 1976V AS/KOK] T= 298.15 K, 1=2 2 HC1O4 1.66 [ 1976V AS/KOK] T= 298.15 K, 1=3 3 HC1O4 1.19 [1979VAS/GLA] T= 298.15 K, 1=2 2 HC1O4 1.68 [1979VAS/GLA] T= 298.15 K, 1=3 3 HC1O4 2.19 [ 1979V AS/GLA] T= 298.15 K, 1=4 4 HC1O4 1.30 [1982SMI/MAR] T= 298.15 K, 1=2 2 1.67 [1982SMI/MAR] T= 298.15 K, 1=3 3 2.24 [1982SMI/MAR] T= 298.15 K, 1=4 4 3.30 [1982SMI/MAR] T= 298.15 K, 1=6 6 1.60 [1982WAG/EVA] T= 298.15 K, 1=3 3 2.02 [1984HOU/KEL] T= 298.15 K, 1=0 0 2.06 [1985BAB/MAT] T= 298.15 K, 1=0 0 4.26 [1987BROAVAN], T= 298.15 K, 1=0 0 2.05 [1988PHI/HAL] T= 298.15 K, 1=0 0 0.88 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ 2 log Pl4: Sn + 4CI- <=> SnCl4 - 1.81 [1975KRA] T= 298.15 K, 1=0 0 1.50 [1976SMI/MAR] T= 298.15 K, 1=0 0 2.30 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.70 [1984HOU/KEL] T= 298.15 K, 1=0 0 4.54 [1987BROAVAN], T= 298.15 K, 1=0 0 1.47 [1988PHI/HAL1 T= 298.15 K, 1=0 0 2+ + log Pi Sn + H2O + Ci <=> SnOHCl°+ H -0.62 [1982WAG/EVA] T= 298.15 K, I=n/a -0.41 [1984HOU/KEL] T=298.15K, 1=0 -0.64 ri988PHI/HAL1 T= 298.15 K, I=n/a- 2+ log Ks0: Sn + 2CI <=> SnCl2(s) 2.3 [1952LAT] T= 298.15 K,I=n/a 1.9 > [1963WIC/BLO] T= 298.15 K,I=n/a -0.1 [1977BAR/KNA] T= 298.15 K,I=n/a -0.1 [1979KUB/ALC] T= 298.15 K,I=n/a 2.1 [1984HOU/KEL] T= 298.15 K, 1=0 0 2.2 [1985BAB/MAT] T= 298.15 K, 1=0 0 2.2 [1985GAL] T= 298.15 K, 1=0 0 2.2 [1988PHI/HAL] T= 298.15 K, 1=0 0 2+ + log K'so: Sn + H2O + Ci <=> SnOHCl(s) + H 2.28 [1980BEN/TEA] T= 298.15 K,I=n/a 2.75 [1982WAG/EVA] T= 298.15 K,I=n/a 3.00 [1984HOU/KEL] T= 298.15 K, 1=0 2.73 [1988PH1/HAL1 T= 298.15 K,I=n/a 1 calculated with a A(G° of -26.42 kJ/mol for Sn2+ (Section 4.4.1) 67 JNC TN8400 99-011 4.8 Tin(II) fluoride system Tin(II) forms complexes and compounds with fluoride. Experimentally determined data are given in Table 4.22. Further thermodynamic data are compiled in Table 4.23 and 4.24. Table 4.22: Experimentally determined equilibrium data compiled for the tin(II) fluoride system. These data were chosen for the evaluation of recommended values in the present report. Method: pot = potentiometry. log (3 Reference Comments KM) Medium Method log Pij: Sn2+ + F- <=> SnF+ 4.14 1 [1961CON/PAU] T= 298.15 K, 1=0.5 0.5 HC1O4 pot l 4.05 [1961CON/PAU] T= 298.15 K, 1=2 2 HC1O4 pot 4.00 [1970BONATAY] T= 298.15 K, 1=1 1 NaC104 pot 2 4.48 [1971BON1] T= 298.15 K, 1=0.85 0.85 NaC104 pot 3.60 [1975NEL/AM] T= 298.15 K, 1=0.1 0.1 NaF pot 2+ logfr.2: Sn + 2F- a SnF2° 6.85 [1970BON/TAY] T= 298.15 K, 1=1 1 NaC104 pot 8.18 2 [1971BON1] T= 298.15 K, 1=0.85 0.85 NaClO4 pot 7.04 [1975NEL/AMI] T= 298.15 K, 1=0.1 0.1 NaF pot 2+ log pli3: Sn + 3F- <=> SnF3- 9.43 [1970BON/TAY] T= 298.15 K, 1=1 1 NaC104 pot 2 10.30 [1971BON1] T= 298.15 K, 1=0.85 0.85 NaC104 pot 9.00 [1975NEL/AMJJ T= 298.15 K, 1=0.1 0.1 NaF pot 1 mean of several determinations, no higher Sn(II) fluoride complexes considered. Calculated using a pK(HF/F") of 2.92 (1=0.5) and 3.06 (1=2). 2 recalculation of measurements of [ 1968HAL/SLA] by [ 1971 BON 1 ]. 68 JNC TN8400 99-011 4.8.1 SnF+, SnF2°, and SnF3- Equilibrium constants for the formation of tin(II) fluoride complexes found in the literature are given in Tables 4.22 and 4.23. From the data given in Table 4.22 the following formation constants for the tin fluoride complexes can be derived with the SIT equation at 1=0 (see also Figure 4.11 to 4.13): 2+ Sn -h F" <=> SnF+ log.P°u=4.47 2+ Sn .h2F SnF2° log 3°i,2 = 7.74 2+ Sn Hh3F SnF3- log (3°i,3 = 9.61 As the data at higher ionic strength (Figures 4.11 to 4.13) show a considerable spread, the given constants should be regarded as preliminary. Sn2+ SnF+ 7.5 -• 7 •• y = 0.32x + 4.47 Q 6.5 + 6 + 5.5 oa 3.5 3 0.5 1 1.5 2.5 lm, molal 2+ + Figure 4.11: Plot of log (3U + 4 D vs. Im for the reaction : Sn + F" <=> SnF at 25 °C. The straight line shows the result of the linear regression: Ae = -0.32; log (3°],] = 4.47. Calculated from data given in Table 4.22. 69 JNC TN8400 99-011 SnF,° 1 1.5 L, molal 2+ Figure 4.12: Plot of log (3U + 6 D vs. Im for the reaction : Sn + 2F" <=> SnF2° at 25 °C. The straight line shows the result of the linear regression: Ae = -0.91; log 3°ii2 = 7.74. Calculated from data given in Table 4.22. Sn2+ + 3P <=> SnF, 13 0.5 1 1.5 2.5 lm> molal 2+ Figure 4.13: Plot of log pu + 6 D vs. Im for the reaction : Sn + 3P <=> SnF3- at 25 °C. The straight line shows the result of the linear regression: Ae = -1.40; log (3°i)3 = 9.61. Calculated from data given in Table 4.22. 70 JNC TN8400 99-011 4.8.2 SnF2(s) Direct measurements of the solubility of SnF2(s) are not available. Based on the data compiled of [1963WIC/BLO] (Table 4.24) it can be expected that SnF2(s) is rather soluble. 4.8.3 Additional equilibrium data compiled for the tin(II) fluoride system Table 4.23: Additional, experimentally determined equilibrium data compiled for tin(II) fluoride system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry. log Pm,n,o Reference Comments I (M) Medium Method log Pij: Sn2+ + F- <=> SnF+ 6.26 ' ri968HAL/SLA1 T= 298.15 K, 1=0.85 0.85 NaC104 pot 2+ log Pia: Sn + 2F- <=> SnF2° 8.76 ' ri968HAL/SLAl T= 298.15 K, 1=0.85 0.85 NaC104 pot 2+ log pu: Sn + 3F- <=> SnFf 9.25 l [1968HAL/SLA1 T= 298.15 K, 1=0.85 0.85 NaCIO, pot these values were recalculated by [1971BON1], see Table 4.22 Table 4.24: Thermodynamic data compiled for the formation of tin(II) fluoride system. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log P Reference Comments I (M) Medium 2+ log Pu: Sn + F- <=> SnF+ 4.08 [1976SMI/MAR] T= 298.15 K, 1=1 1 4.45 ' [1980BON/HEF] T= 298.15 K, 1=1 1 NaC104 4.16 [1980BEN/TEA] T= 298.15 K, 1=0 0 4.62 [1982WAG/EVA] T= 298.15 K, 1=0 0 3.52 [1987BRO/WAN] T= 298.15 K, 1=0 0 6.9 [1992PEA/BER] T= 298.15 K, 1=0 0 2+ log P,_2: Sn + 2F- <=> SnF2° 6.68 [1976SMI/MAR] T= 298.15 K, 1=1 1 7.43 ' [1980BON/HEF] T= 298.15 K, 1=1 1 NaC104 6.59 [1987BROAVAN] T= 298.15 K, 1=0 0 9.7 [1992PE A/BER] T= 298.15 K, 1=0 0 71 JNC TN8400 99-011 Table 4.24: continued 2+ log (3!3: Sn + 3F- <=> SnFy 9.5 [1976SMI/MAR] T= 298.15 K, 1=1 11 11.10 ' [1980BON/HEF] T= 298.15 K, 1=1 1 NaC104 9.33 [1987BROAVAN] T= 298.15 K, 1=0 0 10.20 fl992PEA/BERl T=298.15K, 1=0 0 2+ 2 log pli4: Sn + 4F- <=> SnF4 - 9.81 [1987BROAVAN1 T== 298.15 K, 1=0 2+ log K*so: Sn + 2F- <=> SnF2(s) 5.6 2 ri963WIC/BL01 T= 298.15 K, I=n/a 1 calculated using a pK (HF/F) of 2.94 (1=1). 2 calculated with a Afi° of -26.42 kJ/mol for Sn2+ (Section 4.4.1) 72 JNC TN8400 99-011 4.9 Tin(II) carbonate system No experimental data are available on complex formation between tin(H) and carbonate species. It is probable that Sn2+, similar to Pb2+ (cf. Section 6.6) forms complexes with carbonate and bicarbonate ions. Thermodynamic data calculated by [1987BROAVAN] are compiled in Table 4.25. Table 4.25: Thermodynamic data for the Sn(II) carbonate system, taken from [1987BRO/WAN]. The following table serves only for comparison. log $i i log Pi 2 log Pi 3 log P, 4 References Comments 2+ 2 2m 2n log pm,n: mSn + nCO3 - <=> Snm(CO3)n - 9.72 17.85 24.65 30.21 [1987BRO/WAN] T= 298.15 K, 1=0 2+ 2 log pmn: mSn + nHCOf <=> Snm(HCO3)n ">-" 4.57 7.66 9.51 10.24 fl987BROAVAN1 T= 298.15 K, 1=0 log K*so Reference Comments 2+ log K*so: Sn + CO3 <=> Sn(CO3)(s) 10.95 ri987BROAVANl T= 298.15 K, 1=0 73 JNC TN8400 99-011 4.10 Tin(II) nitrate system Tin(II) forms weak complexes and compounds with nitrate. Experimental data were determined by [1980AND/SAM] (Table 4.26). Further thermodynamic data are compiled in Table 4.27 and 4.28. Table 4.26: Experimentally determined equilibrium data compiled for the tin(II) nitrate system. These data were chosen for the evaluation of recommended values in the present report. Method: pot = potentiometry. !„„ n Reference Comments I (M) Medium Method 1U5 Pm,n,o 2+ log Pu: Sn + NOf <=> SnNO3" 0.44 [1980AND/SAM] T= 298.15 K, 1=1 1 NaNO3 pot 0.41 [1980AND/SAM] T= 298.15 K, 1=2 2 NaNO3 pot 0.14 [1980AND/SAM] T= 298.15 K,I=3 3 NaNO3 pot 0.15 [1980AND/SAM] T= 298.15 K, 1=4 4 NaNO3 pot 0.18 [1980AND/SAM] T= 298.15 K, 1=6 6 NaNO3 pot 2+ log Pi,:>; Sn + 2NOf t=> Sn(NO3)2° 0.45 [1980AND/SAM] T= 298.15 K, 1=2 2 NaNO3 pot 0.05 [1980AND/SAM] T= 298.15 K, 1=3 3 NaNO3 pot -0.06 [1980AND/SAM] T= 298.15 K, 1=4 4 NaNO3 pot 0 [1980AND/SAM] T= 298.15 K, 1=6 6 NaNO3 pot 2+ log PU: Sn + 3NO3- <=> Sn(NO3)f -0.35 [1980AND/SAM] T= 298.15 K, 1=3 3 NaNO3 pot -0.58 [1980AND/SAM] T= 298.15 K, 1=4 4 NaNO3 pot -0.85 [1980AND/SAM] T= 298.15 K, 1=6 6 NaNO3 pot 2+ 2 log pli4: Sn + 4NO3- & Sn(NO3)4 - -0.98 [1980AND/SAM] T= 298.15 K, 1=4 4 NaNO3 pot -1.2 [1980AND/SAM] T= 298.15 K, 1=6 6 NaNO3 pot 74 JNC TN8400 99-011 2 4.10.1 SnNO3+, Sn(NO3)2°, Sn(NO3)3- and Sn(NO3)4 - From the experimentally determined data by [1980AND/SAM] in 1, 2, 3, 4, and 6 M NaNO3 (Table 4.26) the following formation constants for the tin nitrate complexes can be derived (see also Figures 4.14 to 4.17): 2+ Sn + NO3- SnNO3+ log P°i.i = 1-25 2+ Sn + 2NO3- Sn(NO3)2° log P°i,2= 1-74 2+ Sn + 3NO3- Sn(NO3)3- log P°li3 = 1.37 2+ 2+ Sn + 4NO3- Sn(NO3)4 log P°li2 = 0.30 Sn2+ + NO SnNO, 2+ + Figure 4.14: Plot of log |3U + 4 D vs. Im for the reaction : Sn + NO3" <=> SnNO3 at 25 °C. The straight line shows the result of the linear regression: Ae = 0.02; log (3°ii = 1.25. Calculated from data given in Table 4.26. 75 JNC TN8400 99-011 2+ Sn + 2NO3- <=> Sn(NO3)2° 4.5 • 4 Q CD 3.5 y = -0.05x + 1.74 + 3 2.5 CO. O) 2 .o " • " o : "• ' ••"••' 1.5 ° o—— Q_ 1 0.5 n : J 1 __ 1 4 L molal 2+ Figure 4.15: Plot of log p1>2 + 6 D vs. Im for the reaction : Sn + 2NO3- o Sn(NO3)2° at 25 °C. The straight line shows the result of the linear regression: Ae = 0.05; log (3°ii2 = 1.74. Calculated from data given in Table 4.26. Sn(NO3)3 2+ Figure 4.16: Plot of log p1>3 + 6 D vs. Im for the reaction : Sn + 3NO3" <=> Sn(NO3)3" at 25 °C. The straight line shows the result of the linear regression: Ae = 0.12; log (3°i3 = 1.37. Calculated from data given in Table 4.26. 76 JNC TN8400 99-011 2+ ;2 Sn + 4NO3" <=> Sn(NO3)4 - lm, molal 2+ 2 Figure 4.17: Plot of log p1>4 + 4 D vs. Im for the reaction : Sn + 4NO3" <=> Sn(NO3)4 - at 25 °C. The straight line shows the tentative result of the 'linear regression': Ae = 0.11; log p°1>4 « 0.30. Calculated from data given in Table 4.26. 4.10.2 Tin(II) nitrate compounds No thermodynamic data about tin(II) nitrate compounds are available. It can be expected, that they are very easily soluble and will not limit for Srt(II) solubility. 4.10.3 Additional equilibrium data compiled for tin(II) nitrate system Table 4.27: Additional, experimentally determined equilibrium data compiled for tin(II) nitrate system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry. log (3m,n,0 Reference Comments I (M) Medium Method 2+ log (5,.,: Sn + NOy <=> SnNO3 1.1 [1980AND/SAM] T= 298.15 K, 1=0 NaNO3 pot 0.43 ' [1980AND/SAM] T=298.15K, 1=8 pot 2+ 0 log f3u: Sn + 2NO3- <=> Sn(NO})2 0.57 ' [1980AND/SAM] T= 298.15 K, 1=8 pot 77 JNC TN8400 99-011 Table 4.27: continued 2+ log p13: Sn + 3NO3- <=> Sn(NO3)3- 0.46 ' [198QAND/SAM1 T= 298.15 K, 1=8 8 NaNCh pot 2+ 2 log P,,4: Sn + 4NO3 <^> Sn(NO3)4 - -0.01 ' [1980AND/SAM1 T= 298.15 K, 1=8 8 NaNO^ pot 1 the values determined at 1=8 are quite different from the others and are therefore not chosen for extrapolation to 1=0 in this report. Table 4.28: Thermodynamic data compiled for the formation of tin(II) nitrate system. . As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. l°g Pm.n,o Reference Comments I (M) 2+ + log PJJ: Sn + NOf <=> SnNO3 0.44 [1976SMI/MAR] T= 298.15 K, 1=1 1 1.90 ri987BROAVAN1 T= 298.15 K, 1=0 0_ 2+ log PL2: Sn + 2NOy <=> Sn(NO3)2° 2.25 [1987BRO/WAN1 T= 298.15 K, 1=0 2+ log Pu: Sn + 3NOf <=> Sn(NO3)f 1.32 f!987BRO/WAN1 T= 298.15 K, 1=0 2+ 2 log P14: Sn + 4NO3- <=> Sn(NO3)4 - -0.80 [1987BRO/WAN] T= 298.15 K, 1=0 78 JNC TN8400 99-011 4.11 Tin(H) phosphate system 3 [1969AWA/KAS] determined the redox potential of the reaction 3Sn(cr) + 2PO4 " <=> Sn3(PO4)2(s) + 6e" in phosphate solution. From the standard potential of -0.865 V given by 2+ 3 [1969AWA/KAS], a log K*So of 73.93 for the reaction 3Sn + 2PO4 - <^ Sn3(PO4)2(s) (Table 4.29) can be calculated. The usage of this value is not recommended as the ionic strength in the experiments of [1969AWA/KAS] is varied strongly. [1969AWA/KAS] also proposed the formation of tin phosphate complexes without being able to calculate formation constants. No other experimental data are available about complex formation between tin(II) and phosphate. Thermodynamic data given in previous compilations are compiled in Table 4.30. Table 4.29: Experimentally determined equilibrium data compiled for tin(II) phosphate system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry. log K*so Reference Comments I (M) Medium Method 2+ 3 log K*so: 3Sn + 2PO4 - <=> Sn3(PO4)2(s) 73.84' [1969AWA/KASJ T= 303 K, 1=0.0005-2 var Na3PO4 pot 1 3 2+ calculated from E° = -0.865 V for 3Sn(cr) + 2PO4 " <=> Sn3(PO4)2(s)+ 6e- and log K (Sn /Sn(cr)) of 4.63 (Section 4.4.1) Table 4.30: Thermodynamic data for the Sn(II) phosphate system, taken from previous compilations. As pointed out in Section 2 of this report' only experimental data were used for the present evaluation. The following table serves only for comparison. log $! log P2 log P3 log p4 References Comments 2+ + 2 2n log ft: Sn + nPO/' + nH <=> Sn(HPO4) - 18.62 36.40 53.60 70.31 [1987BROAVAN] T= 298.15 K, 1=0 2+ + 2 n log ft: Sn + nPO/- + 2nH <=> Sn(H2PO4) - 21.71 42.52 62.70 82.32 [1987BRO/WAN] T= 298.15 K, 1=0 log K*so Reference Comments 2+ 3 log K*s0: 3Sn + 2PO4 - <=> Sn3(PO4)2(s) 71.9 ' [1985GAL1 T= 298.15 K, I=n/a 1 calculated with a A(G° of -26.42 kJ/mol for Sn2+ (Section 4.4.1) 79 JNC TN8400 99-011 4.12 Tin(II) sulfate system Similar to Pb(II), Sn(II) forms complexes and compounds with sulfate. Experimental data for complex formation of Sn(II) with sulfate were determined by [1981PET/MIL] in 1 M NaNO3 (Table 4.31). Further data are compiled in Table 4.32. Table 4.31: Experimentally determined equilibrium data compiled for the tin(EE) sulfate system. These data were chosen for the evaluation of recommended values in the present report. Method: pot = potentiometry. R Reference Comments I (M) Medium Method Pm,n,o 2+ 2 log Pi,j: Sn + SO4 - & SnSO4° 1.29 [1981PET/MEL] T= 298.15 K, 1=1 1 NaNQ, pot 2+ 2 2 log pli2: Sn + 2SO4 - <^ Sn(SO4)2 - 1.65 [1981PET/MIL] T= 298.15 K, 1=1 1 NaNO, pot 4.12.1 Tin(II) sulfate complexes From the experimentally determined data by [1981PET/MTL] in 1 M NaNO3 (Table 4.31) the following tentative formation constants for the tin sulfate complexes can be calculated: 2+ 2 Sn + SO4 - «=> SnSO4° log (3°u = 2.91, Ae = - 0.01 2+ 2 2 Sn + 2SO4 - <=> Sn(SO4)2 - log p°i'i2 = 2.83, Ae = -0.43 The Ae values are assumed to be the same as calculated for Pb (Section 6.9.1: Lead sulfate complexes). 80 JNC TN8400 99-011 4.12.2 SnSO4(s) SnSC>4(s) might be expected to be moderately soluble, similar to anglesite (PbSC>4(cr); Section 6.9.2). The values compiled in Table 4.32, however, are unusually high, indicating a very low solubility and might be somewhat suspect. In these compilations no primary source of information is given. [1987BRU] conducted a specific computer search in Chemical Abstracts for solubility data about SnSO4(s) which did not turn up any useful information. Therefore, no solubility product for SnSO4(s) is recommended. The solubility of SnSC>4(s) may be much larger than indicated by the values compiled in Table 4.32. Table 4.32: Thermodynamic data compiled for the formation of tin(II) sulfate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log f3m,n,0 Reference Comments I (M) 2+ 2 log P,j: Sn + SO4 - <=> SnSO4° 3.12 [1987BRO/WAN] T= 298.15 K, 1=0 1.35 [1988PHI/HAL1 T= 298.15 K, I=n/a 2+ 2 2 log plX- Sn + 2SO4 - <=> Sn(SO4)2 - 5.02 H987BRO/WAN1 T= 298.15 K, 1=0 2+ 2 4 log Pu: Sn + 3SO4 - <=> Sn(SO4)} - 5.94 ri987BRQ/WAN] T= 298.15 K, 1=0 2+ 2 log P!A: Sn + 4SO4 - <=> Sn(SO4)f- 5.99 [1987BRO/WAN] T= 298.15 K, 1=0 0 2+ 2 log K'so.• Sn + SO4 - <=> SnSO 24.24 ' [1977BAR/KNA] T= 298.15 K, I=n/a 24.24 ' [1979KUB/ALC] T= 298.15 K, I=n/a 49.86 [1985GAL] T= 298.15 K, I=n/a 49.86 [1988PHKHAL] T= 298.15 K, I=n/a 23.93 [1992PEA/BER1 T= 298.15 K, 1=0 0 calculated with a AfG° of -26.42 kJ/mol for Sn2+ (Section 4.4.1) JNC TN8400 99-011 4.13 Tin(II) sulfide system 4.13.1 SnS(herzenbergite) Chemically, SnS(cr), herzenbergite, is expected to be quite insoluble, as indicated by the values compiled in Table 4.33. No determination of solubility data could be found in the literature and in the compilations listed in Table 4.33 no primary source of information is given. Therefore, no solubility product for SnS(cr) is recommended in this report. 4.13.2 Sn2Ss(s) and SnjS^s) Thermodynamic data for the formation of Sn2S3(s) and Sn3S4(s) have been proposed in different compilations (Table 4.33). However, no experimental determination of the solubility of these solids could be found in the literature and in the compilations listed in Table 4.33 no primary source of information is given. Thus, no solubility products for Sn2S3(s) and are recommended. Table 4.33: Thermodynamic data compiled for the formation of tin(II) sulfide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK *so Reference Comments KM) logK • Sn2+ + HS- <=> SnS(s, herzenbergite) + H+ so-' 15.55 1 [1964HIR] T= 298.15 K, I=n/a 14.74 1 [1968SUS/KHO] T= 298.15 K,I=n/a 14.78 1 [1971NAU/RYZ] T= 298.15 K, I=n/a 16.11 1 [1974MBL] T= 298.15 K, I=n/a 12.00 [1976SMI/MAR] T= 298.15 K, 1=0 0 15.86 1 [1978ROB/HEM2] T=298.15K, I=n/a 15.57 [1979KUB/ALC] T=298.15K, I=n/a 15.72 [1980BEN/TEA] T= 298.15 K,I=n/a 14.74 1 [1982WAG/EVA] T= 298.15 K, I=n/a 16.33 [1985GAL] T= 298.15 K, I=n/a 15.86 1 [1985JAC/HEL2] T= 298.15 K,I=n/a 12.00 [1986MYE] T= 298.15 K,I=n/a 14.60 [1988PHI/HAL] T= 298.15 K,I=n/a 2+ + logK *so' 2Sn + 3HS- <=> Sn2S3(s) + 3H + 2e- 41.70 J [1974MEL] T= 298.15 K,I=n/a 41.58 1 [1977BAR/KNA] T= 298.15 K, 1=0 0 41.58 1 [1979KUB/ALC1 T= 298.15 K,I=n/a logK * 3Sn2+ + 4HS- <=> Sn S (s)+ 4//++ 2e- so-' 3 4 57.54 1 [1974MIL] T= 298.15 K,I=n/a 57.57 1 [1977BAR/KNA] T= 298.15 K, 1=0 0 57.57 1 [1979KUB/ALC1 T= 298.15 K, I=n/a calculated with a AfG0 of -26.42 kJ/mol for Sn2+ (Section 4.4.1) 82 JNC TN8400 99-011 4.14 Comments on selected references [1939GOR]: [1939GOR] measured Sn(II) solubility under acidic conditions, assuming only the presence of Sn2+ and SnOH+ and neglecting the formation of polymeric species. They extrapolated their results to 1=0 with the Debye-Hiickel term. [1941GAR/HEI]: [1941GAR/HEI] measured Sn(II) solubility in 0 to 0.4 M HC1 and NaOH 2+ + assuming only the presence of Sn , SnOH , Sn(OH)2° and Sn(OH)3" at variable I. They neglected the formation of polymeric species. Under acidic conditions, the formation of tin(II) chloride complexes cannot be excluded. This work also suffers from the difficulty that the free acidity was not measured but was calculated indirectly from assumed equilibria. In water [1941GAR/HEI] determined the solubility of 6 SnO(cr) to be 5 x 10" M and reported for Sn(OH)2(precip) a solubility of 1.35 x 10" 5 M (determined by [1906GOL/ECK]). From these values and a log (3°i,2 of -7.71 2+ + for the reaction Sn + 2H2O <=> Sn(OH)2° + 2H (Section 4.5.2): 2+ Sn + H2O <=> SnO(s) + 2H+ log K*°so = -2.41 2+ Sn +2H2O<=} Sn(OH)2(precipitated) + 2H+ log K*°s0 =-2.84 The values given by [1941GAR/HEI] themselves are different (Table 4.17 and 4.18) as [1941GAR/HEI] used a different log p1>2 value. [1942GOR/LEI]: [1942GOR/LEI] measured Sn(II) solubility in 0 - 0.1 M HC1O4 at variable I under acidic conditions. The free acidity was not measured but was calculated indirectly from assumed equilibria. [1952VAN/RHO]: Sn(II) concentration = 200 mM. Investigated hydrolysis of tin(II) and complex formation with chloride. Polymeric tin hydroxide species were not taken into account by [1952VAN/RHO]. [1952VAN/RHO] give for the reaction SnOH+ + 0 2+ Cl- <=> SnOHCl a log KliU of 1.04. With a log (3U of -3.81 for Sn + H2O <=> 2+ SnOH+ + H+ at I = 3 (see Table 4.13) a log (3u,i of -2.77 for Sn + H2O + Cl" <=> SnOHCl0 + H+ at I = 3 can be calculated. [1963FEI/SCH]: Review of solubility of metal oxides and hydroxides in water. The data of Sn(OH)4, SnO2(precip) and SnO2(cassiterite) are from unpublished experimental 4+ results [1957EGG]. [1963FEI/SCH] gives a log K*so of 5.0 for Sn + 4H2O ^> + SnO2 + 4H . They state that to them no data concerning the hydrolysis of Sn(IV) are available. Experimental details such as pH, ionic strength, ... are not given. However, considering that Sn(OH)4° dominates Sn(IV) speciation between pH 2 to 8, it was assumed in this report that in the experiments cited by [1963FEI/SCH], Sn(OH)4°, and not Sn4+, is the dominant dissolved species. 83 JNC TN8400 99-011 [1964LIN/TU]: [1964LIN/TU] recalculated the results of [1958TOB] with least-squares analysis and questioned the existence of SnOH+. In place of this species they proposed the existence of Sn2(OH)3+. However, the interpretation of [1958TOB] seems more probable. [1976GOB] confirmed in her work the existence of Sn3(OH)42+ and SnOH+ and rejected the formation of the species Sn2(OH)3+ proposed by [1964LIN/TU]. [1965MES/IRA]: The solubility measurements of [1965MES/IRA] in 1 M NaC104 and in presence of 0.5-10 mM Sn allows one to assign a lower limit for the reactions: 2+ Sn + H2O <=> SnO(cr) + 2H+ log K*So < -2.29 + SnOH+ o SnO(cr) + H log K*Si = 0.28 As the formation of Sn(OH)2° and polynuclear Sn3(OH)42+ is not taken into account, this are only estimates. [1970BAR/KLI]: [1970BAR/KLI] investigated the hydrolysis of tin(TV) in 0.2 - 2.5 M NaOH. In contrast to [1997AMA/CHI] who proposed the formation of Sn(OH)5- and Sn(OH)62" under alkaline conditions, [1970BAR/KLI] found only evidence for the presence of Sn(OH)5". From the experimental details reported in the paper of [1970BAR/KLI] it seems that pH was not measured directly but determined from the amount of NaOH initially added to the system, while the Sn concentration was measured in the solutions after one month shaking in contact with atmospheric CO2 which has a strong influence on both solution composition and resulting pH. Thus the calculated pH is probably significantly larger than a measured pH would be. [1970BAR/KLI] also measured the solubility of cassiterite in water, dilute HNO3 and dilute NaOH. They observed a constant, pH-independent (between pH 2 - 11) solubility of 4 x 10'7 M Sn(IV)/L. Again, the pH seems to be calculated and not to be measured directly. No detection limit for Sn(IV) is indicated which makes it difficult to decide whether the measured concentrations below pH 11 correspond to the detection limit or whether they are real concentrations. [1970KUR/BAR]: [1979KUR/BAR] determined constants for the hydrolysis of Sn(OH)4° at 100 °C. In contrast to observations at 25 °C, protonation of the Sn(OH)4° ion was observed by [1979KUR/BAR] already at a pH of 7. [1971NAZ/ANT]: [1971NAZ/ANT] studied the hydrolysis of tin(IV) under acidic conditions in 1 M KNO3. [1971NAZ/ANT] used a spectrophotometric method in which the competition between the hydrolysis reaction and complexation with salicylfluorone was measured. The total Sn(IV) concentration in these experiments was 10~5 M. Considering the possible interaction of Sn(IV) with the nitrate ions in solution, these log p values can be considered as useful estimates, but not as exact values for the 84 JNC TN8400 99-011 stability constants of Sn(IV) hydrolysis at I = 1. Nevertheless, based on the observation of [1971NAZ/ANT] and [1997AMA/CHI] it is clear that protonation of Sn(OH)4° will occur only at pH < 1. [1973GAB/SRI]: [1973GAB/SRI] determined the redox potential of the reactions Sn(cr) + 2 3OH- o Sn(OH)3- + 2e" and Sn(OH)3- + 3OH" <=> Sn(OH)6 " + 2e- as 0.90 V and 0.910 V, respectively. From their data log (3 values for the hydrolysis of Sn2+ and Sn(OH)4 can be estimated (Table 4.2 and Table 4.14). Assumptions: log K\y = - 13.76, log K (Sn2VSn(cr)) = 5.0 and log K (Sn(OHySn(cr)) of -0.77 (at 1=0.25). These values can be considered only as estimates due to possible error introduced by estimating the redox potential at 1=0.25. [1973KLI/BAR]: The values given in [1973KLI/BAR] for the hydrolysis of Sn(OH)4° are calculated from the data measured by [1971NAZ/ANT] in 1 M KNO3. The solubility of synthetic cassiterite is determined at different temperatures. [1973KLI/BAR] determined in the temperature range 473 - 673 K a higher solubility of cassiterite than [1981DAD/SOR] and [1988BAR/SHA] (cf. Table 4.5). Also at 298 K, the solubility measured by [1973KLI/BAR] of 4 10"7 M is much higher than observed in other references [1926GRU/LIN, 1963FEI/SCH, 1997AMA/CHI] (Table 4.5). The reason of this apparent higher solubility observed by [1973BAR/KLI] is not clear. This paper was not chosen in the present report. [1974GOB]: same results as given in [1976GOB]. [1976GOB]: [1976GOB] studied with potentiometric titration the hydrolysis of the tin(II) ion in 0.02 -2.3 mM in the pH range 2.7 - 3.7. [1976GOB] showed in her work that the species Sn3(OH)42+, SnOH+ and Sn2* are dominant under the conditions of her experiments, similar to [1958TOB]. With her data she could not confirm (nor exclude) the presence of Sn2(OH)22+ which was also proposed by [1958TOB]. [1977MAR]: [1977MAR] studied the hydrolysis of tin(II) in 3 M NaC104 in alkaline solution. Sn(II) = 1 mM; log Kw = 14.18 [1978DIC/LOT1]: [1978DIC/LOT1] determined the redox potential of the reactions Sn(0) + 3OH- <^=> Sn(OH)3- + 2e\ They extrapolated their data to I = 0 with the Pitzer model and obtained a redox potential of 0.88 V. From their data value for the reaction Sn2+ + + 3H2O <=> Sn(OH)3" + 3H a log (3°i;3 of -16.85 can be calculated, assuming a log 2+ K°w = -14.0, log K° (Sn /Sn(cr)) = 4.60 (Section 4.4.1). [1979VAS/GLA]: The values given in [1979VAS/GLA] for the hydrolysis of Sn(OH)4° are calculated from the data measured by [1971NAZ/ANT] in 1 M KN03. [1981DAD/SOR]: [1981DAD/SOR] determined the solubility of synthetic cassiterite in the temperature range 473 - 673 K. [1981DAD/SOR] determined an approximately 5 85 JNC TN8400 99-011 times lower solubility of cassiterite than [1973KLI/BAR]. Extrapolation of the measurements of [1981DAD/SOR] to 298 K gives a log K*So of « 7.5 for cassiterite, a value which is in fair agreement with the a log K*so of 8.06 determined recently by [1997AMA/CHI] for cassiterite (SnO2(cr)). [1981PET/MIL]: [1981PET/MIL] studied with anodic stripping voltammetry the hydrolysis of tin(II) and the complex formation of Sn(II) with CY, Br and SO42" in 0.1 - 1 M NaNO3, NaCl and artificial seawater with a Sn concentration of 10~7 M. The hydrolysis constants measured in NaCl and artificial seawater also include the formation of tin chloride complexes. [1988BAR/SHA]: [1988BAR/SHA] determined at 300 °C a much smaller solubility of synthetic cassiterite than [1973KLI/BAR]. They offer no explanation for these findings. [1997AMA/CHI]: Based on solubility measurements, [1997AMA/CHI] proposed the formation of Sn(OH)s" and Sn(OH)(52" under alkaline conditions. They extrapolated their measurements from I = 0.1 M NaC104 to I = 0 using the wrong sign for the Davies equation. For further comments see [1998ODA/AMA]. [1998ODA/AMA]: See also comments to [1997AMA/CHI]. [1998ODA/AMA] gives the uncorrected log P15 and (3ig values for I = 0.1 M NaC104 derived from the experimental data from [1997AMA/CHI] and corrected these values to I = 0 using the Davies equation. The additional experiments of [ODA/AMA] in 0.1 M NaClC^ confirmed the solubility data determined by [1997AMA/CHI]. [1998ODA/AMA] also showed that the presence of chloride and sulfate did not influence the solubility of SnC>2(precip). 86 JNC TN8400 99 -Oil 5 Antimony Antimony exists in the oxidation states -III, 0, +III and +V. It occurs in nature primarily as the sulfide Sb2S3(stibnite) or as the oxide Sb2O3(valentinite) [1976BAE/MES, 1985PAS, 1995WIB]. When Sb2C<3(cr) is heated in air, additional oxygen is taken up above 300 °C, and Sb2O4 (a mixed compound of Sb(III) and Sb(V)) can be formed [1984BER/BRE]. In water, Sb(III) is stable under reducing conditions and Sb(V) under oxidizing conditions [1976BAE/MES, 1985PAS, 1995WIB]. 5.1 Hydrolysis of antimony (III) Solubility measurements in presence of Sb2C>3(valentinite), e.g. [1952GAY/GAR] and diffusion experiments [1965JAN/HAR1, 1977KEP/TAL] show that Sb(III) exists below pH 2 as Sb(OH)2+, and in basic solutions (pH > 11) as Sb(OH)4~. For a wide pH range the solubility of Sb2O3 does not depend on pH, indicating the predominance of the uncharged Sb(OH)3° species. [1965JAN/HAR1] have shown by spectrophotometric measurements that no polymeric species are formed in significant amounts in 0.1 M Sb(III) solutions. [1965JAN/HAR1] showed that under basic conditions (in 1 to 16 N NaOH), Sb(III) exists exclusively as Sb(OH)4- (or SbO2-), while no other anionic species could be observed. [1970DAW/WIL] observed with spectrophotometric methods at Sb(III) concentrations > 0.1 mM a polymerization of antimony(III) species in sulfate solutions. [1974AHR/BOV] and [1994AKI/ZOT] concluded from solubility measurements that Sb2(OH)42+ and Sb2(OH)6 exist in concentrated (Sb > 0.1 mM) Sb(III) solutions. In the following paragraphs, all log P values refer to Sb(OH)3° as master species as the constant for the protonation of Sb(OH)3° to Sb3+ is not known with sufficient accuracy. Table 5.1: Experimental equilibrium data used for the hydrolysis of antimony(III), according n + to the equilibrium: mSb(OH)3° + nH2O <=> Sbm(OH)3m+n " + nH . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12: 'Comments on selected references'. Method: extr = solvent extraction, sol = solubility, sp = spectrophotometry. The notation log |3b refers to a reaction involving OH~. ] Reference Comments I (M) Medium Method °g Pm;, 3m+n + 3+ log A,0: Sb(OH)3° + 3H <=> Sb + 3H2O 0.43 [1977ANT/NEV] T= 298.15 K, 1=1 1 NaClO4 sp 87 JNC TN8400 99-011 Table 5.1: continued + 2+ log pjj: Sb(OH)3° + 2H s=> SbOH 2H2O 0.23 [1974AHR/BOV] T= 298.15 K, 1=5 5 HNO3 sol 1.04 [1977ANT/NEV] T= 298.15 K, 1=1 1 NaCIO, sp_ + log fiu: Sb(OH)30 + H <=> Sb(OH)2" + H2O 1.37 * [1924SCH] T= 298.15 K, 1=0.2 0.2 HC1O4 sol 1.16 1 [1924SCH] T= 298.15 K, 1=0.5 0.5 HC1O4 sol l 1.19 [1924SCH] T= 298.15 K, 1=0.9 0.9 HC1O4 sol 1.18 l [1924SCH] T= 298.15 K, 1=1.1 1.1 HC1O4 sol 1.43 [1968MIS/GUP] T= 298.15 K, 1=0.02 0.02 HC1O4 sp 1.40 [1968MIS/GUP] T= 298.15 K, 1=0.05 0.05 HC1O4 sp 1.42 [1968MS/GUP] T= 298.15 K, 1=0.1 0.1 HC1O4 sp 1.09 [1974AHR/BOV] T= 298.15 K, 1=5 5 HNO3 sol 1.09 [1974AHR/BOV] T= 298.15 K, 1=5 5 HC1O4 sol 1.23 [1974SHO/MAB] T= 298.15 K, 1=3 3 HC1O4 extr 1.23 [1974SHO/MAB] T= 298.15 K, 1=0.1 0.1 HC1O4 extr 1.05 [1977ANT/NEV] T= 298.15 K, 1=1 1 NaClO4 sp log fP1A: Sb(OH)3° + OH~ & Sb(OH)i 2.07 2 [1948TOU/MOU] T= 308 K, I=diluted 1.99 NaOH sol 2.04 2 [1948TOU/MOU] T= 308 K, I=diluted 0.7 NaOH sol 1.99 2 [1948TOU/MOU] T= 308 K, I=diluted 0.46 NaOH sol 1.93 2 [1948TOU/MOU] T= 308 K, I=diluted - 0.092 NaOH sol 1.93 2 [1948TOU/MOU] T= 308 K, I=diluted 0.042 NaOH sol 2.21 2 [1952GAY/GAR] T= 298.15 K, 1=0.1 0.1 NaOH sol 2.21 2 [1952GAY/GAR] T= 298.15 K, 1=0.075 0.075 NaOH sol 2.202 [1952GAY/GAR] T= 298.15 K, 1=0.04 0.04 NaOH sol 2.142 [1952GAY/GAR] T= 298.15 K, 1=0.04 0.04 NaOH sol 2.192 [1952GAY/GAR] T= 298.15 K, 1=0.02 0.02 NaOH sol 2.06 2 [1973VAS/SHO2] T= 298.15 K, 1=2.37 2.37 NaOH sol 2.04 2 [1973VAS/SHO2] T= 298.15 K, 1=2.03 2.03 NaOH sol 2.05 2 [1973VAS/SHO2] T= 298.15 K, 1=1.69 1.69 NaOH sol 1.99 2 [1973VAS/SHO2] T= 298.15 K, 1=1.36 1.36 NaOH sol 1.98 2 [1973VAS/SHO2] T= 298.15 K, 1=1.02 1.02 NaOH sol 1.96 2 [1973VAS/SHO2] T= 298.15 K, 1=0.68 0.68 NaOH sol 1.95 2 [1973VAS/SHO2] T= 298.15 K, 1=0.34 0.34 NaOH sol 2.06 2 [1973VAS/SHO2] T= 298.15 K, 1=2 2 NaClO4 sol 2.09 [1994AKI/ZOT] T=298, I=dil 0 self medium sol JNC TN8400 99-011 Table 5.1: continued + 2+ log j324: 2Sb(OH)3° + 2H <=> Sb2(OH)4 + 2H2O 2.86 [1974AHR/BOV] T= 298.15 K, 1=5 5 HC1O4 sol log P2i6: 2Sb(OH)3o & Sb2(OH)6° 0.08 [1994AKI/ZOT] T=298, I=dil 0 self medium sol "T calculated with log K°*S3 = 8.72 (section 5.2.1) 2 recalculated in this report from experimental values of [1948TOU/MOU], [1952GAY/GAR] and [1973VAS/SHO]. 5.1.1 Sb3+ [1965JAN/HAR1], [1965JAN/HAR2] concluded from diffusion measurements that only monomeric Sb(OH)2+, and possibly Sb3+, are present in perchloric acid solutions up to H+ concentration of 6 M and [Sb] < 1 mM. Spectrophotometric measurements of [1968MIS/GUP] (Sb concentration = 0.3 mM) indicate the presence of Sb(OH)2+ in the pH range 1-2, and at pH 3+ 2 < 0 (1 and 3 M HC1O4) the presence of a more protonated Sb complex (Sb or SbOH *). From the data given by [1977ANT/NEV] (Table 5.1) a tentative log P\o value can be 3+ extrapolated for the reaction Sb(OH)3° + 3H+ 0 Sb + 3H2O assuming a Ae of 0.07 (from e(Am3+, CIO4-) = 0.49 and E(H+, CIO4-) = 0.14; [1995SIL/BID]): 3+ Sb(OH)3° +3H+ o Sb + 3H2O log (3°u = -0.73 The values calculated from the potentiometric measurements of [1970BON/WAU] and [1970BON] in 5 and 2 M HCIO4 (Table 5.2) agree well with the log pli0 values measured by [1977ANT/NEV]. It is not clear, however, if in the experiments of Bond and co-workers [1970BONAVAU, 1970BON] Sb3+ is the only species present. The potentiometric measurements of [1975HEI/SCH] in 1.5 MHF and H2SO4 give a log p of 0.36 for the 3+ + protonation of Sb(OH)3° to Sb . [1975HEI/SCH] assumed the formation of Sb(OH)2 (at a pH < 0). We consider it as probable, considering the high H+ concentration, that the Sb3+ species is dominant in such solutions (cf. Table 5.2). 89 JNC TN8400 99-011 5.1.2 SbOH2+ 2+ For the reaction Sb(OH)3° + 2H+ <=> SbOH + 2H2O, [1974AHR/BOV] determined with solubility measurements a log Pu of 0.23 in 5 M HNO3. [1977ANT/NEV] determined with solubility and spectrophotometric measurements a log Pu of 1.04 in 1 M HCIO4 (Table 5.1) From these measurements (Figure 5.1) a tentative value can be calculated: 2+ Sb(OH)3°+2H+ SbOH + 2H2O logP°u =0.83 The Ae = 0.19 obtained from only two datapoints (Figure 5.1) agrees satisfactorily with the Ae of 0.11 estimated from E(AmOH2+, CIO4-) = 0.39 and e(H+, CIO4-) = 0.14; [1995SIL/BID]. + 2+ Sb(OH)3°+2H <4 SbOH +2H2O 2.5 2 L m&lal Figure 5.1: Plot of log 01,1 - 2 D vs. Im for the reaction Sb(OH)3° + 2H+ <=> 2H2O at 25 °C. The straight line shows the result of the 'linear regression': Ae 0.19; log P°i,i = 0.83. Calculated from data compiled in Table 5.1. + 5.1.3 Sb(OH)2 Many measurements of the hydrolysis of Sb(OH)3° in dilute acid exist. The formation constant + given in [1952LAT] and [1957PIT/POU] for SbO+ (corresponding to Sb(OH)2 ) and SbO2~ originate from the early work of [1883SCH]. The log pli2 value of 0.90 given by [1985PAS] + + for the reaction Sb(OH)3° + H <=> Sb(OH)2 + 2H2O is taken from the data measured by [1924SCH] in 0.2 - 1.1 M HCIO4. 90 JNC TN8400 99-011 [1924SCH] determined in 0.2 - 1.1 M HC1O4 both the solubility of Sb(III) in perchlorate acid + and the electrode potential of the reaction Sb(cr) + 2H2O <=> Sb(OH)2 + 2H+ + 3e~. Both measurements resulted in similar log (3ii2 values (Table 5.1 and 5.2). [1952GAY/GAR] + measured in 0 - 0.1 M HC1 a mean log pli2 of 1.13 for the reaction Sb(OH)3° + H <=> + Sb(OH)2 + H2O (Table 5.2). As Sb(III) forms complexes with chloride, these measurements were not used for extrapolation in this report. However, they agree well with other measurements reported in the literature. [1968MIS/GUP] determined spectrophotometrically a mean log (3i)2 value of 1.42 in 0.02-0.1 M HCIO4. Based on potentiometric measurements, [1975HEI/SCH] calculated a log p1>2 of 0.53 in presence of I = 1.5 M (HF and H2SO4). Although they corrected their measurements for the interactions of Sb(III) with F~ and SO42-, their measurements differ from the other data reported. Considering the high H+ concentration, the presence of Sb3+ or SbOH24" in the experiments is probable. Extrapolation of the data of [1924SCH (only solubility data), 1968MIS/GUP, 1974AHR/BOV, 1974SHO/MAB, and 1977ANT/NEV (compiled in Table 5.1) gives at I = 0, as shown in Figure 5.2: Sb(OH)3° Sb(OH)2+ + H2O log(3°li2 =1.30 The extrapolation of the experimental results with the SIT term shown in Figure 5.2 gives a Ae value of 0.02 which is in good agreement with the expected Ae value of =0 for the isocoulombic + reaction Sb(OH)3° + H+ <=> Sb(OH)2 + H2O. + + Sb(OH)3°+H o Sb(OH)2 +H2O Figure 5.2: Plot of log (3U + 0 D vs. Im for the reaction Sb(OH)3° + H+ <=> Sb(OH)2+ + H2O at 25 °C. The straight line shows the result of the linear regression: Ae = 0.04; log (3\2 = 1.30. Calculated from data compiled in Table 5.1. 91 JNC TN8400 99-011 The potentiometric measurements of [1924SCH] and [1972VAS/SHO] (in 0.3-2.5 M HC1O4) refer all to the reaction Sb(cr) + 2H2O <=> Sb(OH)2+ + 2H+ + 3e~. For comparison the log of these values were corrected for the reaction Sb(cr) + 3H2O <=> Sb(OH)3 + 3H+ + 3e~ with a factor of 11.99 (cf. Section 5.7: Redox reactions). These values agree very well with the log (312 values measured with other techniques. If the these values were also included in the extrapolation to 1=0 a log (3° 1,2 value of 1.29 results, the nearly the same value as in absence of the potentiometric measurements. 5.1.4 Sb(OH)4- Several authors determined formation constants for the reaction Sb(OH)3° + H2O <=> Sb(OH)4~ + H+. Extrapolation to I = 0 of the experimental data given by [1948TOU/MOU, b 1952GAY/GAR, 1973VAS/SHO2, 1994AKI/ZOT] compiled in Table 5.1 results in a log p \4 of 2.07 for the reaction Sb(OH)3° + OH~ <=> Sb(OH)4- (Figure 5.3). From this log (3\4 can be obtained: Sb(OH)3° + H2O <^> Sb(OH)4- \A =-11.93 The extrapolation of experimental results with the SIT term shown in Figure 5.3 gives a Ae value of 0.05 which is in good agreement with the expected Ae value of 0 for the isocoulombic reaction Sb(OH)3° + OH- <=> Sb(OH)4- Sb(OH)3° + OH" « Sb(OH)4" b Figure 5.3: Plot of log p 1?4 + 0D vs. Im for the reaction Sb(OH)3° + OH~<=» Sb(OH)4- at 25 °C (and 35 °C). The straight line shows the result of the linear regression: Ae = b + 0.05; log p °i,4 = 2.07. For the reaction Sb(OH)3° + H2O <=> Sb(OH)4- + H a log P°ij4 of- 11.93 can be calculated. Calculated from data compiled in Table 5.1. 92 JNC TN8400 99-011 5.1.5 Sb2(OH)42+ + 2+ For the reaction 2Sb(OH)3° + 2H <=> Sb2(OH)4 + 2H2O, a tentative log p2,4 of 2.43 can be extrapolated from the data given by [1974AHR/BOV], As this log (3 value is determined at high ionic strength and only by one author, the usage of this value at low ionic strength may result in a considerable error. The species Sb2(OH)42+ will be important only in concentrated Sb(III) solutions (Sb>0.1 mM). 5.1.6 Sb2(OH)6° For the reaction 2Sb(OH)3° <=> Sb2(OH)6 a log (32,6 of 0.08 is determined by [1994AKI/ZOT] from measurements at 25 - 300 °C. While this complex seems to be important at higher temperature, its concentration is small at 25 °C. 5.1.7 Additional equilibrium data compiled for the hydrolysis of antimony(III) Table 5.2: Additional experimentally determined equilibrium data compiled for the hydrolysis of antimony(IH), These data were not chosen in the present report for the evaluation of recommended stability values. Method: pol = polarography, pot = potentiometry, sol = solubility. log Pm, 3,ra+n Reference Comments KM) Medium Method 3 log Pi,0: Sb(OH)3° + 3W <=> Sb * + 3H2O 0.89 > [1970BON/WAU] T= 303 K, 1=5 5 HC1O4 pol 0.32 ' [1970BON] T= 303 K, 1=2 2 Na/HClO4 pol 2 0.53 '' [1975HEI/SCH] T= 298.15 K, 1=1.5 1.5 HF, H,SO4 pot + + log p12: Sb(OH)3° + H <=> Sb(OH),2 + H2O 1.38 ' [1924SCH] T= 298.15 K, 1=0.2-1.1 0.2 HC1O4 pot 1.18 ' [1924SCH] T= 298.15 K, 1=0.24.1 0.5 HC1O4 pot 1.19 ' [1924SCH] T= 298.15 K, 1=0.2-1.1 0.9 HC1O4 pot 1.19 ' [1924SCH] T= 298.15 K, 1=0.2-1.1 1.1 HC1O4 pot 1.13 [1952GAY/GAR] T= 298.15 K, 1=0.01-0.1 var HC1, NaOH sol 1.02 3 [1952GAY/GAR] T= 298.15 K, 1=0.01 0.01 HC1 sol 1.09 3 [1952GAY/GAR] T= 298.15 K, 1=0.02 0.02 HC1 sol 1.16 3 [1952G AY/GAR] T= 298.15 K, 1=0.03 0.03 HC1 sol 1.15 3 [1952G AY/GAR] T= 298.15 K, 1=0.05 0.05 HC1 sol 1.16 3 [1952GAY/GAR] T= 298.15 K, 1=0.075 0.08 HC1 sol 1.15 3 [1952GAY/GAR] T= 298.15 K, 1=0.1 0.1 HC1 sol 4 1.64 '' [1972V AS/SHO] T= 298.15 K, 1=0.3-2.5 0 HC1O4 pot ] 1.61 [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 0.36 HC1O4 pot l 1.55 [1972V AS/SHO] T= 298.15 K, 1=0.3-2.5 0.71 HC1O4 pot 1.30 ' [1972 V AS/SHO] T= 298.15 K, 1=0.3-2.5 1.78 HC1O4 pot 1.34 ' [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 2.14 HC1O4 pot 1.32 ' [1972V AS/SHO] T= 298.15 K, 1=0.3-2.5 2.49 HC1O4 pot 5 ! 0.53 ' [1975HEI/SCH] T= 298.15 K, 1=1.5 1.5 HF, H,SO4 pot log PK4: Sb(OH)3° + H2O <=> Sb(OH)i + H* -12.05 6 [1948TOU/MOU] T= 298.15 K, 1=0.04-2 0 NaOH sol -11.82 [1952GAY/GAR] T= 298.15 K, 1=0.01-0.1 var HC1, NaOH sol -11.83 6 [1952GAY/GAR] T= 298.15 K, 1=0.01-0.1 0 NaOH sol -11.73 4 •! [ 1973 V AS/SHO 1] T= 298.15 K, 1=0.3-2.5 0 NaOH pot 93 JNC TN8400 99-011 Table 5.2: continued -12.10 6 [1973VAS/SHO2] T= 298.15 K, 1=0.3-2.5 0 NaOH sol -11.99 4-7 [1973VAS/SHO2] T= 298.15 K, 1=0.3-2.5 0 NaOH sol -11.91 [1994AKI/ZOT] T=298, I=dil 0 self medium sol 1 estimated from potentiometric data assuming a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions) 2 we assumed, considering the high H+ concentration, that the Sb3+ species is dominant in such a solution 3 O> calculated with log K S3 = 8.48 [1952GAY/GAR] or = 8.30[1974AHR/BOV] 4 linear extrapolation to 1=0 by [1972VAS/SHO] 5 + [1975HEI/SCH] assumed that the Sb(OH)2 species is dominant 6 recalculated in this report from experimental values and extrapolated to 1=0 with SIT. 7 calculated with log K°*S3 = 8.72 (section 5.2.1). Table 5.3: Thermodynamic data for the antimony(III) hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. logpm, 3m+n Reference Comments I (M) Medium + 3+ log /?,,„.• Sb(OH)3° + 3H <=>Sb + 3H2O 0.35 ' H952LAT1 T= 298.15 K, I=n/a u log p,,: Sb(OH)3° + 2H+ <=> SbOH + 2H2O 0.19 ri976SMI/MAR] T= 298.15 K, 1=5 + log pu: Sb(OH)3° + H* <=> Sb(OH)2 + H2O 1.23 ' [1952LAT] T= 298.15 K, I=n/a 0.87 [1957PIT/POU] T= 298.15 K,I=n/a 1.20 [1976SMI/MAR] T= 298.15 K, 1=0 0 1.41 [1976BAE/MES] T= 298.15 K, 1=0 0 1.47 > [1980BEN/TEA] T= 298.15 K, I=n/a 1.41 [1981BAE/MES] T= 298.15 K, 1=0 0 1.47 ' [1982WAG/EVA] T= 298.15 K,I=n/a 1.17 [1985BAB/MAT] T= 298.15 K, 1=0 0 0.90 [1985PAS] T= 298.15 K, 1=0 0 1.18 [1986ITA/NIS] T= 298.15 K, 1=0 0 1.18 [1989SMI/MAR] T= 298.15 K, 1=3 3 1.42 [1992PEA/BER] T= 298.15 K, 1=0 0 log /V Sb(OH)3° + H2O <=> Sb(OH)-+4 H* -10.63 > [1952LAT] T= 298.15 K,I=n/a -10.99 [1957PIT/POU] T= 298.15 K, 1=3-9 var KOH -11.80 [1976SMI/MAR] T= 298.15 K, 1=0 0 -11.82 [1976BAE/MES] T= 298.15 K, 1=0 0 -11.50 ' [1980BEN/TEA] T= 298.15 K,I=n/a -11.82 [1981BAE/MES] T= 298.15 K, 1=0 0 -11.50 l [1982WAG/EVA] T= 298.15 K,I=n/a -11.79 [1985BAB/MAT] T= 298.15 K, 1=0 0 -11.90 [1985PAS] T= 298.15 K, 1=0 0 -11.79 [1986ITA/NIS] T= 298.15 K, 1=0 0 -11.82 [1992PEA/BER] T=298.15K, 1=0 0 + 2 log (324: 2 Sb(OH)3° + 2H <=> Sb2(OH)4 * + 2H2O 2.88 [1976SMI/MAR] T= 298.15 K, 1=5 0 calculated in this report with a A,G of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 94 JNC TN8400 99 -Oil 5.2 Solid antimony(IH)-oxide/hydroxide Antimony sesquioxide, Sb2C>3(cr) (often also the notation Sb4C>6(cr) is used) exists in the orthorhombic form (a-Sb2O3, valentinite) from room temperature to 846 K. Between 846 K and the melting temperature of 929 K, Sb2C>3(cr) exists in the cubic modification ((3-Sb2C>3, senarmontite). While at room temperature valentinite (a-Sb2O3(cr)) precipitates readily from solutions ([1924SCH], [1939BLO], [1984BER/BRE]), the transformation of the cubic senarmontite to the more stable orthorhombic valentinite does not take place. Where no description of the Sb2C>3(cr) is given, it can normally be assumed that the data given refer to a- Sb2O3(cr) (Table 5.6). Many measurements given in the literature are based on emf (electromotive force) measurements at high temperature extrapolated to 298 K, which may reduce the reliability of the data (these data are given in Table 5.5). When Sb2C>3(cr) is heated in the presence of oxygen above 300 °C, Sb2O4(cr) can be formed [1984BER/BRE], [1987PAN/SRE]. Table 5.4: Experimental equilibrium data used for the determination of precipitation of antimony(UI) oxide/hydroxide. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12: 'Comments on selected references'. Method: pol = polarography, pot = potentiometry, sol = solubility, sp = spectrophotometry. Reference Comments Medium Method log K*S3 KM, log K*S3: 2 Sb(OH)3° <=> a-Sb2O3 (valentinite)+ 3H2O 8.57 i [1924SCH] T= 298.15 K, 1=0.23 0.23 HC1O4 pot 8.72 1 [1924SCH] T= 298.15 K, 1=0.49 0.49 HC1O4 pot 8.26 i [1924SCH] T= 298.15 K, 1=1.13 1.13 HC1O4 pot 8.70 [1948TOU/MOU] T= 308 K, I=dil 0 self medium sol 9.29 [1948TOU/MOU] T= 308 K, I=dil 0 NaOH sol 8.48 [1952GAY/GAR] T= 298.15 K, I=dil 0 self medium sol 8.72 i [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 0.36 HC1O4 pot 8.69 ' [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 0.71 HC1O4 pot 8.77 i [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 1.78 HC1O4 pot 8.71 i [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 2.14 HC1O4 pot 8.69 > [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 2.49 HC1O4 pot 8.30 [1974AHR/BOV] T= 298.15 K, 1=5 5 HC1O4 sol 8.56 [1990SHI/ZOT] T=298, I=dil 0 self medium sol logK*S3: 2 Sb(OH)3° <=> P-Sb2O3(senarmontite) + 3H2O 8.55 [1974AHR/BOV] T= 298.15 K, 1=5 5 HC1O4 sol 8.96 [1990SHI/ZOT] T=298, I=dil 0 self medium sol 8.96 [1994AKI/ZOT] T=298, I=dil 0 self medium sol calculated from potentiometric data with a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions) 95 JNC TN8400 99-011 5.2.1 a-Sb2O3 (valentinite) The solubility of a-Sb2C>3(valentinite) in water was determined by [1939BLO, 1948TOU/MOU and 1952GAY/GAR]. They determined log K*S3 values of 8.08 to 8.70 for the reaction 2Sb(OH)3° <=> Sb2O3(valentinite) + 3H2O. The solutions used by [1939BLO] contained 0.01 M HCl. Since they did not measure pH and did not account for interaction with C1-, the solubility of Sb2O3(valentinite) is probably overestimated. A careful examination of the data of [1948TOU/MOU] showed that the solubility a-Sb2O3 decreased slightly (minimum log K*s3 = 9.29) after the addition of small amounts of NaOH, indicating that in the measurements + in water, Sb2O3 was probably in equilibrium with Sb(OH)2 . The potentiometric measurements of [1924SCH and 1972VAS/SHO], as well as the measurements of [1973BEH/ROS, 1974AHR/BOV, and 1990SHI/ZOT] all resulted in log K*S3 values in the range of 8.3 - 8.8. Extrapolation of the data given in Table 5.4 results in: 2Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O log K*°S3 = 8.72 The extrapolation of the experimental results with the SIT term shown in Figure 5.4 gives a AE value of 0.02 which is in good agreement with the expected Ae value of 0 for the uncharged species involved in the reaction 2Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O. 2Sb(OH)3° <=> Sb2O3+3H2O 10.5 -• valentinite 10 - g 9.5 + 9 o co °OO 8.5 j? 8 y = -0.02x + 8.72 7.5 7 • 6.5 6 4 6 i molal Figure 5.4: Plot of log K*S3 + 0D vs. Im for the reaction 2Sb(OH)3° <=> a-Sb2O3 + 3H2O at 25 °C (and 35 °C). The straight line shows the result of the linear regression: Ae = 0.02; log K*°s3 = 8.72. Calculated from data given in Table 5.4. 96 JNC TN8400 99-011 5.2.2 fi-Sb2C>3 (senarmontite) For the (theoretical) formation of senarmontite a log K*°s3 value of 8.96 is given by [1990SHI/ZOT] based on solubility experiments for the isocoulombic reaction 2Sb(OH)3° <=> (3-Sb2C>3(senarmontite) + 3H2O. This value agrees reasonably well with the value of 8.55 determined by [1974AHR/BOV]. As the value determined by [1990SHI/ZOT] for valentinite agrees very well with our calculation it seems reasonable to use for senarmontite the log K*°s3 value of 8.96 given by [ 1990SHI/ZOT]. Table 5.5: Additional experimentally determined equilibrium data compiled for the precipitation of antimony(III) oxide/hydroxide. These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry, sol = solubility. log K*S3 Reference Comments I (M) Medium Method log K*S3: 2 Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O 8.08 ' [1939BLO] T= 298.15 K,I=diluted 0.01 HC1 sol 8.50 2 [1965FRI/VER] T= 298.15 K,I=n/a 0 n/a n/a 3 4 8.78 ' [1972VAS/SHO] T= 298.15 K, 1=0.3-2.5 0 HC1O4 pot 9.08 5>4 [1973BEH/ROS] T= 298.15 K, I=n/a 0 n/a n/a 9.30 3>4 [1973VAS/SHO1] T= 298.15 K, 1=0.3-2.5 0 NaOH pot 6.74 6-4 fl986AZA/PAN] T= 298.15 K,I=n/a n/a emf log K*S3: 2 Sb(OH)3° <=> fi-Sb2O3(senarmontite) + 3H2O 8JJ3 f!939BLO] T=298.15 K, I=diluted 0.01 HC1 sol 1 + measured in presence of 0.01 M HC1, equilibrium with Sb(OH)2 probable 2 as cited in [1968MIS/GUP] 3 linear extrapolation to 1=0 by [1972VAS/SHO]. 4 calculated in this report with a Afi° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 5 extrapolated from AH and AS values determined at 600-700 K 6 extrapolated from = 800 K Table 5.6: Thermodynamic data for the precipitation of antimony(III) oxide/hydroxide taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK* S3 Reference Comments log K*S3: 2 Sb(OH)3° <=> a-Sb2O3(valentinite) + 3H2O 6.38 [1957PIT/POU] T= 298.15 K, I=n/a 6.36 [1957PIT/POU] T= 298.15 K, I=n/a 8.54 [1976SMI/MAR] T= 298.15 K, 1=0 0 8.30 [1976SMI/MAR] T= 298.15 K, 1=5 5 9.10 ' [1980B EN/TEA] T= 298.15 K, I=n/a 97 JNC TN8400 99 - Oil Table 5.6: continued 9.10 I [1982WAG/EVA] T=298 .15 K, I=n/a 8.14 [1985PAS] T=298 .15 K, I=n/a 8.52 [1986ITA/NIS] T=298 .15 K, 1=0 7.66 2 [1994AKI/ZOT] T=298, I=dil log K*S3: 2 Sb(OH)3° <=> Sb2O3(c) + 3H2O) 8.54 • [1952LAT] T= 298.15 K, I=n/a 8.62 l [1954COU] T= 298.15 K, I=n/a 8.62 • [1963WIC/BLO] T= 298.15 K, I=n/a 10.23 ' [1971NAU/RYZ] T= 298.15 K, I=n/a 8.48 [1976BAE/MES] T= 298.15 K, I=n/a 10.78 ' [1977BAR/KNA] T= 298.15 K, I=n/a 9.07 ' [1978ROB/HEM2] T= 298.15 K, I=n/a 10.01 ' [1979KUB/ALC] T= 298.15 K, I=n/a 8.48 [1981BAE/MES] T= 298.15 K, I=n/a 9.10 ' [1982PAN] T= 298.15 K, I=n/a 7.64 [1985BAB/MAT] T= 298.15 K, I=n/a 8.48 [1992PEA/BER] T= 298.15 K, I=n/a T*55.- 2 Sb(OH)3° 3 <=> P-Sb2O3(senarmontite) + 3H2O 9.28 ' [1954COU] T= 298.15 K, I=n/a 7.84 [1957PIT/POU] T= 298.15 K, I=n/a 7.83 [1957PIT/POU] X- 298.15 K, I=n/a 9.28 ' [1963WIC/BLO] T= 298.15 K, I=n/a 11.63 i [1971NAU/RYZ] 298.15 K, I=n/a 8.54 [1976SMI/MAR] T= 298.15 K,I=5 10.42 l [1980BEN/TEA] T= 298.15 K, I=n/a 10.42 i [1982WAG/EVA] T= 298.15 K, I=n/a 9.46 [1985PAS] T= 298.15 K, I=n/a 12.77 3 [1986ITA/NIS1 T= 298.15 K,I=0 log K*S3: Sb(OH)3° <=> Sb(OH)3 (s) 5.13 l [1971NAU/RYZ] T= 298.15 K, I=n/a 7.40 ' [1980BEN/TEA] T= 298.15 K, I=n/a 7.39 ' ri982WAG/EVA1 T= 298.15 K, I=n/a 1 calculated in this report with a A[G° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 2 probably copied wrongly by [ 1994AKI/ZOT] from literature 3 copied wrongly by [198ITA/NIS] from [1982WAG/EVA] 98 JNC TN8400 99-011 5.3 Antimony(HI) chloride system Sb(III) forms complexes and compounds with chloride. Also the existence of mixed Sb(III) chloride hydroxide complexes and compounds is reported. Experimentally determined data are given in Table 5.7 and 5.8. Thermodynamic data from earlier compilations are given in Table 5.9. Table 5.7: Experimentally determined equilibrium data compiled for the Sb(III) chloride system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12. Method: pol = polarography, pot = potentiometry. log Pm.n Reference Comments I (M) Medium Method log Pi, v Sb3+ + Or <=> SbCl2+ 2.26 [1959PAN/DES] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 2.30 [1970BON/WAU] T= 303.15 K, 1=5 5 HC1O4 pol 3+ + log 07,2: Sb + 2CI- t=> SbCl2 3.49 [1959PAN/DES] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 4.00 [1970BON/WAU] T= 303.15 K, 1=5 5 HC1O4 pol 3.39 [1975BIE/ZIE] T= 298.15 K, 1=4 4 HC1O4 pot 3+ log pU: Sb + 3Ct <=> SbCl3° 4.18 [1959PAN/DES] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 5.78 [1970BON/WAU] T= 303.15 K, 1=5 HC1O4 pol 4.09 [1975BIE/ZIE] T= 298.15 K, 1=4 HC1O. pot 3+ log PJ,4: Sb + 4CI- <^ SbCl4- 4.72 [1959PAN/DES] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 6.78 [1970BON/WAU] T= 303. 15 K, 1=5 5 HC1O4 pol 3+ 2 log Pi,s-- Sb + 5CI- <=> SbCl5 - 4.72 [1959PAN/DES] T= 298. 15 K, 1=4 4 NaClO4, HC1O4 pol 3+ 3 log PJ,6: Sb + 6Ct <=> SbCl6 ~ 4.11 [1959PAN/DES] T= 298. 15 K, 1=4 4 NaC104, HC1O4 pol 99 JNC TN8400 99-011 3+ 2+ + 0 2 3 5.3.1 Reactions ofSb : SbCl , SbCl2 , SbCl3 , SbCl4-, SbCl5 ~ and SbCl6 ~ Antimony forms complexes with chloride in acid solutions. [1959PAN/DES, 1970BON/WAU, 1975BIE/ZIE] determined in concentrated HCIO4 solutions the formation of Sb(III) chloride complexes from the reaction of Sb3+ with Ch (Table 5.7). From the experimental evidence it is clear that Sb(III) chloride complexes will be important only in very acid solution, under other 3+ 2+ + conditions, the hydrolyzed of complexes of Sb , i.e., SbOH , Sb(OH)2 and Sb(OH)3°, will be dominant. [1974SHO/MAB] was able to show that the formation of Sb(III) chloride complexes will become important only at pH values < 2. Antimony forms complexes with chloride. [1959PAN/DES, 1970BONAVAU, 1975BIE/ZIE] determined in concentrated HCIO4 solutions the formation of Sb(III) chloride complexes from the reaction of Sb3+ with Ch (Table 5.7). From these data equilibrium constants can be 3+ calculated (Figures 5.5 and 5.6). After conversion from Sb to Sb(OH)3° (cf. Section 5.1.1) the following tentative values can be calculated: 2 Sb(OH)3° +3H+ + Cl- SbCl + + 3H2O log P°u= 2.78 + Sb(OH)3° +3H+ + 2C1- SbCl2 + 3H2O log P°1>2 = 3.27 Sb3++CI"oSbCI2+ 4.5 • 4 - Q 3.5 - CD + 3 4- D) 2 = 0.05x + 3.51 1.5 1 0.5 •• 0 0 2 4 6 8 lmi molal 3+ 2+ Figure 5.5: Plot of log (3U + 6 D vs. Im for the reaction : Sb + Cl~ <=» SbCl at 25 °C. The straight line shows the result of the 'linear regression': Ae = -0.05; log $°\,\ ~ 3.51. Calculated from data given in Table 5.7. 100 JNC TN8400 99-011 Q 3+ + Figure 5.6: Plot of log (3li2 + 10 D vs. Im for the reaction : Sb + 2C1- <=> SbCl2 at 25 °C. The straight line shows the result of the linear regression: Ae = -0.38; log P0]^ = 4.00. Calculated from data given in Table 5.7. No data are recommended in this report for the formation of SbCl30, SbCLr, SbCls2 , and 3 0 2 SbCl6 ~ as the data reported for SbCl3 and for SbCl4- show quite a spread and for SbCl5 - 3 and SbCl6 - not enough data are available to extrapolate the constants to 1=0. The Ae values of -0.05 and -0.38 obtained in Figures 5.5 and 5.6 can be compared to the As of -0.22 and -0.56 2+ 3+ estimated from the NEA americium data (£(AmCl , C1O4-) = 0.39, e(Am , CIO4-) = 0.49, + + and e(H , C1-) = 0.12, e(Am(0H)2 , CIO4-) = 0.17; [1995SIL/BID]). 5.3.2 SbCl4- and SbOHCl3~ A number of scientists, [1953HAI, 1967VAS/YUS, 1968NOR/KAZ, 1984VAS/SHO], measured the redox potential of the reaction Sb(cr) + 4Ch <=> SbCl4- +3e~. Assuming a log (3 for Sb(OH)3°/Sb(cr) of 11.82 (see Section 5.7: Redox reactions), from these data a constant can be estimated for the reaction Sb(OH)3 + 4C1" + 3H+ <=> SbCl4- + 3H2O (Table 5.8). These constants are not very reliable, as none of the authors could show, that SbCl4~ was in fact the dominant species in their experiments or whether other antimony(III) chloride complexes were 3+ 2 + 2 also present (see Section 5.3.1: Reactions of Sb : SbCl +, SbCl2 , SbCl3°, SbCl4~, SbCl5 - 3 and SbCl6 -). Additionally, the solutions of [1953HAI] and [1967VAS/YUS] had a pH around 0 (H+ concentrations < 1.2 M). It is doubtful that really Sb3+ will dominate the antimony(III) speciation at this pH [1977ANT/NEV]. 101 JNC TN8400 99-011 5.3.3 SbCl3(s) and SbOCl(s) (or Sb4O5Cl2(s)) Direct measurements of antimony(III) chloride and oxychloride solubility are not available. Based on the data compiled in Table 5.9, it can be concluded that both SbCl3(s) and SbOCl(s) are easily soluble. [1975HEN/LON, 1997KOL/HEN] showed that recrystallization of SbCl3(s) with water produces Sb^sC^s) and not SbOCl(s). 5.3.4 Additional equilibrium data compiled for the antimony(III) chloride system Table 5.8: Experimentally determined equilibrium data compiled for the antimony(III) chloride hydroxide system. These data were not chosen in the present report for the evaluation of recommended stability values (cf. Section 5.3.2). Method: pol = polarography, pot = potentiometry. log P Reference Comments Medium Method + ,: Sb(OH)3° + 4CI- + 3H » SbCl4- + 3H2O 2.91 [1953HATJ T= 303 K, 1=4 4 HCI pol 4.18 [1953HAI] T= 303 K, 1=1-6 var HCI pol 4.38 [1967VAS/YUS] T= 298.15 K, 1=0.8-1.2 var HCI pol 3.37 [1968NOR/KAZ] T= 303.15 K, 1=5 5 HC1O4 pot 2.66 ,2 [1984VAS/SHO] T= 298.15 K, 1=4-7.5 0 HCI pot 3.90 [1984VAS/SHO] T= 298.15 K, 1=4.0 4 HCI pot 4.71 [1984VAS/SHO] T= 298.15 K, 1=4.7 4.7 HCI pot 5.69 [1984VAS/SHO] T= 298.15 K, 1=5.5 5.5 HCI pot 6.07 [1984VAS/SHO] T= 298.15 K, 1=5.9 5.9 HCI pot 6.32 [1984VAS/SHO] T= 298.15 K, 1=6.3 6.3 HCI pot 6.97 [1984VAS/SHO] T= 298.15 K, 1=6.9 6.9 HCI pot 7.53 f 1984V AS/SHO] T= 298.15 K, 1=7.4 7.4 HCI pot log Pi.i,3: Sb(OH)3° + 3Ct + 2H+ i=> SbOHClf + 2H2O 1.84 ' ri967VAS/YUSl T= 298.15 K, 1=0.6 0.6 HCI pol 1 estimated from potentiometric data assuming a log K (Sb(OH)3°/Sb(cr)) = 11.99 (see Section 5.7: Redox reactions) 2 extrapolated to 1=0 by [1984VAS/SHO] with extended Debye-Huckel term. Table 5.9: Thermodynamic data for the formation of antimony(III) chloride hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. Reference Comments I (M) Medium 3 2 log Pu:Sb + + Ct <=> SbCl * 2.30 [1976SMI/MAR] T= 298.15 K, 1=4 HC1O, 2.30 [1989SMI/MAR1 T= 303.15 K, 1=5 102 JNC TN8400 99-011 Table 5.9: continued 3+ log pli2: Sb + 2CI- <=> SbCl2" 3.50 [1976SMI/MAR] T= 298.15 K, 1=4 4 HC1O. 4.00 R989SMI/MAR1 T= 303.15 K, 1=5 5 3+ log pu: Sb + 3Cl- <=> SbCl3° 4.20 [1976SMI/MAR] T= 298.15 K, 1=4 HC1O, 5.80 P989SMI/MAR1 T= 303.15 K, 1=5 3+ log p1A: Sb + 4CI- <=> SbCl4- 4.70 [1976SMI/MAR] T= 298.15 K, 1=4 4 HC1O4 6.80 [1989SMI/MAR1 T= 303.15 K, 1=5 5 3+ log P,i5: Sb + 5Cl~ <=> Sb 4.70 [1976SMI/MAR] T= 298.15 K, 1=4 HC1O. 3+ 3 log pL6: Sb + 6CI- <=> SbCl6 ~ 4.10 [1976SMI/MAR1 T= 298.15 K, 1=4 HC1O. + log K*S3: Sb(OH)3° + 3Cl~ + 3H <=> SbCl3(s) + 3H2O 4.3 > [1952LAT] T= 298.15 K, I=n/a 4.5 ' [1963WIC/BLO] T= 298.15 K, I=n/a 4.2 ' [1971NAU/RYZ] T= 298.15 K, I=n/a 4.1 l [1979KUB/ALC] T= 298.15 K, I=n/a 4.2 ' [1980BEN/TEA] T= 298.15 K, I=n/a 4.2 ] [1982WAG/EVA] T= 298.15 K, I=n/a 3.9 [1985PAS] T= 298.15 K, I=n/a + log K*S3: Sb(OH)3° + Cl- + H <=> SbOCl(s) + 2H2O 1.8 ' [1977BAR/KNA] T= 298.15 K,I=n/a 1.8 ' ri979KUB/ALCl T= 298.15 K, I=n/a calculated in this report with a AfG° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 103 JNC TN8400 99-011 5.4 Antimony(IH) fluoride system Antimony(III) forms complexes and compounds with fluoride. Experimentally determined data used in this report are given in Table 5.10. Further constants are compiled in Tables 5.11 and 5.12. Table 5.10: Experimentally determined equilibrium data compiled for the antimony(III) fluoride system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12. Method: pol = polarography, pot = potentiometry. Reference Comments I (M) Medium Pm,n,o Method 3+ 2+ + log Ku: Sb + HF& SbF + H 3.00 [1970BON] T= 303 K, 1=2 2 Na, HC1O4 pol 3+ + log Klt2: Sb + 2HF <=> SbF2 + 2H+ 5.70 [1970BON] T= 303 K, 1=2 2 Na, HC104 pol 3+ + log KJJ: Sb + 3HF <=> SbF3° + 3H 8.30 [1970BON] T= 303 K, 1=2 2 Na, HC1O4 pol + log Kh4: SbF3° SbF4~ + H 2.65 [1970BON] T= 303 K, 1=2 2 Na, HC1O4 pol 3+ 2+ + 5.4.1 Reactions ofSb : SbF , SbF2 , SbF30, and SbF4~ Antimony(in) forms weak complexes with fluoride. From the data determined by [1970BON] at 1=2 (Table 5.10) the tentative formation constants for Sb(III) fluoride complexes can be derived with the SIT equation. The Ae values of 0.04 and -0.04 are estimated from the NEA 3+ 2+ + americium data (£(Am , C1O4-) = 0.49, e(AmF , CIO4-) = 0.39, £(AmF2 , CIO4-) = 0.17, and e(H+, CIO) = 0.14; [1995SIL/BID]). Sb3+ + HF SbF2+ log K°u = 4.03, Ae = 0.04 3+ Sb + 2HF log K°i,2 = 7.02, Ae = -0.04 3+ Sb + 3HF SbF30 log K°i,3 = 9.55, Ae =-0.07 104 JNC TN8400 99-011 Correction of these values with a log Ka of 3.18 for F" + H+ o HF [1992GRE/FUG] and for + 3+ the reaction Sb(OH)3° + 3H <=> Sb + 3H2O (Section 5.1.1) gives the following tentative values: 2 Sb(OH)3°+3H++ F- <=> SbF + + 3H2O log p°u = 6.48 Sb(OH)3o +3H+ + 2F- <=> SbF2+ + 3H2O log (3\2 = 12.65 Sb(OH)3° +3H+ + 3F- «• SbF3° + 3H2O log p\3 = 18.36 No data is recommended in this report for the formation of SbF^. 5.4.2 Reactions ofSbF3°: SbF4~, and SbF3OH- Recently, [1993DEL/MIL] determined constants for the formation of SbF^ from SbF3° and for the hydrolysis of SbF3° to SbF3OH~ (Table 5.11). Unfortunately, their experiments were carried out at varying ionic strength. Extrapolation to 1=0 is therefore not possible. 5.4.3 SbOF°orSb(OH)2F° 0 In Table 5.12 data are compiled for the formation of SbOF or Sb(OH)2F°. Direct experimental measurements of these constants, however, are not available. 5.4.4 SbF3(s) Direct measurements of the solubility of SbF3(s) are not available. Based on the data compiled in Table 5.12 it can be concluded that SbF3(s) is quite soluble. Table 5.11: Additional, experimentally determined equilibrium data compiled for Sb(III) fluoride system. These data were not selected in the present report for the evaluation of recommended stability values. Method: NMR = 19F NMR, tit = titration (pH). log Pm,n,0 Reference Comments I (M) Medium Method log K, ,: Sb3+ + HF <=> SbF2+ + H+ 3.11 ' [1975HEI/SCH] T= 298.15 K, 1=1.5 HF, H,SO4 pot 3+ + log K12: Sb + 2HF <=> SbF2* + 2H 5.11 ' [1975HEI/SCH] T= 298.15 K, 1=1.5 HF, H7SO4 p_ot_ 3+ + log Ku: Sb + 3HF <=> SbFf + 3H 6.23 ' [1975HEI/SCH1 T= 298.15 K, 1=1.5 HF, H,SO4 pot^ 105 JNC TN8400 99-011 Table 5.11: continued + log K14: SbF3° + HF <=> SbFf + H 1.72 2 [1993DEL/MIL] T= 298 K, 1=0.03-1 var self medium NMR + log KUJ: SbF3° + H2O <=> SbF3OH~ + H 2 3.92 [1993DEL/MIL1 T= 298 K, 1=0.03-0.1 var self medium tit ' exact I not known. 2 in this report the presence of HF (instead of F~ as by [1993DEL/MIL]) is assumed, as measurements are carried out at pH 1.6 - 2.8). I not constant. Table 5.12: Thermodynamic data for the Sb(III) fluoride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log Pm.n.o Reference Comments KM;i Medium 3+ 2+ + log K,j: Sb + HF <=> SbF + H 3.00 [1980BON/HEF1 T= 298.15 K, 1=2 2 Na, HC1O, 3+ + + log Ku: Sb + 2HF <=> SbF2 •+• 2H 5.70 [1980BON/HEF1 T= 298.15 K, 1=2 2 Na, HC1O, 3+ + log Ku: Sb + 3HF <=> SbF3° -f 3H 8.30 [1980BON/HEF] T= 298.15 K, 1=2 2 Na, HC1O4 + log KK4: SbF3° + HF<=> SbF4- -¥ H 2.70 [1980BON/HEF1 T= 298.15 K, 1=2 2 Na, HC1O4 + 0 Sb(OH)3° H SbOF + 2//2O 6.51 ' [1982WAG/EVA] T= 298.15 K,I=n/a 6.19 [1985PAS1 T= 298.15 K,I=n/a + log jij : Sb(OH)3° H <=> Sb(OH)2F° + H2O 6.52 ' [1980BEN/TEA] T= 298.15 K, I=n/a 6.52 ' [1982WAG/EVA] T= 298.15 K,I=n/a 6.21 [1985PAS1 T= 298.15 K,I=n/a + log K'S3: Sb(OH)3° + 3F'+ 3H <=> SbF3(s) + 3H2O 10.48 ' [1952LAT] T= 298.15 K, I=n/a 12.96 ' [1963WIC/BLO] T= 298.15 K,I=n/a 12.78 ' [1977BAR/KNA] T= 298.15 K, I=n/a 12.78 ' [1979KUB/ALC] T= 298.15 K,I=n/a 10.17 [1985PAS1 T= 298.15 K, I=n/a calculated in this report with a Afi" of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 106 JNC TN8400 99-011 5.5 Sb(III) sulfate system Antimony(III) is reported to form complexes and compounds with sulfate. Experimental data for complex formation of Sb(III) with sulfate were determined by [1970DAWAVTL] in concentrated H2SO4 solutions (Table 5.13). Additional constants are compiled in Table 5.14. 5.5.1 SbOSO4~ [1970DAWAVTL] proposed the formation of SbOSO4- in 2 - 6 M H2SO4 solutions and of Sb(SO4)2~ in 14 - 16 M H2SO4 solutions. For the calculation of formation constants (Table 5.13) he assumed in 0.2 - 12 M H2SO4 solutions the presence of SbO+, which seems quite doubtful in the light of the results discussed in Section 5.1.1 and 5.1.2 of this report. No formation constants for the Sb(III) sulfate complexes are proposed in this report. Such complexes, however, will be important only in concentrated sulfate solutions [1970DAW/WIL]. 5.5.2 Sb2(SO4)3(s) Thermodynamic values for Sb2(SC>4)3(s) were compiled by [1977BAR/KNA] and [1979KUB/ALC]. Solubility products calculated in this report from these thermodynamic data are shown in Table 5.14. From these data, Sb2(SC>4)3(s) is expected to be very soluble. No direct experimental determination of Sb2(SC>4)3(s) solubility is available. Table 5.13: Experimentally determined equilibrium data compiled for Sb(III) sulfate system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references see text. Method: cat = cation exchange. l°g Pm.n Reference Comments I (M) Medium Method + 2 log pij: SbO + SO4 ~ <=> SbOSOf 0.30 ri970DAW/W]L1 T= 298.15 K, 1=1-4 var H,SO^ cal_ Table 5.14: Thermodynamic data for the Sb(III) sulfate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log Pmno Reference Comments 2 + log K'S3: 2Sb(OH)3° + 3SO4 - + 6H <=> Sb2(SO4)3(s) -5.0 ' [1977BAR/KNA] T= 298.15 K, I=n/a -5.0 ' [1979KUB/ALC1 T= 298.15 K, I=n/a 1 calculated in this report with a A(G° of -643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 107 JNC TN8400 99-011 5.6 Antimony(III) sulfide system Antimony(III) forms complexes with sulfide and also forms a stable salt, stibnite (Sb2S3(cr)). [1988KRU] and [1994AKI/ZOT] studied complex formation with sulfide from measurements at 25 - 300 °C. [1994AKI/ZOT] included also the results determined by [1988KRU] and put up a consistent thermodynamic data set for the Sb(III)-S(II)-OH system. Equilibrium constants are compiled in Table 5.15. Table 5.15: Experimentally determined equilibrium data compiled for the Sb(III) sulfide system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12: 'Comments on selected references'. Method: sol = solubility measurements. log Pm,n,o Reference Comments I(M) Medium Method : 2 log Po,2J 2 Sb(OH)3° +4HS- + 2H+ <=> Sb2S 42.55 1 [1988KRU] T=298, I=dil 0 H2S sol 42.51 2 [1994AKI/ZOT] T=298, I=dil 0 H,S sol + log Pi,2,4.- 2 Sb(OH)3° + 4HS- + 3H <=> HSbtfr + 6H2O 52.07 i [1988KRU] T=298, I=dil 0 H2S sol 52.29 2 [1994AKI/ZOT] T=298, I=dil 0 H,S sol : log /32,2,4 2 Sb(OH)3° + 4HS- + 4H+ <=> H2Sb2S4° + 6H2O ] 56.98 [1988KRU] T=298, I=dil 0 H2S sol 57.02 2 [1994AKI/ZOT] T=298, I=dil 0 H,S sol + log K*S3: 2Sb(OH)3° + 3HS- + 3H <=> Sb2S3(stibnite) + 6H2O 55.14 2 [1994AKI/ZOT] T=298, I=dil 0 H,S sol 1 extrapolated to 1=0 with Debye-Hiickel by [1988KRU]; calculated with a log K*so for stibnite of 55.14. 2 extrapolated to 1=0 with Debye-Hiickel by [1994AKI/ZOT]; calculated from Afi° given by [1994AKI/ZOT]. 2 5.6.1 Sb2S4 ', HSb2S4-, and H2Sb2S4° Sb(III) forms complexes with sulfide (Table 5.15 and 5.16). As the determination of the Sb(III) sulfide complex formation constants at 1=0 by [1988KRU] and [1994AKI/ZOT] are the only 108 JNC TN8400 99-011 experimental value reported in the literature. The mean of these log values (Table 5.15) are recommended in this report as tentative values: 2 42 53 2Sb(OH)3 + 4HS-+ 2H+ <=> Sb2S4 -+6H2O log p°0,2,4 = - 2Sb(OH)3 + 4HS- + 3H+ HSb2S4-+6H2O log P°i'2'4= 52.18 2Sb(OH)3 + 4HS-+ 4H+ H2Sb2S4° + 6H2O log p°2^4 = 57.00 5.6.2 Sb2S3(stibnite) Stibnite, Sb2S3(cr), is quite insoluble, as indicated by the values compiled in Table 5.15 and 5.16. The log K*°s3 value determined by [1994AKI/ZOT] is recommended as tentative value: 2Sb(OH)3° + 3HS-+ 3H+ Sb2S3(stibnite) + 6H2O logK*°S3 =55.14 Table 5.16: Thermodynamic data for the Sb(II) sulfide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log rso Reference Comments KM) log Po,2,4-2Sb(OH)3° +4HS- + 2W <=> Sb^/- + 6H2O 50.0 ' [1980BEN/TEA] T= 298.15 K, 1=0 0 50.0 > [1982WAG/EVA] T= 298.15 K, 1=0 0 49.2 [1985PAS] T= 298.15 K, 1=0 0 49.4 [1986ITA/NIS] T= 298.15 K, 1=0 0 l + Og Po,l,2-Sb(OH)3° + 2HS- + H <^> SbS2- + 3H2O - 25.8 ' [1952LAT1 T= 298.15 K, 1=0 0 + log K*S3: 2Sb(OH)3° + 3HS- + 3H « Sb2S3(stibnite) + 6H2O 66.05 ' [1974MIL] T= 298.15 K, I=n/a 60.19 [1985PAS] T= 298.15 K,I=n/a 60.25 [1986ITA/NIS] T= 298.15 K, 1=0 0 56.68 ' ri992SEAl T= 298, I=n/a + log K*S3: 2Sb(OH)3° + 3HS- + 3H <=> Sb2S3(s) +6H2O 3.07 2- [1971NAU/RYZ] T= 298.15 K,I=n/a 66.05 [1977BAR/KNA] T= 298.15 K,I=n/a 60.80 [1978ROB/HEM2] T= 298.15 K,I=n/a 66.05 [1979KUB/ALC] T= 298.15 K, I=n/a 60.80 [1980BEN/TEA] T= 298.15 K,I=n/a 55.02 [1981JOH/PAP] T= 298.15 K,I=n/a 60.82 ' [1982WAG/EVA1 T= 298.15 K, I=n/a + log K'S3: 2Sb(OH)3° + 3HS- + 3H i=> Sb2S3(am, orange) + 6H2O 53.9 ' [1952LAT1 T= 298.15 K, I=n/a 1 calculated in this report with a AfG° of-643.0 kJ/mol for Sb(OH)3° (section 5.7.1) 2 difference to others values probably due to missing minus sign in the Afi° of 156.1 kJ/mol [1971NAU/RYZ]. 109 JNC TN8400 99-011 5. 7 Redox reactions Antimony exists in the oxidation states -III, 0, +III and +V [1976BAE/MES, 1985PAS, 1995WIB]. The -III oxidation state is thermodynamically not stable [1952LAT, 1985PAS] in presence of aqueous solutions. In water, aqueous species of Sb(III) and Sb(V) are stable. Experimentally determined equilibrium data for the redox potential of antimony chosen in this report for the calculation of equilibrium constants at 1=0 are given in Table 5.17. Additional data are compiled in Table 5.18 and 5.19. Table 5.17: Experimentally determined equilibrium data compiled for the redox potential of antimony. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 5.12: 'Comments on selected references'. Method: pot = potentiometry, sol = solubility measurements. logK Reference Comments I (M) Medium Method + + log K: Sb(OH)2 + 2H + 3e- ^Sb(cr) + 2H2O 10.38 [1972VAS/SHO] T= 298.15 K, 1=0.36 0.36 HCIO4 pot 10.44 [1972VAS/SHO] T= 298.15 K, 1=0.71 0.71 HC1O4 pot 10.68 [1972VAS/SHO] rp 298.15 K, 1=1.8 1.78 HCIO4 pot 10.64 [1972VAS/SHO] T= 298.15 K, 1=2.1 2.14 HCIO4 pot 10.67 [1972VAS/SHO] T= 298.15 K, 1=2.5 2.49 HCIO4 pot log K: Sb(OH)3° + 3e~ + 3H+ 4=> Sb(cr) + 3H2O 11.88 [1990SHI/ZOT] T=298, I=dil 0 self medium sol 11.83 [1994AKI/ZOT] T=298, I=dil 0 self medium sol log K: Sb(OH)6- + 2e~ <^Sb(OH)4~+ 20H~ -16.36 [1923GRU/SCH] T= 293 K, 1=4 4 KOH pot -17.44 [1923GRU/SCH] T= 293 K, 1=5 5 KOH pot -18.22 [1923GRU/SCH] T= 293 K, 1=6 6 KOH pot -18.73 [1923GRU/SCH] 293 K, 1=7 7 KOH pot -18.97 [1923GRU/SCH] T= 293 K, 1=7.5 7.5 KOH pot -19.20 [1923GRU/SCH] T= 293 K, 1=8 8 KOH pot -19.54 [1923GRU/SCH] 293 K, 1=9 9 KOH pot -19.91 [1923GRU/SCH] T= 293 K,1=10 10 KOH pot 110 JNC TN8400 99-011 Table 5.17: continued -19.54 [1953HAI] T=300K, 1=10 M 10 KOH pot -17.68 [1953HAI] T= 300 K, 1=6 M 6 KOH pot -16.36 [1953HAI] T= 300 K, 1=3 M 3 KOH pot -14.98 [1953HAI] T= 300 K, 1=1 M 1 KOH pot 5.7.1 Sb(cr)/Sb(OH)3° The redox couple Sb(cr)/Sb(OH)3° is not well investigated. Sb(cr) has a hexagonal rhombohedral structure and has a gray color [1995WIB]. [1924SCH] showed that in 0.2 to 1.1 M HC1O4 solutions Sb(OH)2+ dominates the speciation. Also the hydrolysis constants calculated in Section 5.1 indicate that in this pH region Sb(OH)2+ dominates the speciation. [1972VAS/SHO] determined with potentiometric measurements under acidic conditions (Table + 5.17) constants for the reaction Sb(OH)2 + 2H+ + 3e- <=> Sb(cr) + 2H2O. The values of [1972VAS/SHO] were extrapolated to I = 0 in Figure 5.7 giving a log K° of + 10.83. Using the log (3°li2 of 1.30 (Section 5.1.3: Sb(OH)2 ), a log K° of 12.13 for the reaction Sb(OH)3° + 3H+ + 3e- <=> Sb(cr) + 3H2O is obtained. + + Sb(OH)2 +3e+2H <^>Sb(cr)+2H2O 13 lm, molal + Figure 5.7: Plot of log K + 3 D vs. Im for the reaction : Sb(OH)2 + 2H+ + 3e- «=> Sb(cr) + 2H2O at 25 °C. The straight line shows the result of the linear regression: Ae = -0.21; log K° = 10.83. Calculated from data given in Table 5.17. Ill JNC TN8400 99-011 [1990SHI/ZOT and 1994AKI/ZOT] determined constants for the reaction Sb(OH)3° + 3H+ + 3e- o Sb(cr) + 3H2O with solubility measurements (Table 5.17) and extrapolated these values to 1=0 with the Debye-Hiickel equation. This value shows a good agreement with the potentiometric data of [1972VAS/SHO]. The mean of the measurements of [1972VAS/SHO and 1990SHI/ZOT, 1994AKI/ZOT] is: Sb(OH)3° + 3H+ + 3e- Sb(cr) + 3H2O logK° = 11.99 0.236 V resulting in a AfG° of -643.0 kJ/mol for Sb(OH)3°. 5.7.2 Sb(III)/Sb(V) The equilibrium between Sb(V) and Sb(III) has been studied electrochemically in KOH solutions by [1923GRU/SCH] and [1953HAI]. The log K values calculated in this report from the experimentally determined redox potentials are listed in Table 5.17 (log K = 2xEmeaSured[V]/ 0.05916). Extrapolation to 1=0 is shown in Figure 5.8 and gives a log Kb° value of -15.36 for b the reaction Sb(OH)6- + 2e- <=> Sb(OH)4~ + 2OH~. This value is similar to the mean log K value of -15.74 (Table 5.19) calculated by [1957PIT/POU] based on the data measured by [1923GRU/SCH]. No other data for the redox equilibria between Sb(III) and Sb(V) have been found in the literature. Sb(OH)6-+2e-oSb(OH)4-+20H- g> -18 + -19 •• -20 -m -21 5 10 15 lmi molal b Figure 5.8: Plot of log K - 2 D vs. Im for the reaction: Sb(OH)6- + 2e~ <=> Sb(OH)4- + 2OH~ at 25 °C. The straight line shows the result of the linear regression: Ae = 0.38; log Kb° = -15.36. Calculated from data given in Table 5.17. 112 JNC TN8400 99-011 b Using the log K ° of-15.36 (Figure 5.10), a log Kw of- 14.00, a log p°1>4 of -11.93 (for the reaction Sb(OH)3° + H2O <=> Sb(OH)4- + H+; see Section 5.1.4) and a log p°lt6 of 2.72 (for + the reaction Sb(OH)5° + H2O <=> Sb(OH)6- + H ; see Section 5.8), the following constant for the redox equilibria between Sb(OH)3° and Sb(OH)5° can be calculated: Sb(OH)5° + 2H++ 2e- Sb(OH)3° + 2H2O logK° = 21.84 E> = 0.646 V 5.7.3 Sb(cr)/Sb(OH)5° From the data obtained in Section 5.7.1 and 5.7.2 also a log K° value for the theoretical equilibrium between Sb(cr)/Sb(OH)5° can be calculated: Sb(OH)5°+5 H++5e- <=> Sb(cr) + 5H2O logK° =33.83 F = 0.400 V corresponding to a AfG° of -992.62 kJ/mol for Sb(OH)5°. 5. 7.4 Additional data compiled for the redox potential of antimony Table 5.18: Additional, experimentally determined equilibrium data compiled for the redox potential of antimony. These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry. logp Reference Comments I (M) Medium Method + + log K: Sb(OH)2 + 2H + 3e- <=> Sb(cr) + 2H2O 10.60 > [1924SCH] T= 298.15 K, 1=0.2 0.2 HC1O4 pot 10.80 ' [1924SCH] T= 298.15 K, 1=0.5 0.5 HCIO4 pot 10.80 ' [1924SCH] T= 298.15 K, 1=0.9 0.9 HCIO4 pot 10.80 > [1924SCH] T= 298.15 K, 1=1.1 1.1 HC1O4 pot 10.69 2 [1924SCH] T= 298.15 K, 1=0.2 0.2 HCIO4 pot 10.82 2 [1924SCH] T= 298.15 K, 1=0.5 0.5 HCIO4 pot 2 11.03 [1924SCH] T= 298.15 K, 1=1.1 1.1 HC1O4 pot 10.34 3 [1972VAS/SHO1 T= 298.15 K, 1=0.3-2.5 0 HC1O, pot 13 JNC TN8400 99-011 Table 5.18: continued log K: Sb(OH)6- + 2e~ <=> Sb(OH)f + 2OH~ 4 -14.47 [1923GRU/SCH] T= 293 K, 1=3 3 KOH pot 5 -15.19 [1953HAI] T= 300 K, 1=1-10 M 0 KOH pot 5 -15.31 [1953HAI] T= 303 K, 1=1-10 M 0 KOH pot 1 values corrected for H+ activity by [1924SCH] 2 + calculated in this report from E° values for the reaction Sb/Sb2O3(cr) as given by [1924SCH] and Sb(III) and H concentrations measured by [1924SCH] in presence of Sb2O3(cr). 3 linear extrapolation to 1=0 by [1972VAS/SHO] 4 this value is very different from the other values and it is not clear if the redox potential is -0.428 or -0.488 V. This value is therefore not chosen in this report 5 as given by [1953HAI]; from data at variable ionic strength Table 5.19: Thermodynamic data for the redox potential of antimony taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions this is indicated. logP Reference Comments KM) Medium + log K: Sb(OH)3° + 3e~ + 3H <=> Sb(cr) + 3H2O 11.26 [1985B AS/MAT] T=298.15,I=0 0 11.67 [1985PAS] T=298.15,I=0 0 11.70 ri986ITA/NIS] T=298.15,I=0 0 log K: Sb(OH)6-+ 2e~ <=>Sb(OH)f + 2OH- -15.72 1 [1957PIT/POU] T= 293 K, 1=3-9 M, 0 KOH -13.49 2-3 [1957PIT/POU] T= 293 K, 1=3-9 M 3 KOH -14.74 2'3 [1957PIT/POU] T= 293 K, 1=3-9 M 4 KOH -15.38 2 [1957PIT/POU] T= 293 K, 1=3-9 M 5 KOH -15.79 2 [1957PIT/POU] T= 293 K, 1=3-9 M 6 KOH -15.86 2 [1957PIT/POU] T= 293 K, 1=3-9 M 7 KOH -15.86 2 [1957PIT/POU] T= 293 K, 1=3-9 M 7.5 KOH -15.86 2 [1957PIT/POU] T= 293 K, 1=3-9 M 8 KOH -15.69 2 [1957PIT/POU1 T= 293 K, 1=3-9 M 9 KOH 1 as calculated by [1957PIT/POU]; mean of data after correction for activity of OH~ 2 data given by [1957PIT/POU], corrected for activity of OH~ (original data from [1923GRU/SCH]) 3 these data were not chosen by [1957PIT/POU] for the calculation of a mean value 114 JNC TN8400 99- Oil 5.8 Hydrolysis of antimony(V) [1963LEF/MAR] studied the hydrolysis of Sb(V) in 0.5 M (CH3)4NH4C1 solutions and Sb(V) concentrations from 0.0013 to 0.346 M. They interpreted their data in terms of a sequence of dodecamers Hi2_y[Sb(OH)6]i2y- (Table 5.21). These data were recalculated later by [1976BAE/MES] (Table 5.20). [1948TOU/MOU] determined the solubility of Sb2O5(precip) in dilute acid. Their results in 4 3 diluted acids are consistent with the existence of Sb^OH^ " (or Sbi2(OH)63 -) as proposed 3+ by [1963LEF/MAR] and [1976BAE/MES]. In more acidic solutions (pH < - 0.25) Sb(OH)2 (or a polynuclear compound having the same Sb(V)/H+ ratio) seems to dominate the solution ([1948TOU/MOU]). Also the observation of [1975ALY/ABD] indicate further hydrolysis of Sb(V) at pH < - 0.3. The data determined by [1963LEF/MAR] are the only experimental data found in the literature. These data were recalculated by [1976BAE/MES] and its use is recommended in this report: + Sb(OH)5° + H20 «- Sb(OH)6- + H log P°,,6 = -2.72 4 + 12Sb(OH)5° + 4H2O <=> Sbi2(OH)64 -+ 4H log B°1264 = 20.34 5 + 12Sb(OH)5° + 5H2O « Sb12(OH)65 -+ 5H log P°lMS = 16.72 + 12Sb(OH)5° + 6H2O « Sb12(OH)666-+ 6H log P°12,66 = 11.89 + 12Sb(OH)5° + 7H2O <=> Sb12(OH)677-+ 7H log P°,,67 = 6.07 Table 5.20: Equilibrium data used for the hydrolysis of antimony(V). These data were chosen as recommended values in the present report. Method: cal = calculated by [1976BAE/MES] based on experimental data reported by [1963LEF/MAR] log R Reference Comments I(M) Medium Method & r'm,5m+ + log P1A: Sb(OH)5° +H2O <=> Sb(OH)6- + H -2.47 i [1976BAE/MES] T= 298.15 K, 1=0.5 0.5 (CH3)4NH4C1 cal -2.72 i [1976BAE/MES] T= 298.15 K, 1=0 0 (CH,)4NH4C1 cal + log Pl2,64: 12Sb(OH)5° + 4H2,0 <=> Sb12(OH)6/- +4H 23.06 ' [1976BAE/MES] T= 298.15 K, 1=0.5 0.5 (CH3)4NH4C1 cal 20.34 1 [1976BAE/MES] T= 298.15 K, 1=0 0 (CH,)4NH4C1 cal log pi2,65- 12Sb(OH)5° + 5H2O <^> Sb!2(OH)655- + 16.72' [1976BAE/MES] T= 298.15 K, 1=0 0 (CHQ4NH4C1 cal 115 JNC TN8400 99-011 Table 5.20: continued 6 + log Pj2,66- 12Sb(OH)5° + 6H2O <=> SbirfOH)^ - + 6H l 11.89 [1976BAE/MES] T= 298.15 K, 1=0 0 (CH,)4NH4C1 cal + log Pi2,67- 12Sb(OH)5° + 7H2O & Sb}2(OH)677- 7H 6.07 l [1976BAE/MES] T= 298.15 K, 1=0 0 (CH,)4NH4C1 cal recalculated by [1976BAE/MES] from data of [1963LEF/MAR] Table 5.21: Experimentally determined equilibrium data compiled for the hydrolysis of antimony(V). These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references see text. Method: tit = pH titration. log Pm,5m+n Reference Comments I (M) Medium Method log j5L6: Sb(OH)5° + H2O <=> Sb(OH)6~ + H* -2.55 ' ri963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CHQ.NH.C1 tit 3 + log P12,63: 12Sb(OH)5° + 3H2O O Sb,2(OH)63 ~ + 3H 24.45 • [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,).NHdCl tit log j3!2M: 12Sb(OH)5° + 4H2O <=> Sb12(OH)6/- + 4H* 22.90 ' fl963LEF/MARl T= 298.15 K, 1=0.5 0.5 tit 5 + log Pn,65-- 12Sb(OH)s° + 5H2O <=> Sb12(OH)65 - + 5H 19.95 ' [1963LEF/MAR1 T= 298.15 K, 1=0.5 0.5 (CH,).NH4C1 tit 6 + log pil66: 12Sb(OH)5° + 6H2O « Sb12(OH)66 ~ + 6H 15.80 • [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,).NH4C1 tit 7 + log p,l67: 12Sb(OH)5° + 7H2O Sb12(OH)67 ~ + 7H 9.85 ' [1963LEF/MAR1 T= 298.15 K, 1=0.5 0.5 (CHQ.NH.Cl tit s 'OS P 12.68- 12Sb(OH)5° + 8H2O <=> Sb12(OH)68 - + 8H+ ^/70> [1963LEF/MAR] T= 298.15 K, 1=0.5 0.5 (CH,)aNHaCl tit Sb = 1.3 - 346 mM; later recalculated by [1976BAE/MES], see Table 5.20 116 JNC TN8400 99-011 Table 5.22: Thermodynamic data for hydrolysis of antimony(V) taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions this is indicated. log Pm,5m+n Reference Comments I (M) Medium + log KIA: Sb(OH)6- + 2H+ <=> Sb(OH)4 + 2H2O -0.54 ! [1957PIT/POU] T= 308 K, 1=0.05-5 HC1 -0.55 2 ri986ITA7NISl T= 298.15 K, I=n/a + log fi,,6: Sb(OH)5° + H2O <=> Sb(OH)6~ + H -2.47 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 (CH3)4NH4C1 -2.72 [1976SMI/MAR] T= 298.15 K, 1=0 0 4 log P)2.64: 12Sb(OH)5° + 4H2O <=* Sb12(OH)64 - + 4H+ 23.06 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 (CH3)4NH4C1 20.34 [1976SMI/MAR1 T= 298.15 K, 1=0 0 + log P,2,65: 12Sb(OH)5° + 5H2O 6 log P12.66: 12Sb(OH)5° + 6H2O <=* Sb!2(OH)66 ' + 67T 11.89 3 [1976SMI/MAR1 T= 298.15 K, 1=0 0 7 + 7: 12Sb(OH)5° + 7H2O <=* Sb12(OH)67 ~ + 7H 6.07 3 [1976SMI/MAR] T= 298.15 K, 1=0 0 1 calculated by [1957P1T/POU] from the data of [1948TOU/MOU] 2 data from [1957PIT/POU] 3 these values are cited wrongly in the NEA database 117 JNC TN8400 99-011 5.9 Sb2O5(precip) The only reported solubility measurements for Sb2O5(precip) is that of [1948TOU/MOU] at 35 °C in water and dilute acid. [1948TOU/MOU] reported a solubility of 2.71xlO"4 M Sb(V) in water. [1976BAE/MES] calculated later from these data (Table 5.23): 2Sb(OH)5° <=> Sb2O5(precip) + 5H2O log K*°so =7.40 [1976BAE/MES] state that in view of the ease with which gels are prepared from antimonic acid solutions, it is not likely that a pure phase was present in the experiments of [1948TOU/MOU]. [1984BER/BRE] find that the precipitation product obtained from the hydrolysis of Sb(V) in dilute HC1 is amorphous to X-ray detection up to 650 °C. Table 5.23: Equilibrium data determined for Sb2Os(precip). These data were chosen as recommended values in the present report. Method: cal = calculated by [1976BAE/MES] based on solubility data reported by [1948TOU/MOU] (Table 5.24). a Reference Comments I (M) Medium Method Pm.i,5m+5 n log K*so: 2Sb(OH)5° <^> Sb2O5(precip) + 5H2O 7.40 l [1976BAE/MES] T= 308 K, 1=0 0 water cal_ 1 recalculated by [1976BAE/MES] from data of [1948TOU/MOU] Table 5.24: Experimentally determined equilibrium data compiled for Sb2O5(precip). These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility. l°g Pm,5m+n Reference Comments I (M) Medium Method log K*so: 2Sb(OH)5° <=> Sb2O5(precip) + 5H2O 7.13 ' [ 1948TOU/MOU1 T= 308 K, I=diluted self medium sol ' approximate value; equilibrium with Sb(OH)5° assumed; calculated in this report based on the solubility of Sb in water = 2.71X10"4 M 118 JNC TN8400 99-011 Table 5.25: Thermodynamic data for Sb2O5(s) taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log pmi5m+n Reference Comments I (M) Medium log K*so: 2Sb(OH)5° <=> Sb2Os(s) + 5H2O 6.88 ' [1952LAT] T= 298.15 K,I=n/a 3.22 ' [1954COU] T= 298.15 K,I=n/a 3.22 ' [1963WIC/BLO] T= 298.15 K,I=n/a 11.41 ' [1971NAU/RYZ] T= 298.15 K, I=n/a 11.43 ' [1977BAR/BCNA] T= 298.15 K,L=n/a 11.42 ' [1979KUB/ALC] T= 298.15 K, I=n/a 5.20 ' [1980BEN/TEA] T= 298.15 K,I=n/a 5.19 ' [1982PAN] T= 298.15 K,I=n/a 5.20 ' [1982WAG/EVA] T= 298.15 K, I=n/a 5.20 ' [1985BAB/MAT] T= 298.15 K, I=n/a 5.21 ' [1985PAS] T= 298.15 K, I=n/a 5.21 ' [1986ITA/NIS] T= 298.15 K,I=n/a + log K*S4: 2Sb(OH)4 <^> Sb2O5(precip) + 4H2O + 2W 8.32 2 [1957PIT/POU] T= 308 K, 1=0.05-5 HC1 1 all these values were calculated assuming a log K (Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3) 2 calculated by [1957PIT/POU] from the data of [1948TOU/MOU] 119 JNC TN8400 99-011 5.10 Sb2O4(cr) and Sb6O13(cr) When Sb2C>3(cr) is heated in the presence of oxygen above 300 °C, orthogonal Sb2C>4(cr) is formed [1971HEG/BAK, 1984BER/BRE, 1987PAN/SRE, 1997KOL/HEN]. Orthorhombic Sb2O4(cr) is generally assumed to be a mixed compound of Sb2C>3(cr) and Sb2C>5(s) [1971HEG/BAK, 1976BAE/MES, 1995WIB]. [1984BER/BRE] observed in nitric acid the formation of orthorhombic oc-Sb2C>3(cr) at room temperature and the formation of Sb2C>4(cr) above 360 °C in air from powdered Sb(cr). [1984BER/BRE] also showed that the precipitation product obtained from Sb(III) and Sb(V) solutions under acidic conditions is different, indicating a kinetic barrier for the equilibria between Sb(III) and Sb(V) under acidic conditions. [1984BER/BRE] found that the precipitation product obtained from the hydrolysis of Sb(V) in dilute HC1 was amorphous to X-ray detection up to 650 °C. After calcination at 735 °C they found that SbgO^cr) was formed. Based on these observations it is not expected that Sb2C>4(cr) or Sb6Oi3(cr) are formed at room temperature and no thermodynamic values for these solids are recommended in this report. Solubility products (with relation to Sb(OH)5°), converted in this report from AfG° values given in the literature, are compiled in Tables 5.26 and 5.27. Table 5.26: Experimentally determined equilibrium data compiled for Sb2O4(s). These data were not chosen in the present report for the evaluation of recommended stability values. Method: emf = electromotive force measurements at high temperatures. log K*so Reference Comments - I(M) Medium Method + log K*so: 2Sb(OH)5°+ 2H + 2er <=> Sb2O4(s) + 6H2O 35.63'-2 [1987PAN/SRE] T= 298.15 K, I=n/a n/a emf 1 this value was calculated assuming a log K (Sb(OH)5°/Sb(cr)) of -33.83 (Section 5.7.3) 2 extrapolated from = 800 K; originally calculated with for A,G° of a-Sb2O3(cr) (-613.03 kJ/mol) (1986AZA/PAN]; in this report corrected for AfG° of a-Sb2O3(cr) (-624.32 kJ/rnol) derived in this report 120 JNC TN8400 99-011 Table 5.27: Thermodynamic data for Sb2O4(s) and Sb6Ol3(s) taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. All these values were calculated assuming a log K (Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3) logPm 5m+n Reference Comments + log K\ 0: 2Sb(OH)5°+ 2H + 2e~ <=? Sb2O4(s) + 6H2O 23.06 [1952LAT] T= 298.15 K, I=n/a 35.23 [1954COU] T= 298.15 K, I=n/a 35.23 [1963WIC/BLO] T= 298.15 K, I=n/a 40.92 [1977BAR/KNA] T= 298.15 K, I=n/a 40.92 [1979KUB/ALC] T= 298.15 K, I=n/a 40.88 [1980BEN/TEA] T= 298.15 K, I=n/a 40.91 [1982PAN] T= 298.15 K, I=n/a 40.87 [1982WAG/EVA] T= 298.15 K, I=n/a 40.89 [1985PAS] T= 298.15 K, I=n/a 40.89 [1986ITAyNISl T= 298.15 K, I=n/a + log r 0: 6Sb(OH)5° + 4H + 4e~ <=* Sb6OI3(s) -I- 77//2O 77.71 fl952LATl T= 298.15 K, I=n/a 1 all these values were calculated assuming a log K (Sb(OH)5°/Sb(cr)) of-33.83 (Section 5.7.3) 121 JNC TN8400 99-011 5.11 Other antimony(V) complexes and compounds For other antimony(V) species and compounds the amount of experimentally determined data available in the literature is rather limited. [1975ALY/ABD] determined stability constants of Sb(V) chlorocomplexes by cation exchange. They give no indication of the speciation of Sb(V) present in 0.1 - 2 M HC1. However, in the pH range -0.3 to 1, the presence of Sb(OH)5° (or of is probable (see Section 5.8). The formation constants given by [1975ALY/ABD] for the formation of Sb(V) chlorocomplexes are compiled in Table 5.28. [1969BRY/IOF] state that SbClg- dominates in concentrated hydrochloric acid, SbCl5OH- in 9 M HC1 and SbCl4OH2- in 6 M HC1. [1974BLA/BUR] observed a solubility of 0.0033 M NaSb(OH)6(s) in water. As pH was not measured it is not possible to calculate a solubility product. [1996KAS/MUK] give standard enthalpies and entropies for alkaline-earth metal antimonates produced in the temperature range of 200-600 °C. From the available data it can be expected that Sb(V) forms weak complexes with chloride (and fluoride) and that easily soluble solids with alkali and earth alkali metal ions can be formed. However, there are not sufficient data available to recommend thermodynamic data for these complexes and solids. Data for other inorganic ligands (e.g., sulfate, nitrate, fluoride, ..) have not reported in the literature. Table 5.28: Experimentally determined equilibrium data compiled for the antimony(V) system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 5.12: 'Comments on selected references'. Method: ex = cation exchange. logp Reference Comments I (M) Medium Method log KUJ: "Sb(V) + Cl- <=> Sb(V)Cl- 0.34 [1975ALY/ABD] T= 298.15 K, I=.l 0.1 HC1 ex 0.32 [1975ALY/ABD] T= 298.15 K, I=.5 0.5 HC1 ex 0.20 [1975ALY/ABD] T= 298.15 K, 1=1 1 HC1 ex 0.46 [1975ALY/ABD] T= 298.15 K, 1=2 2 HC1 ex 1.32 [1975ALY/ABD] T= 298.15 K, 1=4 4 HC1 ex 1.82 [1975ALY/ABD1 T= 298.15 K, 1=6 6 HC1 ex 1 Sb = 1 mM; [1975ALY/ABD] give no indication of the speciation of Sb(V) present in 0.1 -2M HC1. However, at pH 1 4 to -0.3 the presence of Sb(OH)5° (or of Sb]2(OH)M -) is probable (see Section 5.8). [1975ALY/ABD] observed in 4 and 6 M HC1 solutions an increase of the measured log K values, which is consistent of the hydrolysis as observed by [1948TOU/MOU] (Section 5.8). 122 JNC TN8400 99-011 5.12 Comments on selected references [1923GRU/SCH]: [1923GRU/SCH] determined the redox potential of the reactions Sb(cr) + 4OH- <=> Sb(OH)4- + 3e~ and Sb(OH)4- + 2OH~ & Sb(OH)6- + 2e~ as -0.675 V and -0.589 V, respectively, in 10 N KOH. The redox potential of Sb(OH)4- + 2OH~ <=> Sb(OH)<5~ + 2e~ was also determined more diluted solutions (Table 5.17). [1924SCH]: [1924SCH] determined in 0.2 - 1.1 M HC1O4 both the solubility of Sb(III) in perchlorate acid and the electrode potential of the reaction Sb(cr) + H2O <=> SbO+ + 2H+ + 3e~. Both measurements resulted in similar log (31 2 values (Table 5.1 and 5.2). For comparison the log of these values were corrected for the reaction Sb(cr) + 3H2O « Sb(OH)3 + 3H+ + 3e~ with a factor of 11.99 (see Section 5.7: Redox reactions). These values agree very well with the log pi 2 values for the reaction + Sb(OH)3° + H+ <=> Sb(OH)2 + H2O measured with other techniques (Tables 5.1 and 5.2). [1948TOU/MOU]: [1948TOU/MOU] determined the solubility of valentinite (a-Sb2O3) in water and diluted NaOH at 35 °C. From their experimental results, constants for the reactions V2 a-Sb2O3 + 1.5H2O <=> Sb(OH)3° and V2 a-Sb2O3 + 2.5H2O <=> + Sb(OH)4- + H and Sb(OH)3° + H2O «• Sb(OH)4- + H+ can be calculated. Their values agree well with data measured at 25 °C (see Tables 5.1 and 5.4) and were included in our calculations. [1948TOU/MOU] determined the solubility of Sb2O5(precip) in water and in diluted HC1 at 35 °C. They reported a solubility of 2.71x10^ M Sb(V) in water. 4 Their results in diluted acids are consistent with the existence of Sbi2(OH)64 ~ (or 3 Sb12(OH)63 -) as proposed by [1963LEF/MAR] and [1976BAE/MES]. In more 3+ acidic solutions (pH < - 0.25) Sb(OH)2 (or a polynuclear compound having the same Sb(V)/H+ ratio) seems to dominate the solution. [1952GAY/GAR]: [1952GAY/GAR] measured in 0 - 0.1 M HC1 and 0 - 0.1 M NaOH Sb + hydrolysis and determined a mean log P12 of 1.13 for the reaction Sb(OH)3° + H + <=> SbO + 2H2O. As Sb is known to form weak complexes with chloride, measurements in HC1 medium were not used for extrapolation. However, the values in acidic medium agree very well with other measurements reported in the literature (Tables 5.1 and 5.2). In alkaline medium, constants can be calculated as a function of I from their experimental results for the reactions V2 a-Sb2O3 + 1.5H2O <=> Sb(OH)3° and V2a-Sb203 + 2.5H2O <=> Sb(OH)4- + H+ (respectively + Sb(OH)3° + H2O <=> Sb(OH)4- + H ). The constants calculated from the measurements agree well with other measurements given in the literature (Table 5.1) and were included in our calculations. [1957PIT/POU]: [1957PIT/POU] calculated the Sb2O3 solubility from data from different sources. 123 JNC TN8400 99-011 [1968MIS/GUP]: [1968MIS/GUP] determined spectrophotometrically a mean log p\2 value + of 1.42 in 0.02-0.1 M HC1O4 for the reaction Sb(OH)3° + H+ <=> SbO + 2H2O. From their measurements it can be concluded that from pH 2 to 1 (0.02-0.1 M HCIO4) no further hydrolysis of Sb(III) takes place (predominance of SbO+, or + Sb(OH)2 ), while at pH < 0 (1 and 3 M HC1O4) extinction coefficients are different indicating the presence of SbOH2+ or Sb3+. [1970BON/WAU]: [1970BON/WAU] and [1970BON] determined polarographically in 5 and 2 M HCIO4 values for the reaction Sb(cr) <=> Sb3+ + 3e~. From the redox potential given by [1970BONAVAU] and [1970BON] (corrected in this report for the potential of the AgCl/Ag electrode; 0.1988 V in saturated media, 3+ [1996STU/MOR]), log p°i)0 values for the reaction Sb(OH)3° + 3H+ <=> Sb + 3H2O can be calculated (using a log K° of 11.99 for the reaction Sb(cr) + 3H2O <=> Sb(OH)3 + 3H+ + 3e~ Section 5.7.1). The values are given in Table 5.2 and agree quite well with the log Pi.o values measured by [1977ANT/NEV]. It is not clear, however, if in the experiments of Bond and co-workers [1970BONAVAU, 1970BON], Sb3+ is the only species present. [1970BON]: see [1970BON/WAU] [1973BEH/ROS]: [1973BEH/ROS] calculated AfG° values at 298 K for a-Sb2O3 based on determination of AH and AS values at 600-700 K. [1973VAS/SHO2]: [1973VAS/SHO2] determined the solubility of valentinite in dilute NaOH at 25 °C. From their experimental results, constants for the reaction + Sb(OH)3° + H2O o Sb(OH)4~ + H can be calculated as a function of I (Table 5.1). [1974SHO/MAB]: [1974SHO/MAB] studied the hydrolysis of Sb(III) in 0.1 and 3 M 8 NaC104 and at a Sb concentration of < 10~ M. Their results show that no Sb-ClO4 complexes are formed, and that the formation of Sb(III) chloride complexes will become important only at pH values < 2. [1975ALY/ABD]: [1975ALY/ABD] determined stability constants of Sb(V) chlorocomplexes by cation exchange. They give no indication of the speciation of Sb(V) present in 0.1 - 2 M HC1. However, at pH 1 to -0.3 the presence of 4 Sb(OH)5° (or of Sb12(OH)64 -) is probable (see Section 5.8). [1975ALY/ABD] observed in 4 and 6 M HC1 (pH < -0.5 ) solutions an increase of the measured log K values, which is indicates further protonation of Sb(V) and is consistent with the hydrolysis as observed by [1948TOU/MOU] (Section 5.8). [1975HEI/SCH]: [1975HEI/SCH] determined potentiometrically in 1.5 M H2SO4 and HF + the redox potential of the Sb(OH)2 /Sb(cr) couple. From this value an approximate + log pli2 of 0.36 for the reaction Sb(OH)3° + H £=> Sb(OH)2+ + H2O (Table 5.2) 124 JNC TN8400 99-011 was calculated, using log K of 11.99 for the reaction Sb(cr) + 3H2O <=> Sb(OH)3 + 3H+ + 3e~ (see Section 5.7: Redox reactions). Antimony forms complexes with the fluoride and sulfate present in the electrolyte. Although [1975HEI/SCH] corrected their measurements for the interactions of Sb(III) with F~ and SO42-, their measurements differ from the other data reported. Considering the high H+ concentration, the presence of Sb3+ or SbOH2+ is quite probable. Assuming the 3+ presence of Sb gives a log p])0 of 0.53 for the reaction Sb(OH)3° + 3H+ <=> Sb3+ + 3H2O which agrees well with other values found in literature (Table 5.2). From the indications given by [1975HEI/SCH] the exact value of ionic strength of the electrolyte present is not clear (1.5 - 4.5 M). [1977ANT/NEV]: [1977ANT/NEV] studied the hydrolysis of Sb(III) under acidic 5 conditions in 1 M NaC104 containing 2 x 10" M Sb(ffl). [1977ANT/NEV] used a spectrophotometric method in which the competition between the hydrolysis reaction and complexation with gallein was measured. [1994AKI/ZOT]: [1994AKI/ZOT] determined Sb(III) hydrolysis and complex formation with sulfide from measurements at 25 - 300 °C and an Sb(III) concentration of 40 mM. They put up a consistent thermodynamic data set. They also included results determined in other studies in their calculations. The complex formation constants for the Sb(III) sulfide system are based on solubility measurements of stibnite made by [1994AKI/ZOT] and by [1988KRU] at different pH values, temperatures and H2S concentrations. 125 JNC TN8400 99-011 6 Lead The oxidation states 0, +11, +IV are found in naturally occurring lead compounds. The most common oxidation state is +11. The hydrolysis behavior of Pb(II) has been the subject of extensive work. There are at least eight species which exist under widely varying conditions. The plumbous ion, Pb2+, is stable in acid solutions and forms complexes with most negatively charged ions. At Pb concentrations > 10~5 M, polynuclear Pb species play a dominant role; the 4+ 4+ most prominent polynuclear species are Pb4(OH)4 and Pb6(OH)8 [1976BAE/MES]. The stable solid phase in water under most conditions is red PbO(s). The compounds Pb3O4(s) and PbC>2(s) exist in very oxidizing environments. It is known that PbC>2(s) is very insoluble. Pb4+ also hydrolyzes extensively but its hydrolysis products are not well known [1952LAT, 1976BAE/MES, 1985GAL, 1995WIB]. Besides complex formation with organic ligands [1976SMI/MAR], thermodynamic data are available for the formation of Pb2+ complexes or compounds with the following inorganic ligands: hydroxide, chloride, fluoride, nitrate, phosphate, sulfate and sulfide. Equilibrium constants for the hydrolysis of Pb2+ and the complex formation with chloride, fluoride, nitrate, phosphate, sulfate and sulfide are given in the Sections 6.1, 6.3 - 6.10. Also the solubility product of PbO(red), PbO(yellow), Pb(OH)2(precip), PbCl2(s), and PbOHCl(s), PbF2(s), PbClF(s), cerrusite, hydrocerrusite, Plumbonacrite, PbOHNO3(s), different lead phosphate solids, anglesite and galena are calculated from experimental data given in literature (Sections 6.2-6.10). The redox reactions of Pb are discussed in Section 6.11 and the hydrolysis of Pb(IV) in Section 6.12. 6.1 Hydrolysis of lead The hydrolysis of Pb(II) is complex. At least eight different Pb(II) complexes are proposed to be present in aqueous solutions. The most prominent polynuclear species in concentrated lead 5 4+ 4+ 2+ solutions (Pb > 10" M) are Pb4(OH)4 and Pb6(OH)8 [1976BAE/MES]. The Pb ion will dominate lead speciation in acidic solutions; in solutions with pH 12 or higher the anion Pb(OH)3" will be the most important species. Experimental data of the hydrolysis of lead used in this report for extrapolation to I = 0 are given in Table 6.1. Further experimental results reported in the literature are shown in Table 6.2 and log [3 values proposed by authors of previous reviews are collected in Table 6.3. Pedersen [1945PED] has published one of the early works investigating the hydrolysis of Pb(II). Unfortunately, [1945PED and 1954FAU] measured lead hydrolysis in nitrate media. As lead forms complexes with nitrate (see Section 6.6), these data cannot be used directly for the evaluation of the hydrolysis constants of lead. In the 1960s Pb(II) hydrolysis has been determined in 0.3 and 3 M perchlorate media by Olin and co-workers [1960OLI1, 1960OLI2, 1960CAR/OLI, 1961OLI, and 1962PAJ/OLI] in concentrated (1 - 1500 mM Pb) lead solutions. 126 JNC TN8400 99-011 These data have been critically reviewed and extrapolated to I = 0 by [1976BAE/MES]. Several other authors determined lead hydrolysis in 2-3 M perchlorate media [1964HUG, 1976LEE2, 1980KAW/ISH, and 1981KOG/OKA], see Table 6.1. The log p values reported by these authors are all quite similar. Comparison with the data reported by [1965HUG] for 2 M NaNO3 (Table 6.2) shows a strong influence of nitrate on the measured log P values for the hydrolysis of Pb(II). Similarly, the data of [1967SCH/ING] (Table 6.2) illustrate that lead chloride complexes are formed. Consequently, also the data measured by [1973BIL/STU, 1976BIL/HUS, and 1980SYL/BRO] in 0.1 M KNO3 were not used in the present evaluation of formation constants. In the case of nitrate media, the complex formation constants between Pb2+ and nitrate are well defined (see Section 6.6), and theoretically the formation constants can be corrected for the effect of the complex formation with nitrate. However, a closer examination of the data showed a distinct shift of the log P values for the polynuclear Pb(II) species depending on the nitrate 2+ 3+ 4+ 2+ concentration which indicates that not only Pb but also Pb2OH , Pb4(OH)4 , Pb3(OH)4 , Pb3(OH)5+, and Pb6(OH)s4+ interact with nitrate. At low nitrate concentrations, the interaction of Pb(II) with nitrate becomes very small. At NO3- = 0.05 M less than 5% of total Pb(II) are present as PbNCV and less than 0.2% are present as Pb(NO3)2°. Other lead nitrate complexes will be probably even less important. Thus, data measured at NO3" < 0.05 M were included in the calculations (Table 6.1). The data used for the calculation of the formation constants of lead(II) hydroxide complexes valid at I = 0 are listed in Table 6.1 and shown in Figure 6.1 to 6.8. Additional data for the lead(II) hydroxide system are compiled in Table 6.2 and 6.3. These data were not chosen for the calculation of log (3° values in this report. Table 6.1: Experimentally determined equilibrium data compiled for the lead hydroxide 2+ 2m n system, according to the equilibrium* mPb + nH2O <=> Pbm(OH)n - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: Comments on selected references. Method: sol = solubility measurements, tit = titration (pH), and pot = potentiometry. log Pm Reference Comments KM, Medium Method 2+ + + logfr ,:Pb + H20<=^Pb0H + H -1.11 1 [1945PED] T= 291 K, 1=0.06 0.06 Ba(NO3)2 tit -7.88 i [1945PED] T= 291 K, 1=0.03 0.03 Ba(NO3)2 tit ! -7.86 [1945PED] T= 291 K, 1=0.015 0.015 Ba(NO3)2 tit -7.90 [1960OLI1] T= 298.15 K, 1=3 3 NaClO4 pot -7.80 [1960OLI1] T= 298.15 K, 1=0.3 0.3 NaClO4 pot -7.93 [1964HUG] T= 298.15 K, 1=2 2 NaClO4 pot 127 JNC TN8400 99-011 Table 6.1: continued 2 -8.0 [1976LEE2] T= 298.15 K, 1=3 3 NaClO4 pot 3 -7.32 [1978LIN] T= 298.15 K, 1=0.01 0.01 NaC104 pol 3 -7.36 [1978LIN] T= 298.15 K, 1=0.01 0.01 NaC104 pol -7.80 [1993CRU/VAN] T= 298.15 K, 1=1 1 NaClO4 tit 2+ + log fax Pb + 2H2OaPb(OH)2° + 2H -17.08 4 [1939GAR/VEL] T= 298.15 K, 1=0-0.1 0 NaOH sol -17.46 [1960CAR7OLI] T= 298.15 K, 1=3 3 NaC104 pot -17.18 [1960CAR/OLI] T— 298.15 K, 1=0.3 0.3 NaC104 pot 2 -17.02 [1978LIN] T= 298.15 K, 1=0.01 0.01 NaClO4 pol -17.01 2 [1978LIN] T= 298.15 K, 1=0.01 0.01 NaC104 pol 2+ log fax Pb + 3H2O 4=>Pb(OH)3- + 3H+ -28.04 4 [1939GAR/VEL] •y 298.15 K, 1=0-0.1 0 NaOH sol -28.88 [1960CAR/OLI] T= 298.15 K, 1=3 3 NaC104 pot -27.99 [1960CAR/OLI] T= 298.15 K, 1=0.3 0.3 NaC104 pot 3 -28.10 [1978LIN] T= 298.15 K, 1=0.01 0.01 NaC104 pol 5 -29.24 [1987FER/GRE] T= 298.15 K, 1=3 3 NaClO4 sol 2+ log fa,i: 2Pb + H20t=>Pb20. + H+ -6.93 i [1945PED] T= 291 K, 1=0.06 0.06 Ba(NO,)2 tit -7.05 i [1945PED] T= 291 K, 1=0.03 0.03 Ba(NO02 tit -7.13 i [1945PED] T= 291 K, 1=0.015 0.015 Ba(NO,)2 tit -6.45 [1960OLI2] T= 298.15 K, 1=3.5 3.5 NaClO4/Pb(ClO4)2 pot -6.45 [1960OLI2] T= 298.15 K, 1=4 4 NaC104/Pb(C104)2 pot -6.30 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(C104)2 pot -6.39 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(C104)2 pot -6.24 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(C104)2 pot -6.49 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot -6.57 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot -6.45 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot 6 -6.30 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Pb(C104), pot 2+ ,44+ + 4tf+ log faA- 4Pb + 4H2O&Pb4( OH) ] -2025 [1945PED] T= 291 K, 1=0.06 0.06 Ba(NO,)2 tit -20.45 1 [1945PED] T= 291 K, 1=0.03 0.03 Ba(NO,)2 tit -20.54 i [1945PED] T= 291 K, 1=0.015 0.015 Ba(NO3)2 tit -19.25 [1960OLI1] T= 298.15 K, 1=3 3 NaClO4 pot -19.90 [1960OLI1] T= 298.15 K, 1=0.3 0.3 NaC104 pot -19.25 [1960OLI2] T= 298.15 K, 1=3.5 3.5 NaC104/Pb(C104), pot 128 JNC TN8400 99-011 Table 6.1: continued: -19.23 [1960OLI2] T= 298.15 K, 1=4 4 NaC104/Pb(C104)2 pot -19.16 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(C104)2 pot -19.12 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(ClO4)2 pot -19.11 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(C104)2 pot -19.12 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Ba(ClO4)2 pot -18.9 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(ClO4)2 pot -18.95 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot -18.98 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot -19.05 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Mg(C104)2 pot 6 -19.19 [1962PAJ/OLI] T= 298.15 K, 1=4.5 4.5 Pb(C104)2 pot -19.35 [1964HUG] T= 298.15 K, 1=2 2 NaC104 pot 2 -19.1 [1976LEE2] T= 298.15 K, 1=3 3 NaClO4 pot -19.42 [1980KAW/ISH] T= 298.15 K, 1=3 3 LiC104 pot 7 -18.90 [1981KOG/OKA] T= 298.15 K, 1=3 3 LiC104 pot -19.58 [1993CRU/VAN] T= 298.15 K, 1=1 1 NaClO4 tit 2+ 2+ log foA: 3Pb + 4H2O^Pb3( OH)4 + 4H+ -22.87 [1960OLI1] T= 298.15 K, 1=3 3 NaClO4 pot -23.35 [1960OLI1] T= 298.15 K, 1=0.3 0.3 NaClO4 pot 2 -22.6 [1976LEE2] T= 298.15 K, 1=3 3 NaC104 pot -22.78 [1980KAW/ISH] T= 298.15 K, 1=3 3 LiC104 pot 7 -23.03 [1981KOG/OKA] T= 298.15 K, 1=3 3 LiClO4 pot -22.69 [1993CRU/VAN] T= 298.15 K, 1=1 1 NaC104 tit 2+ + + log p3,5-'3Pb + 5H2O <=> Pb3(OH)5 + 5H . -31.62 8 [ 1980S YL/BRO] T= 298.15 K, 1=3 3 NaC104 pot -32.27 [ 1980S YL/BRO] T= 298.15 K, 1=0.3 0.3 NaClO4 pot -30.80 [1993CRU/VAN] T= 298.15 K, 1=1 1 NaC104 tit 2+ 4+ log p6,8-'6Pb + 8H2O <^Pb6(OH)8 + 8H+ -42.14 [1960OLI1] T= 298.15 K, 1=3 3 NaC104 pot -42.66 [1960OLI1] T= 298.15 K, 1=0.3 0.3 NaClO4 pot -42.1 2 [1976LEE2] T= 298.15 K, 1=3 3 NaC104 pot -42.33 [1980KAW/ISH] T= 298.15 K, 1=3 3 LiC104 pot 7 -41.68 [1981KOG/OKA] T= 298.15 K, 1=3 3 LiC104 pot -42.43 [1993CRU/VAN] T= 298.15 K, 1=1 1 NaClO4 tit + corrected in this report for the effect of the formation of PbNO3 and Pb(NO3)2° complexes using the constants derived in Section 6.7.1. Uncorrected values are given in Table 6.2. Pb concentration = 0.3-1.2 mM. calculated with a pKw of 13.91 at 1=0.01. extrapolated to 1 = 0 with SIT in this report, see Section 6.13: Comments on selected references. calculated with a pKw of 14.18 at 1=3. recalculation of data of [1960OLI2]. Pb concentration = 5-40 mM recalculation of data of [ 1960OLI1 ]. 129 JNC TN8400 99-011 6.1.1 PbOH+ In diluted solutions, Pb2+ hydrolyzes to PbOH+, and higher-hydrolyzed species (see also Sections 6.1.2, 6.1.3, and 6.1.4). To calculate the formation constant log p\i at I = 0, the measurements of Olin and co-workers and the data measured in perchlorate medium by [1964HUG, 1976LEE2, 1978LIN, and 1993CRU/VAN] (see Table 6.1) were used. The data measured in nitrate media by [1945PED] in 0.01, 0.02 and 0.04 M nitrate media were corrected for the formation of lead nitrate complexes using the constants calculated in Section 6.7.1 (cf. Table 6.1). Extrapolation of these measurements to I = 0 gave the following formation constant log P°u for PbOH+, as illustrated in Figure 6.1. 2+ Pb +H2O PbOH+ + H+ log p°u=-7.51 Pb2++H,0 «=> PbOH++H+ -b -I i - ••• :.^.;j..;I,,.-U^.i.:..: ,.j:,,:,L—•;.= ;.i>|= ,,;...., ..:..-.,•... , , ^.-ji:^;;,;;.:;;; :,;>•. j -5.5 -6 Q -6.5 • CM + -7 • -jjlllllfli t CO. -7.5 ! -8.5 •••>".;--t i::;jjjiij;-*'i!?l'=vi^iiO'Of1'"'1''^' ^7^'C *1' ' " -•'^^ ij^i.fifii'ri":';;": -9 -9.5 - -m - C) 2 4 C L, molal 2+ Figure 6.1: Plot of log (3U + 2D vs. Im for the reaction Pb + H2O <=> PbOH+ + H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.02; log P°ij = - 7.51. Calculated from data compiled in Table 6.1. The experimental values given by [1980KAW/ISH] and [1989DOR/MAR] (see Table 6.2) were not used. [1980KAW/ISH] themselves classified their value as doubtful due to the low PbOH+ concentration in their experiments. [1989DOR/MAR] did not report experimental details, such as concentration and type of electrolyte present, and no Pb hydroxide species other than the 1:1 complex was included in their calculations. 130 JNC TN8400 99-011 6.1.2 Pb(OH)2° Measurements of lead hydrolysis at higher pH values were carried out by [1939GAR/VEL] and [1960CAR/OLI]. [1939GAR/VEL] measured Pb hydrolysis in 0.002 - 1.2 M NaOH. From the 2+ measurements of [1939GAR/VEL] a log p\2 value at I = 0 of -17.08 for the reaction Pb + 2H2O <=> Pb(OH)2° + 2H+ could be calculated in this report (experimental data of [1939GAR/VEL], extrapolation to 1=0 and conversion to log p\2 value see [1939GAR/VEL] in Section 6.13: Comments on selected references). The extrapolation of the data of [1939GAR/VEL, 1960CAR/OLI and 1978LIN] gives (see also Figure 6.2): 2+ Pb +2H2O Pb(OH)2° + 2H+ log (3\2 = -16.95 2+ + Pb +2H2O <=> Pb(OH)2°+2H -15.5 - -16 - ; =! = .-::--:.r.. :'\'.;.V-^.;:.;:.~~':..^::-v.::.:'>..i!.V -.:•'-• '•. , .: /"--.- V " •.-••-•; Q -16.5 - »- -17.5 - •;.;': Ji'';::"^--iJ:"i?; :\! T::':';':: " " !;-;i;:i:;iiJi;-;:i:;;;";i-.iiV:;i;i-:i-=J: D) - '.'•'• -Vvl'' ?.\>"-\}'~z~Js'..l'.? il'^W O -18 •••••. ' :i£^^WMB^k -18.5 • : ; lilS;iillti;ilii?:: s iSIfil" S: -19 iiIiiii:i:iliiiiiiil®^l 'S?-#ffia«i:' -19.5 •:::;;ii;^;:!^::;n::^ :;;{:y;^:^::;i>;^:s^s;i^,:::;::;::^^:::^:) r;';;; p;; \1 v;;. -20 2 - 4 lm, molal 2+ Figure 6.2: Plot of log pi,2 + 2D vs. Im for the reaction Pb + 2H2O <=> Pb(OH)2° + 2H+ at 25 °C. The straight line shows the result of the linear regression: Ae = -0.0, log (3°i2 = - 16.95. Calculated from data compiled in Table 6.1. Again, the data measured by [1973BHVSTU] and [1976BIL/HUS] in 0.1 M KN03 and by [1967SCH/ING] in 3 M NaCl were not chosen for the extrapolation to I = 0. Also the value reported by [1980KAW/ISH] was not used as the authors themselves classified this value as doubtful due to the low Pb(OH)2 concentration in their experiments. 131 JNC TN8400 99-011 6.1.3 Pb(OH)3- [1939GAR/VEL] measured Pb hydrolysis in dilute, alkaline solutions. Extrapolation of these 2+ measurements to 1=0 with the SIT model gave a log p ] 3 value of -28.04 for the reaction Pb + + 3H2O <=> Pb(OH)3- + 3H (experimental data of [1939GAR/VEL], extrapolation to I = 0 and conversion to log (313 value see: [1939GAR/VEL] in Section 6.13: Comments on selected references). Further data were determined by [1960CAR/OLI, 1978LIN and 1987FER/GRE] in sodium perchlorate media (see Table 6.1). The extrapolation of the data of [1939GAR/VEL, 1960CAR/OLI, 1978LIN and 1987FER/GRE] to I = 0 gives: 2+ Pb +3H2O Pb(OH)3- log =- 28.02 The data measured by [1967SCH/ING] in 3 M NaCl and by [1980WAL/SIN] in wastewater (given in Table 6.2) were not chosen for the extrapolation to I = 0. 2+ + Pb +3H2O « Pb(OH)3-+3H - -25.5 "L'i.iiiiiiriSn'lisli!!^: '•,'?:'?'ll'.rj^'.-'; -26 iiliiiiiliylllillll;. Q -26.5 o llilII|B•||S||jKIlllt + -27 -J: ..".1:'L'-^,n:i;J^:::':-'J:'^':"-•"'•:'-- •::.;. • 1 1 1 CO .,••••;. '>,•.•,"*"•• £ Ji-iMVil-Bi-, !-;:.::-.-; '!';;^.-:.;..!!;, . -27.5 • ;;: -i;r' ;«:£;£5ir;;?=:;=L?; ''V.^A^-' ;ris^ ii:ii; CO. :: : : : : 1 iliSilliililS = ::ii!J".:': i ;i!;' :':i-; ^i:-:ii !::. : : 5 ii;: ii i! : : i i i: i ; :; ; ;;i : -28 ( ;| i!i;r '>-™- '- ; -' " '-"; :;' i "'" = = " - p^i^lil[^ii;!biHi i^ii—^"••-il :v; -:^i:JI'^i(ii* "••f-'i, ;'~;i:^-:::?i^::i-: "":Vv.-^ i :!:; -28.5 -;.:::•; i./;^:H^";'v;:?;i::;ii?*^i^^;-^;,:: : liijllllll •:.;•••/. *..'T~,1T:.'.?;•"~V' >'• •«".*V:i!""'iV.'=H"':="4S"•:jf:"7-^ -29 Sli'ilil.pii-':' •:l:-!;';:v'l~", \>--, ^T:[•".?-!" ::ip^S- -29.5 - " :: •; 'y^'"; ., :L".• -:"ll - -sn - ... .^ ;:, :. . ;;:.::• :^. . ^ • _•. 2 4 lm, molal 2+ Figure 6.3: Plot of log (3i>3 + 0D vs. Im for the reaction Pb + 3H2O <^> Pb(OH)3" + 3H+ at 25 °C. The straight line shows the result of the linear regression: Ae = 0.29, log P°i)3 = - 28.02. Calculated from data compiled in Table 6.1. 2 6.1.4 Pb(OH)4 - [1976HEM] and [1987BROAVAN] proposed a log pli4of-39.7 and -37.2 for the formation of the Pb(OH)42" species (compiled in Table 6.3). Direct experimental data, however, are not available and the experiments of [1939GAR/VEL], [1960CAR7OLI] and [1987FER/GRE] show that below a pH of 14, the species Pb(OH)3" dominates lead speciation under alkaline conditions. 132 JNC TN8400 99-011 3+ 6.1.5 Pb2OH In solutions containing more than approximately 10"5 M Pb, polynuclear lead complexes are formed. Around pH 6 to 7, the Pb2OH3+ species can occur in very concentrated lead solutions (Pb > 0.1 M). Evaluation of a formation constant for Pb2OH3+ is somewhat difficult, as the Pb2OH3+ species exists only in concentrated lead solutions (Pb > 0.1 M) in detectable amounts. [1980SYL/BRO] considered their data determined in presence of 0.1 - 2 mM Pb for the Pb2OH3+ complex as unreliable, due to the low Pb2OH3+ concentration in their experiments. Reliable data are measured by Olin and co-workers [1960OLI2, 1962PAJ/OLI] in I = 3 - 4.5 perchlorate media (Table 6.1). [1945PED] determined log $2,\ values at different ionic strength in nitrate media. The data measured in 0.01, 0.02 and 0.04 M nitrate media were corrected for the formation of lead nitrate complexes using the constants calculated in Section 6.7.1 (cf. Table 6.1). Extrapolation to I = 0 is shown in Figure 6.4 and results in: 2+ 3 2Pb + H2O <=> Pb2OH + + H+ log p°2,i =-7.18 Extrapolation of the data measured by [1945PED] in nitrate medium gives a similar result (see Table 6.2), as at low nitrate concentrations, complex formation of Pb(II) with nitrate becomes negligible. j3+ , u+ -5 -5.5 - : 1 : l ;: 1 1 : 8 1 1 1 1;: : *•?•"• :•••':] :i :^ ^r U :F^i ^.H^= ';^r:^ri-^r^-y'-i\ ;^i ^>^' ^i-^i !^ :^^;'.^!!*!:;-^! : !" ;:•:••!':<>;"/• :'":!: -:r ^: '•'•• -6 - ; ^:i. i. i'-: ;--:":!:";i,'!""' 'i;-i :'.\7:', '*.,'' '.'• ''•:' '[ l-_\ '\ '= : '•:::.^ '.• '. ^'i ; ";=!!i\ ',' }'• '• '*•'' \'-: • ••','. V i:-1-:^1-'^^ :!:?; '.!:.- ..":•":" : : : ; -6.5 '" '.'-- i v .'•'." :',"-'-'',r'".-'"-r-'r'\ '"'" •'.:-'''--'"1 '••'•••'• ••;:">":::.:' '•."•'>':• '.;: .. -=! ;,' • :". •; • ::. 'i" :; :. .•.•:.:-."-•.•. ' -2 D -7 ( CM" EQ. -7.5 D) -8 • -8.5 - -9 ; : : 1 :: : ; :: : : '•''-. ••. ••': '•:-'"•'•'. ::?•'.'•• : =.'.:.,' •,-:; - ;C..;i:-:! =^:-'= '":-- - -:- '/l- -'-'-•-£,"•-'.':'•'•:. • !':'-.L- : •• --v^ ••'•Ji:: . .'='- • '." 1 : : : 1 1 -9.5 •_:.• ••••".";::::-;;:. I..: ";.'. :';:•i^;l.: i•.••-•.•|•^;~"".:;:^• ".-:••• ]r-' ""•.••\y.\r, •.-.••..•• y IL!"-. " ".:.;••;•:••=- . -m - 0 2 4 6 lm, molal 2+ 3+ + Figure 6.4: Plot of log P2,i - 2D vs. Im for the reaction 2Pb + H2O <^> Pb2OH + H at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.04; log P°2,i = - 7.18. Calculated from data compiled in Table 6.1. 133 JNC TN8400 99 -Oil 4 6.1.6 Pb4(OH)4 + Many independent measurements of the formation constants for Pb4(OH)44+ exist [1945PED, 1960OLI2, 1962PAJ/OLI, 1964HUG, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA, 1991CRU/VAN] (cf. Table 6.1). The data measured in 0.01,0.02 and 0.04 M nitrate media by [1945PED] were corrected for the formation of lead nitrate complexes (cf. Table 6.1). Extrapolation of these measurements to 1=0 (see also Figure 6.5) gives: 2+ 4+ 4Pb + 4H2O <=> Pb4(OH)4 + 4H+ log (3°4,4 = - 20.63 2+ 4+ + 4Pb +4H2O <=> Pb4(OH)4 +4H - 10 -18.5 i - -19 - Q -19.5 - - , . • . 1 -20 - o a so. -20.5 jjo Q °^ V -:—S^"""°—^ D) O -21 -21.5 -22 7x: 20.63 -22.5 - : 1 2 4 L, molal 2+ 4+ Figure 6.5: Plot of log p4;4 - 4D vs. Im for the reaction 4Pb + 4H2O <=> Pb4(OH)4 + 4H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.07; log P°4>4 = - 20.63. Calculated from data compiled in Table 6.1. 134 JNC TN8400 99-011 2+ 6.1.7 Pb3(OH)4 2+ The formation constant for Pb3(OH)4 was calculated based on the measurements by [1960OLI, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA, 1993CRU/VAN] as given in Table 6.1. Extrapolation to 1=0 gives (Figure 6.6): 2+ 2+ + 3Pb + 4H2O Pb3(OH)4 + 4H log (3°3)4 = - 22.48 2+ 2+ + 3Pb +4H2O <=> Pb3(OH)4 +4H -20.5 •• I; : ] :;i : :/;: J : .:•:::.. ?LZ> ,; ^ l;,;;\A:.^;. Oi'T bX- — ^^ i4o - ••• ^ '• ;••• ^ -21 •• Q -21.5 --i; -22 -;|; aL -22.5 -M ;il;ft6iillli|gl8lil|IS|:itil^li5|K^V:;;;:; \Sy • -23 -W : : ! ! : -23.5 -§, ,fl3ia::i;.;:;=5:- 3:ffljil':.i i±*JI Ji:yH;3 ft£!p:;Si:-.;-S#i3r. i:':.£:• %r&3=l', . '. ' -24 -§| -24.5 --I :''"' ''}~::':":-1 '"B^tif :g:° =:j':: ;gN'iiM^-:;; -•••alii j j'w -•• • <•• ;?•«•;-:> -: ^" ^;: •• • 0 2 4 6 lm, molal 2+ 2+ Figure 6.6: Plot of log p3,4 + 4D vs. Im for the reaction 3Pb + 4H2O & Pb3(OH)4 + 4H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.16; log P°3,4= - 22.48. Calculated from data compiled in Table 6.1. 135 JNC TN8400 99-011 3+ 6.1.8 Pb3(OH)3 and Pb3(OH)5+ [1980SYL/BRO and 1980KAW/ISH] proposed the presence of additional polymeric species: 3+ Pb3(OH)3 and Pb3(OH)5+. [1980SYL/BRO] showed that the consideration of the additional presence of Pb3(OH)5+ had hardly an influence on the other log |3 values, as recalculated from the experiments of Olin and co-workers by [1980SYL/BRO] (see Table 6.1 and 6.2). The extrapolation of the formation constant of Pb3(OH)5+ based on the values calculated by [1980SYL/BRO] for 1=0.3 and 3 M NaC104 and the values determined by [1993CRU/VAN] gives (Figure 6.7): 2+ 3Pb + 5H2O o Pb3(OH)5+ + 5H+ log P°3,5 = - 30.72 2+ + + 3Pb +5H2O <=> Pb3(OH)5 +5H -28 - -28.5 • -29 - ^::'jJ:\ij$BMi:!il!fi:^ Q -29.5 • CD -30- CO -30.5 - CO. CD -31 • -31.5 l : 1 K i : 1 .' ,~i-'^ ;:'.'-. ;J::.:--L.--:=:'-: .ii'v': :•= :.. 'y V ';-j^'.!/S ! Hi'fj w -'j .•• ' Q/^'^n.->;•': :B:I::',;|' ::-;J : f : : 1 1 ! : :• : ,:.;;1i..: -,- ::i::.^|!.::; ;: -,-i. .: :"^;;';". . V; • "^ •.\/^/ 1 iC/./V^'^V O V/ • / £f''-'--'- ••••- '-"- '' - - • -32 - ••.:••• v\-:t'i;.i:!^::;;]^:.^:;i:::?,.iij:;1;>.;.i^:r::-:»-:;'""^"•/'."•:•.;i:'i:' "..•.- :•.••: ••• : -32.5 - ' '.'BMiM{g§§MMS^ -33 - n, moral 2+ Figure 6.7: Plot of log p\5 + 6D vs. Im for the reaction 3Pb + 5H2O <=> Pb3(OH)5+ + 5H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.19, log P°3i5 = - 30.72. Calculated from data compiled in Table 6.1. The only experimentally determined formation constant for Pb3(OH)33+ has been determined by [1980KAW/ISH] in 3 M LiClO4. This value is not extrapolated to 1=0 and no log (3 value for this complex is recommended in this report. In any case, such a complex is not important in solutions containing less than 1 mM Pb. 136 JNC TN8400 99 -Oil 4+ 6.1.9 Pb6O(OH)6 Several independent measurements exist for the formation constant of PbeO(OH)64+, log p6,8 ([1960OLI, 1976LEE2, 1980KAW/ISH, 1981KOG/OKA, 1993CRU/VAN]; Table 6.1). Extrapolation to 1=0 gives: 2+ 4+ 6Pb + 8H2O Pb6(OH)8 + 8H+ log {3°6i8 = - 42.68 [1981ISH/OHT] proposed, based on calorimetric measurements, that the complex in fact is Pb6O(OH)64+, which corresponds to X-ray diffraction data indicating a central oxygen atom surrounded by four lead atom [1968JOH/OLI]. 2+ 4+ + 6Pb +8H2O «• Pb6(OH)8 +8H -40 -40.5 - -41 Q -41.5 •• O + -42 + CO. -42.5 ; O -43 - -43.5 -44 y = 0.08x - 42.68 -44.5 -f -45 0 2 4 6 lm, molal 2+ 4+ Figure 6.8: Plot of log |36,8 + 0D vs. Im for the reaction 6Pb + 8H2O <=> Pb6(OH)8 + 8H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.08, log p°6,8 = - 42.68. Calculated from data compiled in Table 6.1. 137 JNC TN8400 99-011 6.1.10 Additional equilibrium data compiled for the lead hydroxide system Table 6.2: Additional experimentally determined equilibrium data compiled for the lead hydroxide system, 2+ 2 n + according to the equilibrium: mPb + nH2O <=> Pbm(OH)n " + nH . These data were not chosen in the present report for the evaluation of recommended stability values. Other reasons for not selecting these references are given at the end of this table or in Section 6.13: 'Comments on selected references'. As pointed out in Section 6.1, studies carried out in nitrate, chloride and sulfate media were not chosen as these ions form complexes with Pb(II). Method: pot = potentiometry, sol = solubility measurements, tit = pH titration. log Reference Comments Medium Method 2+ log pu: Pb + H2O <=> PbOH* + H* -6.18 ' [1939GAR/VEL] T= 298.15 K, 1=0-0.1 NaOH sol 2 -8.74 [1945PED] T=291 K, 1=1.2 1.2 Ba(NO3)2 tit 2 -8.35 [1945PED] T=291 K, 1=0.6 0.6 Ba(NO3)2 tit 2 -8.18 [1945PED] T= 291 K, 1=0.3 0.3 Ba(NO3)2 tit 2 -8.07 [1945PED] T= 291 K, 1=0.15 0.15 Ba(NO3)2 tit 2 3 -7.85 ' [1945PED] T=291 K, 1=0.06 0.06 Ba(NO3)2 tit 2 3 -7.94 ' [1945PED] T=291 K, 1=0.03 0.03 Ba(NO3)2 tit 2 3 -7.89 - [1945PED] T= 291 K, 1=0.015 0.015 Ba(NO3)2 tit 5 -7.73 [1945PED] T=291 K, 1=0.0 0 Ba(NO3)2 tit 5 -8.66 [1954FAU] T= 293 K, 1=0.6 0.6 Ba(NO3)2 tit 5 -8.37 [1954FAU] T= 293 K, 1=0.06 0.06 Ba(NO3)2 tit 6 -8.84 [1965HUG1] T= 298.15 K, 1=2 2 NaNO3 pot -6.96 [1973BDL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot -6.96 [1976BEL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 7 -8.31 [1980KAW/ISH] T= 298.15 K, 1=3 3 LiC104 pot -7.86 [1980SYL/BRO] T= 298.15 K, 1=0.1 0.1 KNO3 pot -8.17 [1980WAL/SIN] T= 298 K, I=dil H2SO4 sol -6.84 [1989DOR/MAR] T=298.15K,I=dil Pb(NO3)2, NaHCOj tit -8.72 [1993CRU/VAN] T= 298.15 K, 1=1 1 KNO3 tit -7.94 10 [1993CRU/VAN1 T= 298.15 K, 1=1 1 KNO, tit + log 2H2O <=> Pb(OH)2° + 2H -20.58 " [1967SCH/ING] T= 298.15 K, 1=3 3 NaCl pot -16.42 [1973BDL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot -16.42 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 7 -16.37 [1980KAW/ISH] T= 298.15 K, 1=3 3 LiC104 pot -17.85 [1980WAL/SIN1 T= 298 K, I=dil H,SO4 sol 2+ log [)u: Pb + 3H2O <=> Pb(OH)/ + 3H+ -32.58 n [1967SCH/ING] T= 298.15 K, 1=3 3 NaCl pot -28.3 [198OWAL/SIN1 T= 298 K, I=dil H,SOd sol 2+ 3 log p2J: 2Pb + H20 <=> Pb2OH * - 2 -7A2 [1945PED] T=291 K, 1=1.2 1.2 Ba(NO3)2 tit 2 -7.05 [1945PED] T=291 K, 1=0.6 0.6 Ba(NO3)2 tit 2 -7.05 [1945PED] T=291 K, 1=0.3 0.3 Ba(NO3)2 tit 2 -7.05 [1945PED] T=291 K, 1=0.15 0.15 Ba(NO3)2 tit 138 JNC TN8400 99-011 Table 6.2: continued -7.10 [1945PED] T=291 K, 1=0.06 0.06 Ba(NO3)2 tit 2 3 -7.15 - [1945PED] T=291 K, 1=0.03 0.03 Ba(NO3)2 tit 2 -7.19 - [1945PED] T=291 K, 1=0.015 0.015 Ba(NO3)2 tit 4 -7.28 [1945PED] T= 291 K, 1=0.0 0 Ba(NO3)2 tit 8 -6.34 [1960OLI2] T= 298.15 K, 1=4.5 4.5 Pb(C104)2 pot 6 -7.11 [1965HUG1] T= 298.15 K, 1=2 2 NaNO3 pot -6.26 [ 1980S YL/BRO] T= 298.15 K, 1=0.1 0.1 KNO3 pot -6.79 [1993CRU/VAN1 T= 298.15 K, 1=1 1 KNO, tit 4+ + log fty 4P&+ + 4H2O <=> Pb4(OH)4 + 4H 2 -20.97 [1945PED] T=291 K, 1=1.2 1.2 Ba(NO3)2 tit 2 -20.80 [1945PED] T=291 K, 1=0.6 0.6 Ba(NO3)2 tit 2 -20.65 [1945PED] T=291 K, 1=0.3 0.3 Ba(NO3)2 tit 2 -20.64 [1945PED] T=291 K, 1=0.15 0.15 Ba(NO3)2 tit 2 3 -20.59 - [1945PED] T=291 K, 1=0.06 0.06 Ba(NO3)2 tit 2 3 -20.66 - [1945PED] T=291 K, 1=0.03 0.03 Ba(NO3)2 tit 2 3 -20.66 . [1945PED] T=291 K, 1=0.015 0.015 Ba(NO3)2 tit 4 -20.96 [1945PED] T=291 K, 1=0.0 Ba(NO3)2 tit s -18.75 [1954FAU] T=293 K, 1=0.6 0.6 Ba(NO3)2 tit 5 -18.05 [1954FAU] T=293 K, 1=0.06 0.06 Ba(NO3)2 tit 8 -19.28 [1960OLI2] T=298 .15 K, 1=4.5 4.5 Pb(ClO4)2 pot 6 -21.72 [1965HUG1] T=298 .15 K, 1=2 2 NaNO3 pot -20.4 • [1980S YL/BRO] T=298 .15 K, 1=0.1 0.1 KNO3 pot -21.01 [1993CRU/VAN] T='298.15 K, 1=1 1 KNO3 tit -19.01 10 [1993CRU/VAN1 T=298 .15 K, 1=1 1 KNO, tit 2+ 3+ + log fty 3Pb + 3H2O <=> Pb3(OH)3 + 3H -8.31 7 [1980KAW/ISH] T= 298.15 K, 1=3 LiCIO, pot 2+ 2+ + log 3Pb 4H2O Pb3(OH)4 + 4H -23.91 [1980SYL/BRO] T= 298.15 K, 1=0.1 0.1 KNO3 pot -24.33 [1993CRU/VAN] T= 298.15 K, 1=1 1 KNO3 tit 10 -22.83 ri993CRU/VAN] T= 298.15 K, 1=1 1 KNO, tit 2+ + log P3,5: 3Pb + 5H2O <=> Pb3(OH)5 + 5FT -31.75 f!980SYL/BRO1 T= 298.15 K, 1=0.1 0.1 KNO, pot 2+ 4+ log p6S: 6Pb + 8H2O <=> Pb6(OH)8 + 8H+ -43.38 [1980SYL/BRO] T= 298.15 K, 1=0.1 0.1 KNO3 pot -44.91 [1993CRU/VAN] T= 298.15 K, 1=1 1 KNO3 tit -41.55 10 ri993CRUA^ANl T= 298.15 K, 1=1 1 KNO, tit polynuclear species not considered, I not constant, calculated with a log K*so of -12.68 for PbO (Section 6.2.1) 3 formation of PbNO3' complexes, only small concentration of Pb2OH * present in his experiments. Pb concentration = 5-400 mM values corrected for the effect of the formation of Pb-nitrate complexes are given in Table 6.1. after extrapolation to 1=0 with SIT (this report) 4+ [1954FAU] assumed Pb4(OH)4 to be only species present at Pb > 10 mM. Pb concentration = 0.25-200 mM. Pb concentration =10-200 mM, formation of lead nitrate complexes not considered + [1980KAW/ISH] themselves classified their value as doubtful due to the low PbOH and Pb(OH)2 concentration in their experiments. corrected value in [1962PAJ/OLI] 9 see Section 6.1.8 10 values corrected for the effect of the formation of Pb-nitrate complexes by [1993CRU/VAN]. Pb = 1-50 mM. 11 Pb concentration =10-500 mM, formation of lead chloride complexes not considered; pKw used = 14.18 139 JNC TN8400 99-011 Table 6.3: Thermodynamic data for the hydrolysis of lead taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log Pm,n Reference Comments KM) Medium 2+ log (3U: Pb + H2O <=> PbOH* + H+ -7.90 ' [1961OLI] T=298.15K, 1=3 3 NaClO4 -7.80 • [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 2 -7.81 [1974VAD] T= 298.15 K, 1=3 3 NaC104 2 -7.84 [1974 V AD] T= 298.15 K, 1=3 3 NaC104 -7.71 [1976BAE/MES] T= 298.15 K, 1=0 0 -6.57 [1976HEM] T= 298.15 K, I=n/a. -7.70 [1976SMI/MAR] T= 298.15 K, 1=0 0 -7.8 [1976SMI/MAR] T= 298.15 K, 1=0.3 0.3 -7.9 [1976SMI/MAR] T= 298.15 K, 1=3 3 -7.22 [1980SCH] T = 298.15 K, 1=0 0 2 -7.89 [1980SYL/BRO] T= 298.15 K, 1=3 3 NaClO4 2 -7.79 [1980SYL/BRO] T= 298.15 K, 1=0.3 0.3 NaClO4 -7.71 [1981BAEMES] T= 298.15 K, 1=0 0 -6.17 [1981STU/MOR] T= 298.15 K, 1=0 0 -7.71 [1981TUR/WHI] T= 298.15 K, 1=0 - 0 -7.70 [1982SMI/MAR] T= 298.15 K, 1=0 0 -7.8 [1982SMI/MAR] T= 298.15 K, 1=0.3 0.3 -7.9 [1982SMI/MAR] T= 298.15 K, 1=3 3 -6.18 [1982WAG/EVA] T= 298.15 K, 1=0 0 -7.71 [1983LAN] T= 298.15 K, 1=0 0 -7.70 [1984TAY/LOP] T= 295 K, I=dil -6.17 [1985BAB/MAT] T= 298.15 K, 1=0 0 -7.71 [1987BROAVAN] T= 298.15 K, 1=0 0 -8.00 [1988BYR/KUM] T=291 K, 1=0.7 0.7 -7.64 [1988PHI/HAL] T= 298.15 K, 1=0 0 -7.60 [1989SMI/MAR] T= 298.15 K, 1=0 0 -7.8 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 -7.8 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -8.0 [1989SMI/MAR] T= 298.15 K, 1=2 2 -7.9 [1989SMI/MAR1 T= 298.15 K, 1=3 3 2+ log $l2: Pb + 2H2O <=> Pb(OH)2° +2H* -17.46 ' [1961OLI] T= 298.15 K, 1=3 3 NaClO4 -17.18 ' [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 -17.1 [1963FEI/SCH] T= 298.15 K, 1=0 0 -17.12 [1976BAE/MES] T= 298.15 K, 1=0 0 -15.82 [1976HEM] T= 298.15 K,I=n/a. -17.1 [1976SMI/MAR] T= 298.15 K, 1=0 0 -17.2 [1976SMI/MAR] T= 298.15 K, 1=0.3 0.3 -17.5 [1976SMI/MAR] T= 298.15 K, 1=3 3 -16.91 [1980SCH] T = 298.15 K, 1=0 0 -17.12 [1981BAE/MES] T= 298.15 K, 1=0 0 -17.12 [1981TUR/WHI] T= 298.15 K, 1=0 0 -17.1 [1982SMI/MAR] T= 298.15 K, 1=0 0 140 JNC TN8400 99-011 Table 6.3: continued -17.2 [1982SMI/MAR] T= 298.15 K, 1=0.3 0.3 -17.5 [1982SMI/MAR] T= 298.15 K, 1=3 3 -17.12 [1983LAN] T= 298.15 K, 1=0 0 -17.75 [1984TAY/LOP] T= 295 K, I=dil -17.16 [1985BAB/MAT] T= 298.15 K, 1=0 0 -16.55 [1987BROAVAN] T= 298.15 K, 1=0 0 -17.41 [1988BYR/KUM] T= 291 K, 1=0.7 0.7 sw -17.06 [1988PHI/HAL] T= 298.15 K, 1=0 0 -17.1 [1989SMI/MAR] T= 298.15 K, 1=0 0 -17.2 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -17.5 [1989SMI/MAR] T= 298.15 K, 1=3 3 2+ log Pu: Pb + 3H2O <=> Pb(OH); -27.98 [1952LAT] T= 298.15 K, 1=0 0.015 -28.89 ' [1961OLI] T= 298.15 K, 1=3 3 NaC104 -27.99 ' [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 -28.1 [1963FEI/SCH] T= 298.15 K, 1=0 0 -28.06 [1976BAE/MES] T= 298.15 K, 1=0 0 -28.08 [1976HEM] T= 298.15 K, I=n/a. -28.1 [1976SMI/MAR] T= 298.15 K, 1=0 0 -28.0 [1976SMI/MAR] T= 298.15 K, 1=0.3 0.3 -28.8 [1976SMI/MAR] T= 298.15 K, 1=3 3 -54.20 [1977PAU] T= 298.15 K, 1=0 -28.02 [1979PAT/OBR] T = 298.15 K, 1=0 0 -28.07 [1980SCH] T = 298.15 K, 1=0 0 -28.07 [1980BEN/TEA] T= 298.15 K, 1=0 0 -28.06 [1981BAE/MES] T= 298.15 K, 1=0 0 -28.05 [1981STU/MOR] T= 298.15 K, 1=0 0 -28.06 [1981TUR/WHI] T= 298.15 K, 1=0 0 -28.1 [1982SMI/MAR] T= 298.15 K, 1=0 0 -28.0 [1982SMI/MAR] T= 298.15 K, 1=0.3" 0.3 -28.8 [1982SMI/MAR] T= 298.15 K, 1=3 3 -28.08 [1982WAG/EVA] T= 298.15 K, 1=0 0 -28.07 [1983LAN] T= 298.15 K, 1=0 0 -28.09 [1984TAY/LOP] T= 295 K, I=dil -28.07 [1985BAB/MAT] T= 298.15 K, 1=0 0 -27.98 [1985GAL] T= 298.15 K, 1=0 0 -26.39 [1987BROAVAN] T= 298.15 K, 1=0 0 -28.01 [1988PHI/HAL] T= 298.15 K, 1=0 0 -28.1 [1989SMI/MAR] T= 298.15 K, 1=0 0 -28.0 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -28.8 [1989SMI/MAR1 T= 298.15 K, 1=3 3 2+ log P,4: Pb + 4H2O <=> Pb(OH)/- + 4H+ -39.703 [1976HEM] T= 298.15 K, I=n/a -37.19 [1987BRO/WAN] T= 298.15 K, 1=0 141 JNC TN8400 99-011 Table 6.3: continued 2+ 3+ log P21: 2Pb H2O Pb2OH + H* -6.40 ' [1961OLI] T= 298.15 K, 1=4.5 4.5 Pb(C104)2 -6.36 [1976BAE/MES] T= 298.15 K, 1=0 0 -6.40 [1976SMI/MAR] T= 298.15 K, 1=0 0 -6.30 [1976SMI/MAR] T= 298.15 K, 1=3 3 -6.36 [1980SCH] T = 298.15 K, 1=0 0 -6.40 [1982SMI/MAR] T= 298.15 K, 1=0 0 -6.30 [1982SMI/MAR] T= 298.15 K, 1=3 3 -6.36 [1983LAN] T= 298.15 K, 1=0 0 -6.40 [1984TAY/LOP] T= 295 K, I=dil -6.35 [1985BAB/MAT] T= 298.15 K, 1=0 0 -6.34 [1987BRO/WAN] T= 298.15 K, 1=0 0 -6.40 [1989SMI/MAR] T= 298.15 K, 1=0 0 -6.40 [1989SMI/MAR] T= 298.15 K, 1=3 3 2+ log P4,4: 4Pb + 4H2O <=> Pb4(OH)/ + + 4H+ -19.25 ' [1961OLI] T= 298.15 K, 1=3 3 NaC104 -19.90 ' [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 2 -19.27 [1974VAD] T= 298.15 K, 1=3 3 NaC104 2 -19.26 [1974V AD] T= 298.15 K, 1=3 3 NaCIO4 -20.88 [1976BAE/MES] T= 298.15 K, 1=0 0 -20.9 [1976SMI/MAR] T= 298.15 K, 1=0 0 -19.9 [1976SMI/MAR] T= 298.15 K, 1=0.3 0.3 -19.2 [1976SMI/MAR] T= 298.15 K, 1=3 3 -20.87 [1980SCH] T = 298.15 K, 1=0 0 2 -19.26 [1980SYL/BRO] T= 298.15 K, 1=3 3 NaC104 2 -19.94 [1980SYL/BRO] T= 298.15 K, 1=0.3 0.3 NaC104 -19.27 [1980BEN/TEA] T= 298.15 K, 1=0 0 -20.9 [1982SMI/MAR] T= 298.15 K, 1=0 0 -19.9 [1982SMI/MAR] T= 298.15 K, 1=0.3 0.3 -19.1 [1982SMI/MAR] T= 298.15 K, 1=3 3 -19.27 [1982WAG/EVA] T= 298.15 K, 1=0 0 -20.88 [1983LAN] T= 298.15 K, 1=0 0 -20.89 [1984TAY/LOP] T= 295 K, I=dil -21.02 [1987BROAVAN] T= 298.15 K, 1=0 0 -18.98 [1988PH17HAL] T= 298.15 K, 1=0 0 -20 [1989SMI/MAR] T= 298.15 K, 1=0 0 -19.6 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 -19.9 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -19.8 [1989SMI/MAR] T= 298.15 K, 1=0.5 0.5 -19.3 [1989SMI/MAR] T= 298.15 K, 1=2 2 -19.2 [1989SMI/MAR1 T= 298.15 K, 1=3 3 2 7+ log ft,: 3Pb + 4H2O Pb3(OH)4 + 4H+ -22.87 > [1961OLI] T= 298.15 K, 1=3 3 NaC104 -23.35 ' [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 2 -22.90 [1974 V AD] T= 298.15 K,I=3 3 NaC104 -22.92 2 [1974V AD] T= 298.15 K, 1=3 3 NaCIO, -23.88 [1976BAE/MES] T= 298.15 K, 1=0 0 -23.9 [1976SMI/MAR] T= 298.15 K, 1=0 0 -23.3 [1976SMI/MAR] T= 298.15 K, 1=0.3 0.3 142 JNC TN8400 99-011 Table 6.3: continued -22.9 [1976SMI/MAR] T= 298.15 K, 1=3 3 -23.86 [1980SCH] T = 298.15 K, 1=0 0 -23.35 [198OBEN/TEA] T= 298.15 K, 1=0 0 2 -23.01 [ 1980S YL/BRO] T= 298.15 K, 1=3 3 NaC104 2 -23.18 [ 1980S YL/BRO] T= 298.15 K, 1=0.3 0.3 NaC104 -23.9 [1982SMI/MAR] T= 298.15 K, 1=0 0 -23.3 [1982SMI/MAR] T= 298.15 K, 1=0.3 0.3 -22.7 [1982SMI/MAR] T= 298.15 K, 1=3 3 -23.34 [1982WAG/EVA] T= 298.15 K, 1=0 0 -23.88 [1983LAN] T= 298.15 K, 1=0 0 -23.89 [1984TAY/LOP] T= 295 K, I=dil -23.77 [1987BRO/WAN] T= 298.15 K, 1=0 0 -23.65 [1988PHI/HAL] T= 298.15 K, 1=0 0 -23.9 [1989SMI/MAR] T= 298.15 K, 1=0 0 -23.3 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -22.7 [1989SMI/MAR] T= 298.15 K, 1=3 3 2+ + + log p3J: 3Pb + 5H2O <=> Pb3(OH)5 + 5H -30.29 [1987BRO/WAN] T= 298.15 K, 1=0 0 -7.86 |"1988PHI/HAL1 T= 298.15 K, 1=0 0 2+ 4+ log PM: 6Pb + 8H2O <=> Pb6(OH)8 + 8H+ -42.14 [1961OLI] T= 298.15 K, 1=3 3 NaC104 -42.66 [1961OLI] T= 298.15 K, 1=0.3 0.3 NaC104 2 -42.13 [1974 V AD] T= 298.15 K, 1=3 3 NaC104 2 -42.12 [1974VAD] T= 298.15 K, 1=3 3 NaC104 -43.61 [1976BAE/MES] T= 298.15 K, 1=0 0 -43.6 [1976SMI/MAR] T= 298.15 K, 1=0 0 -42.7 [1976SMI/MAR] T= 298.15 K, 1=0.3- 0.3 -42.1 [1976SMI/MAR] T= 298.15 K, 1=3 3 -43.61 [1980SCH] T = 298.15 K, 1=0 0 2 -42.13 [1980SYL/BRO] T= 298.15 K, 1=3 3 NaC104 2 -42.76 [1980SYL/BRO] T= 298.15 K, 1=0.3 0.3 NaC104 -42.66 [1980BEN/TEA] T= 298.15 K, 1=0 0 -43.6 [1982SMI/MAR] T= 298.15 K, 1=0 0 -42.7 [1982SMI/MAR) T= 298.15 K, 1=0.3 0.3 -41.9 [1982SMI/MAR] T= 298.15 K, 1=3 3 -42.66 [1982WAG/EVA] T= 298.15 K, 1=0 0 -43.61 [1983LAN] T= 298.15 K, 1=0 0 -43.59 [1984TAY/LOP] T= 295 K, I=dil 0 -43.6 [1987BROAVAN] T= 298.15 K, 1=0 0 -43.35 [1988PHI/HAL] T= 298.15 K, 1=0 0 -43.6 [1989SMI/MAR] T= 298.15 K, 1=0 0 -42.7 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 -42.1 [1989SMI/MAR] T= 298.15 K, 1=3 3 1 same values as given in [1960CAR/OL1], [1960OLI1], or [1960OL12]. 2 recalculation of the data of [1960OLI1]. 3 cited by J1976HEM] from a reference in Russian. 143 JNC TN8400 99-011 6.2 Solid lead-oxide/hydroxide phases PbO(s) occurs in two crystalline forms: litharge (red), the stable form under ambient conditions, and massicot (yellow). The existence of a hydrated form, freshly precipitated form is reported by [1928RAN/SPE]. 6.2.1 PbO(litharge) and PbO(massicot) The solubility of litharge and massicot has been determined by [1928RAN/SPE] and [1939GAR/VEL], who measured lead solubility in alkaline medium (see Table 6.5). Both showed that in alkaline solutions, the red PbO is more stable than the yellow form. The solubility constants for litharge and massicot can more accurately be determined from the potentiometric data given by [1922APP/REI] and [1923SMI/WOO] (Table 6.4). From their data: 2+ Pb + H2O <=> PbO(s, red) + 2H+ log K*°so= - 12.68 2+ Pb + H2O <=> PbO(s, yellow) + 2H+ log K*°so = - 12.96 These values also are in good agreement with the solubility measurements listed in Table 6.5 for massicot and litharge and with the log K*°so of -12.65 and -12.71 (Table 6.6) proposed for litharge (red PbO(cr)) by [1979PAT/OBR] and [1984TAY/LOP]. While litharge is thermodynamically favored at room temperature, at temperature above 488 °C the yellow massicot is more stable. However, the precipitation of massicot can also been observed at room temperature [1984TAY/LOP, 1995WD3] as the conversion rate between the two phases is small. Table 6.4: Experimentally determined equilibrium data compiled for the for the precipitation of lead hydroxide/oxide, according to the equilibrium: mPb2+ + nH^O <=> 2 n Pbm(OH)n " + nH+. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: sol = solubility measurements, and pot = potentiometry. log K* Reference Comments I (M) Medium Method 2+ log K*so: Pb + H2O a PbO (red, litharge) + 2H+ -12.67 > [1922APP/REI] T= 298.15 K, I=n/a 0 NaOH pot -12.68 2 [1923SMI/WOO] T= 298.15 K, I=n/a 0 no pot 144 JNC TN8400 99 - Oil Table 6.4: continued 2+ log K*so: Pb + H2O t=> PbO(yellow, massicot) + 2H+ 1 -12.96 [1922APP/REI] T= 298.15 K, I=n/a 0 NaOH pot 2+ log K*so: Pb + 2H2O <^Pb(OH)2 (freshly precipitated) + 2H+ -13.05 [1980SCH] T = 298.15 K, 1=0 0 sol 1 calculated by [1922APP/REI] from potentiometric measurements 2 calculated in this report with a log K of -4.25 for the reaction Pb2++ 2e" <=> Pb(cr) (Section 6.11.1) 6.2.2 Precipitated lead hydroxide Based on the calculations for litharge and massicot and the observations of [1928RAN/SPE] (Table 6.5) and [1980SCH] (Table 6.4) a log K*°So of -13.05 is proposed for freshly precipitated lead hydroxide. The log K*°So of -13.15 determined by [1980WAL/SIN] for precipitated lead hydroxide or oxide is slightly larger due to the presence of anionic species in the wastewater examined. 6.2.3 Pb(OH)2(s) In many compilations, a log K*°so of approximately -8.1 is indicated for the reaction Pb2+ + 2H2O <=> Pb(OH)2(s) + 2H+. This value refers to crystalline Pb(OH)2(cr) which is formed at 300 °C [1995MAR/MAC]. Literature data [1980SCH, 1995MAR/MAC] indicate that this solid does not precipitate directly from solutions at room temperature. The absence of Pb(OH)2(cr) in precipitates is consistent with [1956ROB/THE] who pointed out the difficulty of precipitating Pb(OH)2(cr) directly. Their method of preparing Pb(OH)2(cr) involved heating overnight to 280-300 °C. 6.2.4 Additional data for lead hydroxide/oxide compounds Table 6.5: Additional experimentally determined equilibrium data compiled for the formation of lead hydroxide/oxide compounds. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text in Section 6.2 and in Section 6.13: 'Comments on selected references'. Method: emf = electromotive force measurements at high temperature, sol = solubility measurements. log K*s Reference Comments Medium Method 2+ + log K'so- Pb + H2O <=> PbO (red) + 2H -12.68 ' [1922APP/REI] T= 298.15 K, 1=1 NaOH sol -12.63 2 [1928RAN/SPE] T= 298.15 K, I=n/a NaOH sol -12.55 2 [1939GAR/VEL1 T= 298.15 K, 1=0-0. NaOH sol 145 JNC TN8400 99-011 Table 6.5: continued 2+ log K*so: Pb + H2O <=> PbO (yellow) + 2H+ -12.90 ' [1922APP/REI] T= 298.15 K, 1=1 1 NaOH sol -12.75 2 [1928RAN/SPE] T= 298.15 K,I=n/a 0 NaOH sol -12.95 2 [1939GAR/VEL] T= 298.15 K, 1=0-0.1 0 NaOH sol 2+ + log K*so: Pb + 2H2O <=> Pb(OH)2 (fresh precipitated) + 2H -12.91 2 [1928RAN/SPE] T= 298.15 K, I=n/a NaOH sol -13.15 [ 1980WAL/SIN1 T= 298 K, I=dil 0 H,SO,. sol 2+ + log K*so: Pb + 2H2O <=> Pb(OH)2 (crystalline, formed atT = 300 °C) + 2H -7.51 [1972NRI] T= 298.15 K, 1=0.1 0J KOH sol 2+ + log K*so: Pb + H2O <=> PbO (crystallinity not defined) + 2H -13.18 3'4 [1968CHA/FLE] T= 298.15 K, I=n/a n/a emf -12.55 3 [1984BAN1 T= 298.15 K, I=n/a n/a emf 1 calculated with a log p13 value of -28.32 (1=1; Section 6.1.3) and a log Kw of -13.79 (1=1) 2 the solubility constants for litharge and massicot can more accurately be determined from the potentiometric data given in Table 6.4 3 extrapolated from measurements at 600 - 1000 K 4 criticized by [1984BAN] Table 6.6: Thermodynamic data for the fonnation of lead hydroxide/oxide compounds taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log K*so Reference Comments I (M) Medium 2+ + log K*so: Pb + H20 <=> PbO(red) + 2H -12.64 [1952LAT] T= 298.15 K, I=n/a -12.62 [1954COU] T= 298.15 K, I=n/a -12.64 [1960NAS/MER] T= 298.15 K, I=n/a -12.70 [1963FEI/SCH] T= 298.15 K, 1=0 -12.74 [1963WIC/BLO] T= 298.15 K, I=n/a -12.69 [1971NAU/RYZ] T= 298.15 K, I=n/a -12.72 [1976BAE/MES] T= 298.15 K, I=n/a -12.70 [1976SMI/MAR] T= 298.15 K, I=n/a 27.93 [1977PAU] T= 298.15 K, I=n/a -12.67 [1978ROB/HEM2] T= 298.15 K, I=n/a -12.63 [1979KUB/ALC] T= 298.15 K, I=n/a -12.65 ' [1979PAT/OBR] T = 298.15 K, 1=0 -12.68 [1980BEN/TEA] T= 298.15 K, I=n/a -12.69 [1982PAN] T= 298.15 K, I=n/a -12.64 [1982PAU] T= 298.15 K, I=n/a -12.73 [1982WAG/EVA] T= 298.15 K, I=n/a -12.68 [1983LAN1 T= 298.15 K, I=n/a 146 JNC TN8400 99 -Oil Table 6.6: continued -12.71 [1984TAY/LOP] T= 295 K, I=dil 0 HCO3" -12.72 [1984VEE/TAR] T= 298.15 K, I=n/a -12.66 [1985BAB/MAT] T= 298.15 K, I=n/a -12.65 [1985GAL] T= 298.15 K, I=n/a -12.61 fl988PHI/HAL] T= 298.15 K, I=n/a 2+ + log K*so: Pb + H2O <=> PbO(yellow) + 2H -12.79 [1952LAT] T= 298.15 K,I=n/a -12.77 [1954COU] T= 298.15 K,I=n/a -12.9 [1963FEI/SCH] T=298.15K, 1=0 -12.86 [1963WIC/BLO] T=298.15K, I=n/a -12.88 [1971NAU/RYZ] T= 298.15 K, I=n/a -12.90 [1976SMI/MAR] T= 298.15 K,I=n/a -12.79 [1977PAU] T= 298.15 K,I=n/a -12.78 [1978ROB/HEM2] T=298.15K, I=n/a -12.79 [1980BEN/TEA] T= 298.15 K, I=n/a -12.87 [1982PAN] T= 298.15 K,I=n/a -12.79 [1982PAU] T= 298.15 K, I=n/a -12.91 [1982WAG/EVA] T=298.15K, I=n/a -12.79 [1983LAN] T= 298.15 K,I=n/a -12.91 [1983SAN/BAR] T= 298.15 K,I=n/a -12.80 [1985BAB/MAT] T= 298.15 K, I=n/a -12.79 [1985GAL] T= 298.15 K,I=n/a -12.72 [1988PHI/HAL1 T= 298.15 K,I=n/a 2 log K'so: Pb + + 2H2O<=>Pb(OH)>(fresh: precipitated) + 1 -13.63 [1952LAT] ' T= 298.15 K,I=n/a -13.63 [1960NAS/MER] T= 298.15 K, I=n/a -13.1 [1963FEI/SCH] T= 298.15 K, 1=0 , -27.99 [1977PAU] T= 298.15 K,I=n/a -12.99 [1980BEN/TEA] T= 298.15 K,I=n/a -13.63 [1982PAU] T= 298.15 K, I=n/a -12.98 [1982WAG/EVA] T= 298.15 K,I=n/a -13.64 [1985BAB/MAT] T= 298.15 K,I=n/a -13.63 [1985GAL] T= 298.15 K,I=n/a -12.91 [1988PHI/HAL] T= 298.15 K,I=n/a 2+ + log K*so: Pb + 3H2O <=> PbOPb(OH)2 + 4H -27.10 [1976SMI/MAR] T= 298.15 K, I=n/a 2+ + • K*so: Pb + 2H2O <=> Pb(OH)2 (crystalline, formed atT' = 300 °C) + 2H -8.11 [1971NAU/RYZ] T= 298.15 K, I=n/a 0 n/a -7.74 [1976NRI] T= 298.15 K, I=n/a 0 n/a -8.14 [1980BEN/TEA] T= 298.15 K,I=n/a 0 n/a -8.13 [198OSCH] T = 298.15 K, 1=0 0 n/a -8.14 [1981STU/MOR] T= 298.15 K, I=n/a 0 n/a -8.15 [1982WAG/EVA] T= 298.15 K, I=n/a 0 n/a -8.08 [1988PHI/HAL1 T= 298.15 K, I=n/a 0 n/a 147 JNC TN8400 99-011 Table 6.6: continued 2+ + log K*so: Pb + H2O t=>PbO (crystallinity not defined) + 2H -12.88 [1964HIR] T= 298.15 K, I=n/a -12.71 [1973BAR/KNA] T= 298.15 K, I=n/a -12.72 [1981BAE/MES] T= 298.15 K, I=n/a -12.90 [1981STU/MOR] T= 298.15 K, I=n/a -12.63 [1985CHA/DAV] T= 298.15 K, I=n/a -8.93 [1987BRQAVAN1 T= 298.15 K, I=n/a 1 [1979PAT/OBR] selected thermodynamic data for lead and then tested their dataset with tested with experimental data from different sources. 148 JNC TN8400 99-011 6.3 Lead chloride system 6.3.1 Lead chloride complexes The existing literature on lead chloride complexes is extensive. Comparison of the results reported by different workers is complicated by the fact that the experimentally determined values depend on the composition of the respective medium [1984BYR/MIL, 1984MHVBYR]. For the present report, experimental values determined in chloride or perchlorate media (Table 6.7) were extrapolated with the SIT equation to I = 0 (Figures 6.9, 6.10, 6.11 and 6.12): Pb2+ + Cr <=> PbCl+ log p°,,i = 1.55 2+ Pb + 2Cr <=} PbCl2° log p°i)2 = 2.00 2+ Pb + 3Cr <=> PbCl3- log P°i>3 = 2.01 2+ 2 Pb + 4CT <=> PbCl4 - log P°i,4 = 1.35 A number of investigators extrapolated the formation constants of lead chloride complexes to 1=0. Some of these data have also been included in the calculation of the formation constant to expand the data basis (see Table 6.7). The included data have been either determined at low ionic strength [1930RIG/DAV, 1955PIG/PAR, 1955NAN, 1992LOZ/SCH] where the method of extrapolation to 1=0 is expected to have only a small influence, or are based on a large number of experiments which have been carried out at different ionic strength [1984BYR/MEL, 1980MEL/BYR, 1984SEW, 1991MAG/FUE]. These results have been extrapolated by the respective authors to 1=0 using the Pitzer equation. These data were also included in our calculations assuming that ideally at 1=0, the same formation constant is obtained with the Pitzer correction as with the SIT approach. The omission of all constants reported at 1=0 would result inlogP°i,i, p°i)2, P°i,3, and P°i i4 values of 1.59, 2.10, 2.07 and 1.26, respectively. The log P values for the formation of Pb chloride'complexes determined in LiClO4 medium are generally larger than formation constants determined in NaQC>4 medium. Thus, the log P values are plotted in the SIT plot using: a) the log P values measured in all medium to determine log P°and Ae. b) only the log P values measured in LiClO4 to determine a Ac^cio value for the IJCIO4 medium. c) only the log P values measured in NaClO4 to determine a Ae^oo value for the NaClO4 medium. 149 JNC TN8400 99-011 Table 6.7: Experimentally determined equilibrium data compiled for the lead chloride system, 2+ 2m n according to the equilibrium: mPb + nCl" <=> Pbm(Cl)n " . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: con = conductivity measurements, el = electrophoresis, kin = kinetic measurements, pol = polarography, pot = potentiometry, sol = solubility measurements, and sp = spectrophotometry. Reference Comments .(M> Medium Method log Pm, n 2+ b« A.r-pb +cr<=>Pbcr 1.52 i [1930RIG/DAV] T= 298.15 K, 1=0.001-0.02 0 NaCl con 1.57 2, 3 [1955BIG/PAR] T= 298.15 K, 1=0.0.02-0.1 0 Pb(C104)2 sp 3 1.59 [1955NAN] T= 298.15 K, 1=0.001 0 NaClO4 con 4 1.23 [1964MER/KUL] T= 298.15 K, 1=3 3 LiClO4 sol 1.11 [1972FED/SHI] T= 298.15 K, 1=0.1 0.1 LiClO4 pot 0.90 [1972FED/SHI] T= 298.15 K, 1=0.5 0.5 LiClO4 pot 0.85 [1972FED/SHI] T= 298.15 K, 1=1 1 LiC104 pot 1.04 [1972FED/SHI] T- 298.15 K, 1=2 2 LiC104 pot 1.15 [1972FED/SHI] T= 298.15 K, 1=3 3 LiC104 pot 1.34 [1972FED/SHI] T= 298.15 K, 1=4 4 LiClO4 pot 1.28 5 [1972VBE] T= 298.15 K, 1=4 4 NaClO4 pot 0.94 6 [1973BON/HEF] T= 298.15 K, 1=1 1 NaClO4 pot 1.23 [1973HUT/HIG] T= 298.15 K, 1=0.1 0.1 NaClO4 kin 1.08 [1973HUT/HIG] T= 298.15 K, 1=1 1 NaClO4 kin 0.93 7 [1977SIP/VAL] T= 298.15 K, 1=6.7 0.7 NaCl pol 0.83 8 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 8 0.69 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 0.90 8 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 0.61 8 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaC104 pol 0.70 8 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaClO4 pol 1.59 3 [1980PRA/PRA] T- 298.15 K, 1=0.01-0.03 0 HC1 pot 0.91 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1/HC1O4 sp 0.84 [1981BYR/YOU] T= 298.15 K, 1=1 1 MgCl2 sp 0.92 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1 sp 0.89 [1982BEN/MEU] T= 298.15 K, 1=0.5 0.5 NaClO4 sp 0.98 [1982BEN/MEU] T- 298.15 K, 1=1 1 NaClO4 sp 0.82 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaC104 sp 0.83 [1982BEN/MEU] T= 298.15 K, 1=0.5 0.5 NaClO4 pot 0.85 [1982BEN/MEU] T= 298.15 K, 1=1 1 NaC104 pot 0.92 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaC104 pot 1.05 [1982BEN/MEU] T= 298.15 K, 1=3 3 NaClO4 pot 1.15 [1982BEN/MEU] T= 298.15 K, 1=4 4 NaC104 pot 150 JNC TN8400 99-011 Table 6.7: continued 1.00 [1982ROH] T= 296 K, 1=0.7 0.7 NaC104 el 1.33 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 HC1/HC1O4 sp 1.54 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 NaCl/NaC104 sp 1.37 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 MgCl2 sp 1.38 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 CaCl2 sp 1.14 [1984LOZ/SCH] T= 298.15 K, 1=1 1 HC1O4 sol 1.11 [1984LOZ/SCH] T= 298.15 K, 1=0.5 0.5 HC1O4 sol 1.15 [1984LOZ/SCH] T= 298.15 K, 1=0.2 0.2 HC1O4 sol 1.48 1° [1984ML/BYR] T= 298.15 K, 1=0 0 different sp 1.41 n [1984SEW] T= 298.15 K, 1=0 0 NaCl, HC1 sp 1.60 i2 [1991MGA/FUE] T= 298.15 K, 1=0-6 0 NaCl sol 1.60 3 [1992LOZ/SCH] T= 298.15 K, I=dil 0 PbCl, con 2+ b + 2Cr t=> PbCl2° 4 1.87 [1964MIR/KUL] T= 298.15 K, 1=3 3 LiC104 sol 1.56 [1972FED/SHI] T= 298.15 K, 1=0.1 0.1 LiC104 pot 1.30 [1972FED/SHI] T= 298.15 K, 1=0.5 0.5 LiC104 pot 1.26 [1972FED/SHI] T= 298.15 K, 1=1 1 LiC104 pot 1.40 [1972FED/SHI] T= 298.15 K, 1=2 2 LiC104 pot 1.70 [1972FED/SHI] T= 298.15 K, 1=3 3 LiC104 pot 2.06 [1972FED/SHI] T= 298.15 K, 1=4 4 LiC104 pot 5 1.57 [1972 VIE] T= 298.15 K, 1=4 4 NaC104 pot 1.08 6 [1973BON/HEF] T= 298.15 K, 1=1 1 NaC104 pot 1.35 ^ [1977SIP/VAL] T= 298.15 K, 1=0.7 0.7 NaCl pol 8 1.19 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaC104 pol 8 1.10 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaC104 pol 1.26 8 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 8 1.25 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaC104 pol 1.38 8 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaC104 pol 1.21 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1/HC1O4 sp 1.06 [1981BYR/YOU] T= 298.15 K, 1=1 1 MgCl2 sp 1.23 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1 sp 1.13 [1982BEN/MEU] T= 298.15 K, 1=0.5 0.5 NaC104 sp 1.30 [1982BEN/MEU] T= 298.15 K, 1=1 1 NaClO4 sp 1.33 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaClO4 sp 1.34 [1982BEN/MEU] T= 298.15 K, 1=0.5 0.5 NaC104 pot 1.24 [1982BEN/MEU] T= 298.15 K, 1=1 1 NaC104 pot 1.27 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaC104 pot 1.53 [1982BEN/MEU] T= 298.15 K, 1=3 3 NaClO4 pot 1.94 [1982BEN/MEU] T= 298.15 K, 1=4 4 NaClO4 pot 1.04 [1982ROH] T= 296 K, 1=0.7 0.7 NaClO4 el 9 1.76 [1984BYR/MIL] T= 298.15 K, 1=0 0 HC1/HC1O4 sp 9 2.08 [1984BYR/MIL] T= 298.15 K, 1=0 0 NaCl/NaClO4 sp 151 JNC TN8400 99-011 Table 6.7: continued 9 1.77 [1984BYR/MIL] T= 298.15 K, 1=0 0 MgCl2 sp 9 1.81 [1984BYR/MIL] T= 298.15 K, 1=0 0 CaCl2 sp 1.01 [1984LOZ/SCH] T= 298.15 K, 1=1 1 HC1O4 sol 0.92 [1984LOZ/SCH] T= 298.15 K, 1=0.5 0.5 HC1O4 sol 0.97 [1984LOZ/SCH] T= 298.15 K, 1=0.2 0.2 HC1O4 sol 2.03 10 [1984MDDL/BYR] T= 298.15 K, 1=0 0 different sp 1.97 n [1984SEW] T= 298.15 K, 1=0 0 NaCl, HC1 sp 1.66 12 [1991MGA/FUE] T= 298.15 K, 1=0-6 0 NaCl sol 1.74 3 [1992LOZ/SCH] T= 298.15 K,I=dil 0 PbCL con log /3!3: PI* +set etna; 4 1.98 [1964MIR/KUL] T= 298.15 K, 1=3 3 LiC104 sol 1.20 [1972FED/SHI] T= 298.15 K, 1=1 1 LiClO4 pot 1.40 [1972FED/SHI] T= 298.15 K, 1=2 2 LiClO4 pot 1.95 [1972FED/SHI] T= 298.15 K, 1=3 3 LiClO4 pot 2.40 [1972FED/SHI] 298.15 K, 1=4 4 LiClO4 pot 5 2.28 [1972VIE] T= 298.15 K, 1=4 4 NaC104 pot 1.72 6 [1973BON/HEF] T= 298.15 K, 1=1 1 NaC104 pot 8 0.86 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaC104 pol 8 0.75 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 8 0.74 [1980LOV/BRA] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 8 1.71 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaC104 pol 1.79 8 [1980LOV/BRA] T= 298.15 K, 1=3 3 NaClO4 pol 1.16 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1/HC1O4 sp 0.92 [1981BYR/YOU] 298.15 K, 1=1 1 MgCl2 sp 1.18 [1981BYR/YOU] T= 298.15 K, 1=1 1 HC1 sp 1.17 [1982BEN/MEU] T= 298.15 K, 1=1 1 NaC104 sp 1.23 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaC104 sp 1.09 [1982BEN/MEU] T= 298.15 K, 1=1 1 NaClO4 pot 1.43 [1982BEN/MEU] T= 298.15 K, 1=2 2 NaC104 pot 1.83 [1982BEN/MEU] T= 298.15 K, 1=3 3 NaClO4 pot 1.97 [1982BEN/MEU] T= 298.15 K, 1=4 4 NaClO4 pot 1.25 [1982ROH] T= 296 K, 1=0.7 0.7 NaC104 el 1.72 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 HC1/HC1O4 sp 2.58 9 [1984BYR/MIL] T= 298.15 K, 1=0 0 NaCl/NaC104 sp 9 1.72 [1984BYR/MIL] T= 298.15 K, 1=0 0 MgCl2 sp 9 1.75 [1984BYR/MIL] T= 298.15 K, 1=0 0 CaCl2 sp 1.52 [1984LOZ/SCH] T= 298.15 K, 1=1 1 HC1O4 sol 1.57 [1984LOZ/SCH] T= 298.15 K, 1=0.5 0.5 HC1O4 sol nn 1.79 [1984LOZ/SCH] 298.15 K, 1=0.2 0.2 HC1O4 sol 1.86 1° [1984MIL/BYR] T= 298.15 K, 1=0 0 different sp 1.66 ii [1984SEW] T= 298.15 K, 1=0 0 NaCl, HC1 sp 2.27 12 [1991MGA/FUE] T= 298.15 K,1=0-6 0 NaCl sol 152 JNC TN8400 99-011 Table 6.7: continued log A/ Pb2+ + 4CI <=> PbCl/' 4 1.72 [1964MIR/KUL] T= 298.15 K, 1=3 3 LiC104 sol 0.85 [1972FED/SHI] T= 298.15 K, 1=2 2 LiC104 pot 1.18 [1972FED/SHI] T= 298.15 K, 1=3 3 LiC104 pot 1.90 [1972FED/SHI] T= 298.15 K, 1=4 4 LiClO4 pot 5 1.43 [1972VIE] T= 298.15 K, 1=4 4 NaC104 pot 1.46 ii [1984SEW] T= 298.15 K, 1=0 0 NaCl, HC1 sp 1.27 12 [1991MGA/FUE] T= 298.15 K, 1=0-6 0 NaCl sol 2+ log KSo: Pb + 2Ct & PbCl2(s) 4 5.00 [1964MIR7KUL] T= 298.15 K, 1=3 3 LiC104 pot 4.78 4 [1964MIR/KUL] T= 298.15 K, 1=3 3 LiCl pot 4.11 [1984LOZ/SCH] T= 298.15 K, 1=1 1 HC1O4 sol 4.01 [1984LOZ/SCH] T= 298.15 K, 1=0.5 0.5 HC1O4 sol 4.07 [1984LOZ/SCH] T= 298.15 K, 1=0.2 0.2 HC1O4 sol 2.80 I2 [1991MGA/FUE] T= 298.15 K, 1=0-6 0 NaCl sol 2+ log Kso: Pb + H2O + Cl t=>PbOHCl(cr) + Ct -1.27 13 [1976NAS/LIN] T= 298.15 K, 1=:0.052 0.05 NaC104 sol 13 -1.33 [1976NAS/LIN] T= 298.15 K, 1==0.554 0.55 NaC104 sol 13 -1.33 [1976NAS/LIN] T= 298.15 K, 1==1.06 1.06 NaC104 sol 13 -1.28 [1976NAS/LIN] T= 298.15 K, 1==1.08 1.08 NaClO4 sol 13 -0.82 [1976NAS/LIN] T= 298.15 K, 1==2.078 2.08 NaC104 sol 14 -0.88 [1976NAS/LIN] T= 298.15 K, 1=:0.0158 0.016 NaClO4 sol 14 -1.25 [1976NAS/LIN] T= 298.15 K, 1==0.029 0.029 NaC104 sol 14 -1.31 [1976NAS/LIN] T= 298.15 K, 1==0.53 0.53 NaC104 sol 14 -1.26 [1976NAS/LIN] T= 298.15 K, 1==1.033 1.033 NaC104 sol 14 -0.62 [1976N AS/LIN] T= 298.15 K, 1=2.031 2.031 NaCIO, sol corrected to I = 0 by [193ORIG/DAV] with a simplified Debye-Huckel equation 50-600 mM Pb corrected to 1=0 with the Davies equation by the respective authors only results used where < 10% LiCl present the reported constants are dependent on the Cl concentration Pb concentration = 0.19 mM Pb concentration = 0.001 mM Pb concentration = 2E-8 M corrected to 1=0 with extended Debye-Huckel equation with parameters were fitted by [1984BYR/MIL] 10 corrected to 1=0 with Pitzer's equation by [1984MIL/BYR], same experimental data as [1984BYR/MIL]. 11 Pb = 0.1 mM, extrapolated to 1=0 with Pitzer equation by [1984SEW]. Also measurements at higher temperature, AH, AS, ACp given. 12 corrected to 1=0 with Pitzer's equation by [1991MGA/FUE]. 13 measured in presence of lead nitrate 14 measured in presence of lead chloride 153 JNC TN8400 99-011 Pb2+ + Cl- <=> PbCI+ 2+ Figure 6.9: Plot of log p\i + 4 D vs. Im for the reaction : Pb + Cl- <=> PbCl+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.12; log f3\i = 1.55; AeLjci04 = -0.16 (in LiC104), AeNaCiO4 = -0.10 (in NaC104). Calculated from data compiled in Table 6.7. 2+ Pb + 2Ch <=> PbCI2° 5 - 4.5 - ] s 1; : : : ; ; 1 t : "'•.'''"•'~".'"".\''" . '-'•• ;"*' T-;-!!: :?D '-:;' '' '*:::ii i,!!r'j;'; ;.'?l'::'^i^.H-ii'iisi^^.^'v.-vii^^i/:^:'i;;":: ':;-.•• • -":-::-it.:j"T''^^•;••;"-; • : 1 ; ; : '.'.'. ;••••;;>.;•':-.»"i •i/vv:-:!: j.r'; :;:;iv. l;^i';^F:::-\"i^:tj,:'.-L:^^ ?:™ :'i: ':!i:i - i': ?;••:;: :i::-. • .:•••.!.:..»•;•:•.: " . :-\- •••.: ;••:; 4 1 ;i i : : - = : : '::;.'L:. .-"• •''^'"ii.-t)-;;^' ^' "'.-vj:.'; l' !7" ^/^^^*:iii :i ': -|=. ":i "-'-ri T" ^ r r'"" l- H"| ^ | : f.; -Lf.r i r-'- •;•;! : '.'..\- -'- ':•.'•:.' •. .'. '•;;".• Q 3.5 CD + 3 T;i?|i^fsj..::..i v. .1;::::i'|^W;i,•,,;•; CO. 2.5 : ;: 2| lo g 1.5 - iii^-v':fh^,i'i5t.;:V:;iK 1 : ; r ,-«*«-•;.:• !• ''iffi.::-.:;-: >:, ;:. .v'n>:;s;v-^^py.:^^:,,- • ;;.,.,-•• .:•,:-.;.:••.:•:. : : ! ; : : s :: 1 •I'-.?.:.-; ;•••;;•.-,• :.i:;:;2«' "i.::.,: a:?KW' pra|iftt?«K"' ^ ; - s-p'.*: ; :^:« «i «•«;- : . 0.5 0 - 2 4 lm, molal 2+ Figure 6.10: Plot of log (5li2 + 6 D vs. Im for the reaction : Pb + 2C1- <=> PbCl2° at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.27; log P°i2 = 2.00; AeLiCiO4 = -0.32 (in LiC104), AeNaciO4 = -0.24 (in NaC104). Calculated from data compiled in Table 6.7. 154 JNC TN8400 99-011 2+ Pb + 3CI- & PbCI3 4.5 -• 4 - LiCIO4 y = 0.35x + 2.01 NaCIO4 = 0.29x + 2.01 y = 0.31x+2.01 2 4 lm> molal 2+ Figure 6.11: Plot of log pii3 + 6 D vs. Im for the reaction : Pb + 3C1- <=> PbCl3- at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.31; log P°i 3 = 2.01; AeLiciO4 = -0.35 (in L1CIO4), AeNaCiO4 = -0.29 (in NaC104). Calculated from data compiled in Table 6.7. 2+ 2 Pb + 4CI- <^ PbCI4 " lm, molal 2+ 2 Figure 6.12: Plot of log p1>4 + 4D vs. Im for the reaction : Pb + 4C1" <=> PbCl4 - at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.21; log (3\4 = 1.35; AeLiCiO4 = -0.24 (in LiC104), AeNaCiO4 = -0.15 (in NaClO4). Calculated from data compiled in Table 6.7. 155 JNC TN8400 99-011 6.3.2 PbCl2(s) Lead also forms salts with chloride. Extrapolation of the values given by [1964MIR/KUL, 1984LOZ/SCH and 1991MGA/FUE] (see Table 6.7) is shown in Figure 6.13: 2+ Pb + 2C1" PbCl2(s) log K*°so = 4.81 This value agrees well with the value of 4.77 proposed by [1980CLE/JOH], who made a careful and extensive review of the solubility of lead salts. 2+ Pb +2CI" « PbCI2(s) jij4p|ilMga|g||a||p;S;ii|S|];|S|;: 2+ Figure 6.13: Plot of log K*So + 6D vs. Im for the reaction : Pb + 2C1- <=> PbCl2(s) at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.44; log K*°So = 4.81. Calculated from data compiled in Table 6.7. 6.3.3 Mendipite: Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr) Based on data given in [1980BEN/TEA] and [1980MAN], log K*So for mendipite (Pb2(OH)3Cl(cr) or Pb4(OH)6Cl2(cr)) is approximately -8. Primary experimental data however are not available. Thus, no solubility product for mendipite is recommended in this report. 156 JNC TN8400 99-011 6.3.4 Laurionite andparalaurionite For laurionite (PbOHCl(cr)) a log K*so of 15.36 can be calculated from the thermodynamic data given in [1971NAU/RYZ] and [1988PHI/HAL]. [1987BRU], however, writes that this value is probably based on an erroneous copying (AjG° -480.3 instead of -408.3) of a value which already originally was too low. Experimental results of the formation of PbOHCl(cr) are given in the careful work of [1976NAS/LIN] (Table 6.7). [1976NAS/LIN] do not state if they use laurionite or paralaurionite in their experiments. [1987BRU], however, states that the difference in solubility between laurionite (PbCl2-Pb(OH)2(cr)) and paralaurionite (PbCl2-PbO-H2O(cr)) is too small to make any apparent difference. Extrapolation of the measurements of [1976NAS/LIN] to 1=0 gives (Figure 6.14): 2+ *o_ Pb + Cl- + H2O PbOHCl(cr) + logK so = - 0.62 This value is similar to the value of -0.29 given in different compilations [1982WAG/EVA, 1983LAN, 1983SAN/BAR] for laurionite or paralaurionite. The formation of lead complexes with other halides (bromide, iodide) has also been observed in different studies, but is not addressed in the present report. 2+ + Pb +C|-+H2O<^>PbOHCI(cr)+H 1 2 3 lmi molal 2+ Figure 6.14: Plot of log K*s0 + 4D vs. Im for the reaction : Pb + Cl- + H2O <=> PbOHCl(cr) + H+ at 25 °C. The straight line shows the result of the linear regression: Ae = -0.11; log K*°so = -0.62. Calculated from data compiled in Table 6.7. 157 JNC TN8400 99-011 6.3.5 Additional equilibrium data compiled for the lead chloride system Table 6.8: Additional experimentally determined data for the lead chloride system, according to the 2+ 2m n equilibrium: mPb + nCl" <=> Pbm(Cl)n ' . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are indicated at the end of the table. Method: con = conductivity measurements, el = electrophoresis, pol = polarography, pot = potentiometry, sol = solubility measurements, and sp = spectrophotometry. log ?„ Reference Comments Medium Method log Pb2+ + Cl <=> PbCl* 1.57 ' [1955BIG/PAR] T= 291.15 K, 1=0 0 n/a n/a 1.16 2 [1971BON1] T= 298.15 K, 1=0.2-1.8 Cl- pol 1.16 > [1970FED/SAM] T= 298.15 K, 1=3 3 LiClO4 n/a 3 1.52 [1972FED/SHI] T= 298.15 K, 1=0 0 LiC104 pot 4 1.48 [1972FED/SHI] T= 298.15 K, 1=0 0 LiClO4 pot 5 0.89 [1982BEN/MEU] T= 298.15 K, 1=0 0 NaC104 pot 1.61 fl984LOZ/SCH] T= 298.15 K, 1=0 0 HC1O, sol 2+ log pu: Pb h 2CI & PbCl2° 1.26 2 [1971BON1] T= 298.15 K, 1=0.2-1.8 Cl- pol 3 2.19 [1972FED/SHI] T= 298.15 K, 1=0 0 LiClO4 pot 4 2.08 [1972FED/SHI] T= 298.15 K, 1=0 0 LiClO4 pot 5 1.32 [1982BEN/MEU] T= 298.15 K, 1=0 0 NaC104 pot 6 1.67 [1984LOZ/SCH] T= 298.15 K, 1=0 0 HC1O4 sol 2+ log /5U: Pb + 3CV <=> PbClj x 1.94 [1960FRI/SAR] T= 298.15 K, 1=5 NaC104 n/a 1.45 2 [1971BON1] T= 298.15 K, 1=0.2-1. Cl- pol 3 1.97 [1972FED/SHI] T= 298.15 K, 1=0 0 LiC104 pot 4 1.81 [1972FED/SHI] T= 298.15 K, 1=0 0 LiC104 pot 5 0.83 [1982BEN/MEU] T= 298.15 K, 1=0 0 NaC104 pot 2.62 6 [1984LOZ/SCH1 T= 298.15 K, 1=0 0 HC1O, sol 2+ log p,4: Pb + 4Ci <=> 3 0.77 [1972FED/SHI] T= 298.15 K, 1=0 0 LiClO4 pot 0.85 4 [1972FED/SHI] T= 298.15 K, 1=0 0 LiClO, pot 2+ log Kso: Pb + 2CI PbCl2(s) 7 4.90 [1972VIE] T= 298.15 K, 1=4 4 NaC104 sol 6 4.77 H984LOZ/SCH1 T=298.15K, 1=0 0 HC1O4 sol 158 JNC TN8400 99 -Oil Table 6.8 : continued 2+ log Kso: Pb + H20 + Cl & PbOHCl(s) + Cl -0.62 [1976NAS/LIN] T= 298.15 K, 1=0 NaCIO, sol 1 no experimental details are known 2 I not constant -* extrapolated to 1=0 with SIT model 4 extrapolated to 1=0 with Vasilev equation (= Debye-Hiickel) 5 linear? extrapolated to 1=0 by [1982BEN/MEU] 6 corrected to 1=0 with SIT by [1984LOZ/SCH] 7 the reported solubility constant is dependent on Cl concentration 8 extrapolated to 1=0 with Debye-Hiickel Table 6.9: Thermodynamic data for the lead chloride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. Reference Comments KM) Medium 2+ + log pa: Pb + Ct <=> PbCl 1.60 [1969HEL] T= 298.15 K, 1=0 0 1.62 [1976HEM] T= 298.15 K,I=n/a 1.59 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.90 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.90 [1976SMI/MAR] T= 298.15 K, 1=1 1 1.02 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.17 [1976SMI/MAR] T= 298.15 K, 1=3 ^ 3 1.29 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.64 [1980BEN/TEA] T=298.15K, 1=0 0 1.58 [1981TUR/WHI] T= 298.15 K, 1=0 0 1.64 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.64 [1983LAN] T= 298.15 K, 1=0 0 1.32 [1987BRO/WAN] T= 298.15 K, 1=0 0 1.53 i [1987BRU] T= 298.15 K, 1=0 0 0.94 ' [1987BRU] T= 298.15 K, 1=1 1 1.17 l [1987BRU] T= 298.15 K, 1=3 3 0.86 [1988BYR/KUM] T= 291 K, 1=0.7 0.7 sea water 1.37 [1988PHI/HAL] T= 298.15 K, 1=0 0 1.55 [1989SMI/MAR] T= 298.15 K, 1=0 0 1.08 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 0.87 [1989SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.86 [1989SMI/MAR] T= 298.15 K, 1=0.7 0.7 0.90 [1989SMI/MAR] T= 298.15 K, 1=1 1 1.00 [1989SMI/MAR] T= 298.15 K, 1=2 2 1.12 [1989SMI/MAR] T= 298.15 K, 1=3 3 1.23 [1989SMI7MAR] T= 298.15 K, 1=4 4 159 JNC TN8400 99-011 Table 6.9: continued 2+ log pu: Pb + 2CV <=> PbCl2° 1.78 [1969HEL] T= 298.15 K, 1=0 0 2.44 [1976HEM] T= 298.15 K,I=n/a 1.80 [1976SM1/MAR] T= 298.15 K, 1=0 0 1.30 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 1.30 [1976SMI/MAR] T= 298.15 K, 1=1 1 1.40 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.70 [1976SMI/MAR] T= 298.15 K, 1=3 3 2.00 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.82 [1981TURAVHI] T= 298.15 K, 1=0 0 1.84 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.84 [1983LAN] T=298.15K, 1=0 0 1.57 [1987BRO/WAN] T= 298.15 K, 1=0 0 1.93 ' [1987BRU] T= 298.15 K, 1=0 0 1.16 ' [1987BRU] T= 298.15 K, 1=1 1 1.72 ! [1987BRU] T= 298.15 K, 1=3 3 1.16 [1988BYR/KUM] T=291 K, 1=0.7 0.7 sea water 1.93 [1988PHI/HAL] T= 298.15 K, 1=0 0 2.20 [1989SMI/MAR] T=298.15K, 1=0 0 1.40 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 1.20 [1989SMI/MAR] T= 298.15 K, 1=0.5 0.5 1.20 [1989SMI/MAR] T= 298.15 K, 1=0.7 0.7 1.20 [1989SMI/MAR] T= 298.15 K, 1=1 1 1.30 [1989SMI/MAR] T= 298.15 K, 1=2 2 1.60 [1989SMI/MAR] T= 298.15 K, 1=3 3 1.80 [1989SMI/MAR] T= 298.15 K, 1=4 4 2+ log Pu: Pb + 3CI <=> PbCli 1.68 [1969HEL] T= 298.15 K, 1=0 0 2.04 [1976HEM] T= 298.15 K,I=n/a 1.70 [1976SMI/MAR] T= 298.15 K, 1=0 0 1.40 [1976SMI/MAR] T= 298.15 K, 1=1 1 1.50 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.90 [1976SMI/MAR] T= 298.15 K, 1=3 3 2.30 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.47 [1980BEN/TEA] T= 298.15 K, 1=0 0 1.71 [1981TUR/WHI] T= 298.15 K, 1=0 0 1.47 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.76 [1983LAN] T= 298.15 K, 1=0 0 0.88 [1987BRO/WAN] T= 298.15 K, 1=0 0 1.84 ' [1987BRU] T= 298.15 K, 1=0 0 1.19 > [1987BRU] T= 298.15 K, 1=1 1 1.95 ' [1987BRU] T= 298.15 K, 1=3 3 1.06 [1988BYR/KUM] T=291 K, 1=0.7 0.7 sea water 1.60 [1988PHI/HAL] T= 298.15 K, 1=0 0 1.80 [1989SMI/MAR] T= 298.15 K, 1=0 0 1.30 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 1.10 [1989SMI/MAR] T= 298.15 K, 1=0.5 0.5 1.10 [1989SMI/MAR] T= 298.15 K, 1=0.7 0.7 1.10 [1989SMI/MAR] T= 298.15 K, 1=1 1 1.40 [1989SMI/MAR] T= 298.15 K, 1=2 2 1.90 [1989SMI/MAR] T= 298.15 K, 1=3 3 2.20 [1989SMI/MAR] T= 298.15 K, 1=4 4 160 JNC TN8400 99-011 Table 6.9: continued 2+ log PM: Pb + 4CI <=> PbCl/- 1.38 [1969HEL] T= 298.15 K, 1=0 0 1.40 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.70 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.20 [1976SMI/MAR] T= 298.15 K, 1=3 3 1.70 [1976SMI/MAR] T= 298.15 K, 1=4 4 1.40 [1981TUR/WHI] T= 298.15 K, 1=0 0 1.46 [1983LAN] T= 298.15 K, 1=0 0 -0.71 [1987BRO/WAN] T= 298.15 K, 1=0 0 1.24 : [1987BRU] T= 298.15 K, 1=3 3 1.38 [1988PHI/HAL] T= 298.15 K, 1=0 0 1.10 [I989SMI/MAR] T= 298.15 K, 1=0 0 0.70 [1989SMI/MAR] T= 298.15 K, 1=2 2 1.00 [1989SMI/MAR] T= 298.15 K, 1=3 3 1.40 f 1989SM1/MAR1 T= 298.15 K, 1=4 4_ 2+ log Kso: Pb + 2Ct « PbCl2(s) 4.78 [1952LAT] T= 298.15 K, I=n/a 4.80 [1963W1C/BLO] T= 298.15 K, I=n/a 5.02 [1968ROB/WAL] T= 298.15 K, I=n/a 4.80 [1971NAU/RYZ] T= 298.15 K, I=n/a 5.04 [1973BAR/KNA] T= 298.15 K, I=n/a 4.78 [1976SMI/MAR] T= 298.15 K, 1=0 0 5.00 [1976SMI/MAR] T= 298.15 K, 1=3 3 4.78 [1977PAU] T= 298.15 K, I=n/a 4.79 [1978ROB/HEM2] T= 298.15 K, I=n/a 4.81 [1979KUB/ALC] T= 298.15 K, I=n/a 4.79 [1980BEN/TEA] T= 298.15 K, I=n/a 4.77 [1980CLE/JOH] T=298.15 K, I=n/a 4.67 [1980MAN/DEU] T=298.15 K, 1=0 , 0 4.78 [1982PAU] T= 298.15 K, I=n/a 4.81 [1982WAG/EVA] T= 298.15 K, I=n/a 4.82 [1983LAN] T= 298.15 K, I=n/a 4.82 [1983SAN/BAR] T= 298.15 K, I=n/a 4.81 [1985CHA/DAV] T= 298.15 K, I=n/a 4.78 [1985GAL] T= 298.15 K, I=n/a _4J?0 ri988PHKHAL] T= 298.15 K, I=n/a 2+ + log Kso: Pb + H2O + Cl- <=> PbOHCl(s) + H -1.00 [1973BAR/KNA] T= 298.15 K, I=n/a -0.30 [1980CLE/JOH] T=298.15K, I=n/a -0.28 [1982PAU] T=298.15K, I=n/a -0.29 [1982WAG/EVA] T= 298.15 K, I=n/a 2+ log Kso: Pb + H2O + Cl <=> PbOHCl(laurionite) + Cl 15.37 [1971NAU/RYZ] T= 298.15 K, I=n/a -0.29 [1983SAN/BAR] T= 298.15 K, I=n/a 15.36 [1988PHI/HAL1 T= 298.15 K,I=n/a 161 JNC TN8400 99-011 Table 6.9: continued 2+ log Km: 2Pb + 3H2O + Ci <=> Pb2(OH)3Cl(s), mendipite -8.90 2 [1980BEN/TEA] T= 298.15 K, I=n/a -8.00 3 [1980MAN/DEU] T=298.15 K, 1=0 2+ log Ks0: 4Pb + 6H2O + 2Ct <=> (Pb(OH)2)3PbCl2(s) -17.59 [1982WAG/EVA1 T= 298.15 K, I=n/a 1 mean calculated by [1987BRU] from data from different sources. 2 mendipite (Pb2(OH)3Cl(cr)) solubility as given in [1987BRU], 3 data selected from literature, corrected to 1=0 by [1980MAN/DEU]. 162 JNC TN8400 99-011 6.4 Lead fluoride system In the presence of fluoride, lead forms different complexes. The complex formation constants in binary lead fluoride system are well documented (Table 6.10). Table 6.10: Experimentally determined equilibrium data compiled for the lead fluoride system, 2+ 2m n according to the equilibrium: mPb + nF- <=> Pbm(F)n " . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: fe = fluoride selective electrode, pol = polarography, pot = potentiometry, sol = solubility measurements. log Kn Reference Comments KM, Medium Method log fa,: Pb2+ + F- <=>PbF+ 1.26 [1963MES/HUM] T=291 K, 1=2 2 NaClO4 pol 1.48 1 [1965BOT/CIA] T= 298 K, 1=1 1 NaC104 pol 1.53 2 [1970BON/HEF] T= 288 K, 1=1 1 NaC104 pol 1.62 2 [1970BON/HEF] T=288K,I=1 1 NaC104 fe 2 1.73 [1970BON/HEF] T= 288 K, 1=0.1 0.1 NaC104 fe 3 1.40 [1971BON2] T= 298 K, 1=1 1 NaClO4 fe 4 1.40 [1972HEF] T= 298 K, 1=1 1 NaC104 fe 1.46 [1973BON/HEF] T= 298 K, 1=1 1 NaClO4 fe -* 2+ log fa,: Pb + 2F- <=> PbF2° 2.55 [1963MES/HUM] T= 291 K, 1=2 2 NaC104 pol 2 2.59 [1970BON/HEF] T= 288 K, 1=1 1 NaClO4 pol 2.52 3 [1971BON2] T= 298 K, 1=1 1 NaClO4 fe 2.52 [1973BON/HEF] T= 298 K, 1=1 1 NaClO4 fe 2+ log Kso: Pb + 2F- <^>PbF2(s) 7.43 5 [1963MES/HUM] T= 298 K, 1=0 0 n/a pot 6.60 [1963MES/HUM] T=291 K, 1=2 2 NaC104 sol 6 7.44 [1963MES/HUM] T=291 K, 1=0 0 NaClO4 sol 7 7.67 [1981 CIA] T= 298 K, 1=0 0 NaClO4 sol 7 6.60 [1981 CIA] T= 298 K, 1=1.05 1.05 NaClO4 sol Pb < 10 mM Precision lower than in other studies of Bond and Hefter [1971BON2, 1972HEF, 1973BON/HEF]. Pb concentration = 10 mM. Pb concentration = 0.4 mM. Pb concentration = 31 mM. recalculation of a value given in [1947BRO/DEV] extrapolation of own measurements to 1=0 no experimental details given 163 JNC TN8400 99-011 + 6.4.1 PbF andPbF2° The formation constants for Pb fluoride complexes reported in literature [1963MES/HUM, 1965BOT/CIA, 1970BON/HEF, 1971BON, 1972HEF, and 1973BON/HEF] (Table 6.10) are extrapolated to 1=0 ( Figures 6.15 and 6.16): pb2+ ,HF « PbF+ log p°i ,1 = 2.27 2+ Pb -h2F <=> PbF2° log P°i ,2 = 3.01 Pb2+ + F- a PbF+ 4.5- 4 -• Q 3.5 y ==-0.02x_+2.27 + 3 + o o g 1.5 - 1 0.5 0 0 0.5 1 1.5 2.5 lm, molal ^ 2+ Figure 6.15: Plot of log pu + 4 D vs. Im for the reaction : Pb + F- o PbF+ at 25 °C. The straight line shows the result of the linear regression: Ae = 0.02; log P°ii = 2.27. Calculated from data compiled in Table 6.10. 164 JNC TN8400 99-011 2+ Pb + 2F- <=> PbF2° 5 - 4.5 •.y'ljjllllplilil^filllllllllllilll 4 Q 3.5 |^-:;.il|||J^||||||f||||||:: CD + 3 CM -2.5 CO. 1.5 1 0.5 • n 0.5 1 1.5 2.5 lm, molal 2+ Figure 6.16: Plot of log (3U + 6 D vs. Im for the reaction : Pb + 2F- <=>-PbF2° at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.04; log (3°i 2 = 3.01. Calculated from data compiled in Table 6.10. 6.4.2 PbF2(s) 2+ Extrapolation of log K*So values for the reaction: Pb + 2F- <=> PbF2(s) from [1962MES/HUM] and [1981CIA] results in (Figure 6.17): 2+ Pb + 2F PbF2(s) log K*°so = 7.52 This value is in very good agreement with the value of 7.48 proposed by [1980CLE/JOH], who made a careful and extensive review of the solubility of lead salts. 165 JNC TN8400 99-011 2+ Pb + 2F- <=> PbF2 (s) 10 • 9.5 - 9 'r-;V''" „• •«" * V Q 8.5 - CD 8 - + J* * "' , ——-^ o 7.5 < 7 •-I K lo g 6.5 - _j;= 0:1^ + 7,52 " 6 - • -- •' \ 5.5- « • r~^* • ^^* ?\ ' t' ** ^' F. - Z , 1 1 , ( • 0.5 1 1.5 2.5 m, molal 2+ Figure 6.17: Plot of log K*So + 6 D vs. Im for the reaction : Pb + 2F" PbF2(s) at 25 °G. The straight line shows the result of the linear regression: Ae = - 0.19; log K*°so = 7.52. Calculated from data compiled in Table 6.10. Table 6.11: Thermodynamic data for the lead fluoride system taken from previous compilations. As pointed out in Section 2 of this report only experimental^ data were used for the present evaluation. The following table serves only for comparison. Reference Comments KM) 2+ + log pu:Pb +F <=>PbF 1.44 [1976SMI/MAR] T= 298.15 K, 1=1 1 1.26 [1976SMI/MAR] T= 298.15 K, 1=2 2 1.44 [1980BON/HEF] T=298K, 1=1 1 2.06 [1981TUR/WHI] T= 298.15 K, 1=0 0 1.44 [1982SMI/MAR] T= 298.15 K, 1=1 1 1.26 [1982SMI/MAR] T= 298.15 K, 1=2 2 0.77 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.01 [1985BAB/MAT] T= 298.15 K, 1=0 0 2.09 [1987BROAVAN] T= 298.15 K, 1=0 0 1.08 [1988PHI/HAL] T= 298.15 K, 1=0 0 166 JNC TN8400 99-011 Table 6.11: continued 2+ log pu: Pb + 2F- <=> PbF2° 2.54 [1976SMI/MAR] T= 298.15 K, 1=1 1 2.55 [1976SMI/MAR] T= 298.15 K, 1=2 2 2.53 [1980BON/HEF] T= 298 K, 1=1 1 3.42 [1981TUR/WHI] T= 298.15 K, 1=0 0 2.53 [1982SMI/MAR] T= 298.15 K, 1=1 1 2.55 [1982SMI/MAR] T= 298.15 K, 1=2 2 1.60 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.75 [1985BAB/MAT] T= 298.15 K, 1=0 0 3.85 [1987BRO/WAN] T= 298.15 K, 1=0 0 -0.89 [1988PHI/HAL] T= 298.15 K, 1=0 0_ 2+ log Pu: Pb + 3P «• PbFf 5.39 [1987BROAVAN] T= 298.15 K, 1=0 0 2.71 ri988PHI/HAL1 T= 298.15 K, 1=0 0_ 2+ 2 log p!4: Pb + 4F- <=> PbF4 - 6.78 ri987BROAVANl T= 298.15 K, 1=0 2+ log Kso: Pb + 2F- <=> PbF2(s) 7.64 [1973BAR/KNA] T= 298.15 K, I=n/a 7.44 [1976SMI/MAR] T= 298.15 K, 1=0 0 6.26 [1976SMI/MAR] T= 298.15 K, 1=1 1 6.60 [1976SMI/MAR] T= 298.15 K,I=2 2 7.48 [1980CLE/JOH] T= 298.15 K, 1=0 0 7.68 [1988PHI/HAL] T= 298.15 K, I=n/a 2+ log Kso: Pb + 2F- <=> a-PbF2 5.67 [1952LAT] T= 298.15 K, I=n/a 5.64 [1963WIC/BLO] T= 298.15 K, I=n/a 7.06 [1971NAU/RYZ] T= 298.15 K, I=n/a 7.54 [1979KUB/ALC] T= 298.15 K, I=n/a 7.60 [1982WAG/EVA] T= 298.15 K, I=n/a 7.54 [1985CHA/DAV] T= 298.15 K, I=n/a 7.64 [1985CHA/DAV] T= 298.15 K, I=n/a 5.64 [1985GAL] T= 298.15 K, I=n7a 167 JNC TN8400 99-011 6.5 Mixed lead fluoride chloride system Independently, [1971BON2] and [1972HEF] determined the formation constant of mixed PbCIF complexes in 1 M NaClC>4 medium and [1968AND] determined the solubility of PbClF(cr) (Table 6.12). Table 6.12: Experimentally determined equilibrium data compiled for the lead chloride fluoride 2+ 2m n system, according to the equilibrium: mPb + nQ~ + oF~ <=> PbmCl nF0 - "°. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: fe = fluoride selective electrode, pol = polarography,. log Pm,n Reference Comments Medium Method 2+ 0 log Pl2: Pb + CI-+ F- <=$ PbCIF ] 2.72 [1971BON2] T= 298 K, 1=1 1 NaC104 pol 2 2.78 [1972HEF] T= 298 K, 1=1 1 NaClO4 fe 2.87 2 [1972HEF] T= 298 K, 1=1 1 NaClO. fe 2+ log KSO: Pb + CI-+ F- ^PbClF(cr) 8.00 [1968 AND] T= 293 K, 1=0.6 0.6 NaClO4 sol 8.15 [1968AND] T= 293 K, 1=0.2 0.2 NaClO4 sol 8.02 [1968AND] T= 293 K, 1=0-0.2 0.25 NaClO4 sol 8.10 [1968 AND] T= 293 K, 1=0-0.2 0.2 NaClO4 sol 8.22 [1968 AND] T= 293 K, 1=0-0.2 0.1 NaClO4 sol 8.40 [1968 AND] T= 293 K, 1=0-0.2 0.025 NaC104 sol 8.54 [1968 AND] T= 293 K, 1=0-0.2 0.01 NaClO4 sol 8.64 [1968 AND] T= 293 K, 1=0-0.2 0.005 NaC104 sol 1 Pb concentration = 0.4 mM 2 Pb concentration = 1.8 mM 168 JNC TN8400 99-011 6.5.1 PbFCP Correction with the SIT term, assuming a Ae of -0.43 (as observed for PbFCl(cr) in the following section), yields a value of: PbFCl0 log = 3.55 Pb2+ + Cl- +F- PbCIF0 0.5 1 1.5 lm, molal 2+ 0 Figure 6.18: Plot of log pi,u + 6 D vs. Im for the reaction : Pb + Cl- + P <=> PbFCl at 25 °C. Ae = - 0.43 (see text); log $°vj,\ = 3.94. Calculated from data compiled in Table 6.12. 6.5.2 Matlockite (PbClF(cr)) Based on data given in different compilations (Table 6.14), a log K*so for matlockite (PbClF(cr)) of approximately 8.6 can be calculated. The solubility has been determined experimentally by [1968AND] (Table 6.12). Extrapolation to 1=0 is shown in Figure 6.19. + Cl- + F- <=> PbClF(matlockite) log K*°so= 8.82 169 JNC TN8400 99-011 Pb;2+ PbCIF(cr) 10 y = 0.43x + 8.82 5.5 +• , 5 o 0.2 0.4 0.6 0.8 lm, molal 2+ Figure 6.19: Plot of log K*so + 6 D vs. Im for the reaction : Pb + Cl- + F" o PbClF(matlockite) at 25 °C. The straight line shows the result of the linear regression: Ae = -0.43; log K*°so= 8.82. Calculated from data compiled in Table 6.12. Table 6.13: Additional experimentally determined data for the lead chloride fluoride system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility measurements. log K*so Reference Comments KM) Medium Method 2+ log K*so: Pb + CI-+ F- <=>PbClF(cr) 8.82 ' T1968AND1 T= 293 K, 1=0-0.2 0 NaClOi sol extrapolated to 1=0 as a function of I°-5/(l+I0-5) Table 6.14: Thermodynamic data for the lead chloride fluoride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K*so Reference Comments 2+ log K*so: Pb + Cl- + F- <=> PbClF(s) 8.37 [1971NAU/RYZ] T= 298.15 K,I=n/a 8.94 [1982WAG/EVA] T= 298.15 K, I=n/a 8.66 [1983LAN] T= 298.15 K,I=n/a 8.41 fl988PHI/HAL1 T= 298.15 K,I=n/a 170 JNC TN8400 99-011 6.6 Lead carbonate system Under environmental conditions carbonate is an important ligand for lead and has a significant influence on lead solubility under neutral to alkaline conditions. The formation constants for lead carbonate complexes and in particular, the solubility products of lead carbonates (cerrusite, hydrocerrusite) reported in the literature are controversial. The experimental data of [1975ERN/ALL, 1976BIL/HUS, 1977SIP/VAL, 1982BIL/SCH, 1987FER/GRE and 1989DOR/MAR] were used to obtain formation constants valid at 1=0. Some of the constants [1973BIL/STU, 1976BIL/HUS] are determined in 0.1 M KNO3 medium. These values were also included in the calculations as the complex formation between lead and carbonate is much stronger than between lead and nitrate and as the constants can be corrected for the interaction between Pb and nitrate using the constants derived in Section 6.7.1. The data measured by Byrne and co-workers [1981BYR, 1988BYR/KUM] in seawater (interaction with Cl) and [1980WAL/SIN] measured in wastewater were not included in the evaluation. The experimentally determined data used to obtain the formation constants valid at 1=0 are given in Table 6.15. Table 6.15: Experimentally determined equilibrium data compiled for the lead carbonate 2+ 2 2m 2n system, according to the equilibrium: mPb + nCO3 -<=> Pbm(CO3)n(OH)o - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: sol = solubility measurements, sp = spectropho tome try, and pot = potentiometry, pol = polarography. Reference Comments KM) Medium Method log (3m,n 2+ 2 log (5 > + CO3 - <=> PbCO3° 1.2 6.37 [1975ERN/ALL] T= 298.15 K, 1=0.1 0.1 KNO3 pot 1, 2 6.47 [1975ERN/ALL] T= 298.15 K, 1=0.1 0.1 KNO3 pot 6.57 1 [1976BEL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 1 6.27 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 5.62 [1976BIL/HUS] T= 298.15 K, 1=0.7 0.7 NaC104 pol 3 5.37 [1977SIP/VAL] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 3 5.34 [1977SIP/VAL] T= 298.15 K, 1=0.7 0.7 NaClO4 pol 5.40 [1982BIL/SCH] T= 298.15 K, 1=0.3 0.3 NaClO4 sol 7.40 [1989DOR/MAR] T= 298.15 K, I=dil 0 no sol 171 JNC TN8400 99-011 Table 6.15: continued 2+ 2 2 log (3]2: Pb + 2CO3 - <=>Pb(CO3)2 - l 9.97 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 9.27 l [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 4 8.2 [1976BIL/HUS] T= 298.15 K, 1=1 1 NaClO4 pot 3 8.37 [1977SIP/VAL] T= 298.15 K, 1=0.7 0.7 NaC104 pol 8.73 [1977SIP/VAL] T= 298.15 K, 1=0.7 0.7 NaC104 pol 8.86 [1982BIL/SCH] T= 298.15 K, 1=0.3 0.3 NaC104 sol 4 8.90 [1987FER/GRE] T= 298.15 K, 1=3 3 NaC104 sol 2+ 2 log Kso: Pb + CO3 - <=> PbCOrfcerrusite) 13.13 6 [1961NAS/MER] T= 298.15 K, 1=0.2-2 0 KNO3 sol 12.71 i [1973BIL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot 12.15 [1982BEL/SCH] T= 298.15 K, 1=0.3 0.3 NaClO4 sol 13.20 [1984TAY/LOP] T= 295 K, I=dil 0 HCO3 sol 13.18 [1989DOR/MAR] T= 298.15 K, I=dil 0 no sol 2+ 2 + log Kso: 3Pb + 2CO3 - + 2H2O t=> Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H 17.64 [1984TAY/LOP] T= 295 K, I=dil 0 HCCy sol 2+ 2 log Kso: 10Pb + 6CO3 - + 7H2O <^PbI0(CO3)6(OH)6O(plumbonacrite) + 8H+ 41.21 [1984TAY/LOP] T= 295 K, I=dil 0 HCO,- sol + corrected in this report for the effect of the formation of PbNO3 and Pb(NO3)2° complexes using the constants derived in Section 6.7.1. Uncorrected values are given in Table 6.16. Pb concentration = 0.0025 mM. no PbHCO3" complexes are formed. Pb concentration = 0.001 mM. recalculation of the experimental data of [1965BAR/BAR] by [1976BIL/HUS] pKw used for calculations in this report (3 M NaClO4) = 14.18 extrapolated to 1=0 by [1961NAS/MER] as a function of I with a poiynom. 172 JNC TN8400 99-011 6.6.1 PbCO3° The formation constant for the PbCC>30 complex at 1=0 was derived from measurements by [1975ERN/ALL, 1976BIL/HUS, 1982BIL/SCH, 1977SIP/VAL, 1989DOR/MAR] (Figure 6.20). The resulting log pV^i of 7.30 is in good agreement with the value of 7.10 and 7.14 used by [1980SCH] and [1984TAY/LOP], respectively. 2+ 2 Pb + CO3 PbCO3° log P°u= 7.30 Pb2+ + CO,2" <=> PbCO, lmi molal 2+ Figure 6.20: Plot of log (3U + 8 D vs. Im for the reaction : Pb +CO3- <=> PbCO3° at 25 °C. The straight line shows the result of the linear regression: Ae = 0.56; log (3°u = 7.30. Calculated from data compiled in Table 6.15. 173 JNC TN8400 99-011 2 6.6.2 Pb(CO3)2 - Experimental values reported in the literature show some spread (see Table 6.15 and Figure 6.21). The largest difference, however, is observed for two values determined by the [1976BIL/HUS] in 0.1 M KNO3 medium with anodic pulse voltammetry and differential pulse polarography. Extrapolation of the results of [1976BIL/HUS, 1977SIP/VAL, 1982BIL/SCH, and 1987FER/GRE] to 1=0 is shown in Figure 6.21 and gives: 2 Pb(CO3)2 - log p°i,2 = 10.13 The resulting log P°i2 of 10.13 is comparable with the value of 10.33 and 10.62 used in the reviews of [1980SCH] and [1984TAY/LOP]. 2+ 2 2 Pb + 2CO3 " ** Pb(CO3)2 " 1 O 12.5 12 3x> 16.13 , _. Q 11.5 00 • '** ^ •"%* 11 - + : : O ;!> - "*:~- } '-, - ••' • v ' O . 10.5 CO. "Illll-—-^-^ 10 : lo g ° ° v"' ' , 9.5 "• • 9 - 8.5 - 8 - I- +- I 1 2 3 lm, molal 2+ 2 2 Figure 6.21: Plot of log pu + 8 D vs. Im for the reaction : Pb + 2CO3 - <=> Pb(CO3)2 " at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.13; log P°i,2 = 10.13. Calculated from data compiled in Table 6.15. 174 JNC TN8400 99-011 4 6 6.6.3 Pb(CO3)3 - and Pb(CO3)4 - 4 6 [1986BROAVAN] calculated formation constants for Pb(CO3)3 - and Pb(CO3)4 \ These formation constants are theoretical values and the existence of these species is not proven. 6.6.4 PbCO3OH- 2+ 2 [1987FER/GRE] determined experimentally a log (5 value for the reaction: Pb + CO3 " + H2O <=> PbCO3OH- + H+ of -3.28 at 1=3 (see Table 6.16). Correction with the SIT term of a single value measured at I = 3 to I = 0 is not possible without making uncertain assumptions. Thus, no log P value for the formation of PbCC^OH" is recommended in this report. 6.6.5 PbHCOf, Pb(HCO3)2° and Pb(HCO3)f Several lead bicarbonate complexes have been reported. For the PbHCO3+ complex, the only actual measurement is that of [1989DOR/MAR] (Table 6.16), who give a log p°u of 2.63 for 2+ the reaction Pb + HCO3- <=> PbHCO3+ or, using a log p value of 10.33 for the reaction CO3- 2+ + H+ <=> HCO3- [1995SIL/BID], a log p°Uil = 12.96 for the reaction Pb + CO3- + H+ <=> PbHCC>3+. As this is only formation constant determined for PbHCC>3+ and as in the experiments of [1989DOR/MAR] less than 10% of the total dissolved Pb are present as PbHCC>3+, the usage of this value is not recommended in this report. [1976BIL/HUS] showed that the results of [1965BAR/BAR and 1967BAR], who proposed the existence of Pb(HCO3)2° and Pb(HCC>3)3~ (given In Table 6.16), can be explained satisfactorily with the formation of a Pb(CC>3)22" complex [1976BIL/HUS], which makes the formation of higher bicarbonate complexes not probable. 175 JNC TN8400 99-011 6.6.6 PbCO3(cerrusite) A number of lead carbonate solids are reported in the literature. From the data of [1961NAS/MER], [1973BIL/STU], [1982BIL/SCH], [1984TAY/LOP], [1989DOR/MAR] (Table 6.15) a log K*°So for cerrusite (PbCO3(cr)) can be derived (Figure 6.22). PbCO3(cerrusite) log K*°so = 13.23 This value can be compared with the log K*°So of 13.10 selected and verified with experimental results by [1995MAR/MAC] (Table 6.16). 2+ 2 Pb +CO3 - PbCO3(cerrusite) 15 2+ Figure 6.22: Plot of log K*so + 8 D vs. Im for the reaction : Pb +CO3" <=> PbCO3(cerrusite) at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.76; log K*°so = 13.23. Calculated from data compiled in Table 6.15. 176 JNC TN8400 99-011 6.6.7 Pb3(CO3)2(OH)2(hydrocerrusite) In the case of hydrocerrusite (Pb3(CO3)2(OH)2(cr)), two distinct groups of log K*So values can be seen in Figure 6.23, Tables 6.15 and 6.16. From the data of [1928RAN/SPE, 1973BIL/STU, 1982BIL/SCH] (Table 6.16) one extrapolates a log K*So of 19.31 for hydrocerrusite at 1=0. However, more recent measurements by [1983SCH/GAR] (Table 6.16) and [1980PAT] indicate that this log K*so for hydrocerrusite is too large, i.e., underestimating lead solubility. [1984TAY/LOP] determined with X-ray analysis (see also comments in Section 6.13) the following solubility product at 1=0 for hydrocerrusite: 2+ 2 3Pb + 2CO3 - + 2H2O <=> Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+ log K*°so = 17.64 This value is in accordance with the values proposed by [1983SCH/GAR] and [1995MAR/MAC] (cf. Table 6.16). The difference between the two groups of data can not be explained satisfactorily at the moment. We propose to use a log K*°so of 17.64 for hydrocerrusite. 2+ 2 3Pb + 2CO3 " + 2H2O Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+ 21 0.2 0.4 0.6 0.8 lm, molal 2+ 2 Figure 6.23: Plot of log K*So + 18 D vs. Im for the reaction: 3Pb + 2CO3 " + 2H2O <=> Pb3(CO3)2Pb(OH)2(hydrocerrusite) + 2H+ at 25 °C. log K*°so = 17.64. Calculated from data compiled in Table 6.15. 177 JNC TN8400 99-011 6.6.8 Plwnbonacrite [1984T AY/LOP] showed with X-ray analysis the existence of plumbonacrite (Pbio(C03)6(OH)60(cr)) under certain conditions in alkaline medium (see also comments in Section 6.13). They determined a log K*°So of 41.21 at 1=0 (Table 6.15), which differs strongly from the value given by [1981HAA/WIL] (Table 6.16), who concluded that plumbonacrite is metastable with respect to hydrocerrusite and litharge. [1981HAA/WTL] did not report any experimental details or details of the calculations, while the value given by [1984TAY/LOP] is based on extensive observation of interconversion reactions of plumbonacrite. In this report the value determined by [1984TAY/LOP] is proposed: 2+ 2 10Pb + 6CO3 - + 7H2O 4=> Pb10(CO3)6(OH)6O(plumbonacrite) + 8H+ log K*°s0 = 41.21 6.6.9 PbCO3PbO(s) and PbCO3(PbO)2(s) Based on data given in different thermodynamic compilations (Table 6.17), log K*so values for PbCO3PbO(s) and PbCO3(PbO)2(s) can be calculated Primary experimental data, however, are not available and the conditions under which these minerals are formed do not seem to be known. 6.6.10 Phosgen ite A log K*so for phosgenite (Pb2CO3Cl2(cr)) of approximately 19.9 can be calculated from the data compiled in previous compilations (Table 6.17). Primary experimental data are not available and the conditions which these minerals are formed do not seem to be known. 178 JNC TN8400 99-011 6.6.11 Additional data for the lead carbonate system Table 6.16: Additional experimentally determined data for the lead carbonate system, according to the 2+ 2 2m 2n + equilibrium: mPb + nCO3 " + oH2O <=> Pbm(CO3)n(OH)o - -° + H . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text in Section 6.6 and in Section 6.13: 'Comments on selected references'. Method: pol = polarography, pot = potentiometry, sol = solubility measurements, and sp = spectrophotometry. Reference Comments Medium Method 2+ 2 log pu: Pb + CO3 - & PbCO3° 7 ' [1969BAR] T= 298.15 K, 1=1 1 NaC104 sol 2 6.4 [1973BBL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot 3 4 6.2 - [1975ERN/ALL] T= 298.15 K, 1=0.1 0.1 KNO3 pot 6.3 3- 4 [1975ERN/ALL] T= 298.15 K, 1=0.1 0.1 KNO3 pot 6.40 3 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 6.10 3 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 6.34 [1980WAL/SIN] T = 298.15 K,I=dil 0 H2SO4 sol 7.10 [1980SCH] T = 298.15 K, 1=0 0 n/a sol 4.00 5 [1981BYR] T= 298.15 K, 1=0.7 0.7 SW sp 6.00 6 [1987FER/GRE] T= 298.15 K, 1=3 3 NaCIO, sol 2+ 2 2- log J5U: Pb + 2CO3 - <=> Pb(CO3)2 8.2 7 [1959FAU/BON] T=291 K, 1=1.7 1.7 KNO3 pol 9 1 [1969BAR] T= 298.15 K, 1=1 1 NaC104 sol 9.8 2 [1973BEL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot 9.8 3 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNOj pot 9.1 3 [1976BIL/HUS] T= 298.15 K, 1=0.1 0.1 KNO3 pot 10.8 3 [1976BIL/HUS] T= 298.15 K, 1=0 0 KNO3 pot 10.33 [1980SCH] T = 298.15 K, 1=0 0 n/a sol 8 10.40 [1987FER/GRE1 T= 298.15 K, 1=0 0 NaClO4 sol 2+ 2 + ,,u: Pb + CO3 - + H2O <=> Pb(CO3)OH-+ H 9 -3.28 [1987FER/GRE1 T= 298.15 K, 1=3 NaC104 sol 2+ + log Pu: Pb + HCO3<=> PbHCO3 10 2.63 [1989DOR/MAR] T= 298.15 K,I=dil no sol 2+ log pl2: Pb + 2HCO3<=> Pb(HCO3)2° 4.77 '' [1965BAR/BAR] T= 293 K, 1=1 NaHCOj sol 4.78 "• 12. ri967BAR] T= 298 K, 1=1 NaHCO, sol 179 JNC TN8400 99 -Oil Table 6.16: continued : log Pu 3HCO3<=> Pb(HCO3)3 u 5.19 [1965BAR/BAR] T=293K,I=1 NaHCO3 sol 5.20 "• 12 [1967BAR1 T= 298 K, 1=1 NaHCO, sol Table 6.16: continued 2 f log Kso: Pb + CO/" <=> PbCO3(cerrusite) 12.54 ' [1973BIL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot l3 12.60 [1980WAL/SIN] T= 298 K, I=dil 0 H2SO4 sol 13.10 fl995MAR/MACl T = 298.15 K, 1=0 0 no sol 2+ 2 + log Kso: 31>b + 2CO3 - + 2H2O «• Pb3(CO3)2Pb(OH)2(hydrocerrusite) -\-2H 19.04 !4 [1928RAN/SPE] T= 298.15 K, I=n/a 0 NaOH sol 3 17.11 [1973BEL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot 17.63 15 [1973BIL/STU] T= 298.15 K, 1=0.1 0.1 KNO3 pot 16.56 [1982BIL/SCH] T= 298.15 K, 1=0.3 0.3 NaC104 sol 16.25 [1982BBL/SCH] T= 298.15 K, 1=0.3 0.3 NaC104 sol 16.88 15 [1982BIL/SCH] T= 298.15 K, 1=0.1 0.1 KNO3 sol 17.40 9 [1982BIL/SCH] T= 298.15 K, 1=0.1 0.1 KNO3 sol 17.80 [1983SCH/GAR] T = 298.15 K, 1=0.0075 0.0075 CaCO3 sol 17.50 [1995MAR/MAC] T = 298.15 K, 1=0 0 no sol 2+ 2 log Kso: 10Pb + 6CO3 - + 7H2O <=> Pbl0(CO3)6(OH)6O(plumbonacrite) + 8W 8.76 16 [1981HAA/WIL1 T= 298.15 K,I=n/a n/a sol approximate values, calculated by [1969BAR] with a log Kso*° of 12.83 for cerrusite. same values also reported in [1976BIL/HUS] values corrected for the effect of the formation of Pb-nitrate complexes are given in Table 6.15. Pb concentration = 0.0025 mM SW = seawater. Formation of lead chloride complexes probable 2+ 2 log P,, for the reaction Pb + CO3 " <=> PbCO3° is only an approximate value. pK^ in 3 M NaClO4 = 14.18. measured in 1.7 M nitrate medium, formation of Pb nitrate complexes probable. Pb concentration = 0.1 mM. extrapolated to 1=0 with SIT term by [1987FER/GRE]. pKw used for calculations in this report (3 M NaClO4) = 14.18 - only value reported in the literature, relatively small amount of PbHCO3 present in the experiments of [1989DOR/MAR] recalculated by [1976BIL/HUS] assuming the formation of Pb carbonate complexes . given in [1969BAR]. determined in wastewater, see also comments in Section 6.13. 14 calculated with a log p!3 value of -28.02 (1=0; Section 6.1.3) 15 + corrected in this report for the effect of the formation of PbNO3 and Pb(NO3)2° complexes using the constants derived in Section 6.7.1. Uncorrected values are given in Table 6.16. 16 extrapolated to 1=0, no experimental details reported. 180 JNC TN8400 99-011 Table 6.17: Thermodynamic data for the lead carbonate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log Pm,n Reference Comments I(M) Medium 2+ 2 log PlA: Pb + CO3 - <=> PbCO3° 7.5 [1972ZIR/YAM] T= 298.15 K, 1=0 0 7.24 [1976HEM] T= 298.15 K,I=n/a 7.50 [1980MAN/DEU] T= 298.15 K, 1=0 0 7.00 [1981TURAVHI] T= 298.15 K, 1=0 0 6.38 [1983LAN] T= 298.15 K, 1=0 0 7.14 [1984TAY/LOP] T= 295 K, I=dil 0 6.81 [1987BROAVAN] T= 298.15 K, 1=0 0 6.13 [1987BRU] T= 298.15 K, 1=0- 1 0 5.07 [1988BYR/KUM] T= 298 K, 1=0.7 0.7 SW 5.10 [1988BYR/KUM] T= 298 K, 1=0.7 0.7 SW 6.53 [1988PHI/HAL] T= 298.15 K, 1=0 0 5.40 [1989SMI/MAR] T= 298 K, 1=0.3 0.3 2+ 2 log Pu: Pb + CO/- <=> Pb(CO3)2 10.64 [1976HEM] T= 298.15 K,I=n/a 0 8.20 [1976SMI/MAR] T= 291.15 K, 1=1.7 1.7 8.62 [1980MAN/DEU] T=298.15K,I=0 0 10.29 [1981TUR/WHI] T= 298.15 K, 1=0 0 9.66 [1983LAN] T= 298.15 K, 1=0 0 10.62 [1984TAY/LOP] T= 295 K, I=dil 0 12.29 [1987BRO/WAN] T= 298.15 K, 1=0 0 9.02 [1987BRU] T= 298.15 K, 1=0-1 0 7.26 [1988BYR/KUM] T= 298 K, 1=0.7 0.7 SW 10.08 [1988PHI/HAL] T= 298.15 K, 1=0 0 8.86 [1989SMI/MAR] T= 298 K, 1=0.3 03 2+ 2 log pu: Pb + 3CO3 - <=> Pb(CO3)/- 16.70 IT987BRO/WAN1 T= 298.15 K, 1=0 2+ 2 6 log p, „: Pb + 4CO3 - <=> Pb(CO3)4 - 20.14 [1987BROAVAN1 T= 298.15 K, 1=0 2+ + log ft,,: Pb + HCOf<=> PbHCO3 2.9 [1972ZIR/YAM] T= 298.15 K, 1=0 0 2.30 [1983SCH/GAR] T = 298.15 K, I=n/a 2.78 [1987BRO/WAN] T= 298.15 K, 1=0 0 5.48 [1988PHLHAL1 T= 298.15 K, 1=0 0 181 JNC TN8400 99-011 Table 6.17: continued 2+ log Pu,: Pb + 2HCO3<=> Pb(HCO3)2° 4.43 [1987BROAVAN] T= 298. 15 K, 1=0 0 8.76 ri988PHI/HALl T= 298. 15 K,1=0 0 2+ log pu: Pb + 3HCOf<=> Pb(HCO3)f 5.20 [1987BRQ/WAN1 T= 298.15 K, 1=0 2+ 2 log P1A: Pb + 4HCOf<=> Pb(HCO3)4 - T= 298.15 K, 1=0 2+ log Kso: Pb + CO/- <=> PbCO3(cerrusite) 12.82 [1952LAT] T= 298.15 K,I=n/a 13.24 [1959UGG] T= 298.15 K,I=n/a 13.46 [1971NAU/RYZ] T= 298.15 K, I=n/a 12.83 [1973BAR/KNA] T= 298.15 K,I=n/a 13.13 [1976HEM] T= 298.15 K, 1=0 0 13.13 [1976SMI/MAR] T= 298.15 K, 1=0 0 11.01 [1976SMI/MAR] T= 298.15 K, 1=1 1 12.97 [1977PAU] T= 298.15 K,I=n/a 12.80 [1978ROB/HEM2 T= 298.15 K,I=n/a 12.91 [1979KUB/ALC] T= 298.15 K,I=n/a 12.85 [1979PAT/OBR] T = 298.15 K, 1=0 0 13.45 [1980BEN/TEA] T= 298.15 K,I=n/a 13.13 [1980CLE/JOH] T=298.15 K, I=n/a 13.13 [1980MAN/DEU] T=298.15K,I=0 0 13.11 [1980SCH] T = 298.15 K, 1=0, 0 12.83 [1981 STUMOR] T= 298.15 K,I=n/a 12.97 [1982PAU] T= 298.15 K,I=n/a 12.82 [1982WAG/EVA] T= 298.15 K,I=n/a 12.78 [1983LAN] T= 298.15 K,I=n/a 12.81 [1983SAN/BAR] T= 298.15 K,I=n/a 12.80 [1984NRI2] T= 298.15 K,I=n/a 10.91 ' [1984SVE] T = 298.15 K, 1=0 0 12.95 [1985BABMAT] T= 298.15 K, I=n/a 12.89 [1985GAL] T= 298.15 K,I=n/a 12.82 [1985MUL] T= 298.15 K,I=n/a 13.13 [1986FLE/JOH] T= 298.15 K,I=n/a 12.15 [1986FLE/JOH] T= 298.15 K,I=n/a 13.64 [1987BROAVAN] T= 298.15 K,I=n/a 12.86 [1988PHI/HAL] T= 298.15 K,I=n/a 13.13 [1989SMI/MAR] T= 298.15 K, 1=0 0 12.00 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 12.15 [1989SMI/MAR] T= 298.15 K, 1=0.3 0.3 11.01 [1989SMIMAR1 T= 298.15 K, 1=1 1 182 JNC TN8400 99-011 Table 6.17: continued 2+ 2 log Kso: 3Pb + 2CO3 - + 2H2O <=> Pb3(C03)2Pb(OH)2(hydrocerrusite) I-2//- + 19.00 [1952LAT] T= 298.15 K,I=n/a 16.77 [1965GAR7CHR] T= 298.15 K, I=n/a 18.93 [1971NAU/RYZ] T= 298.15 K, I=n/a 19.04 2 [1976BAE/MES] T= 298.15 K, 1=0 0 19.04 2 [1976SMI/MAR] T= 298.15 K, 1=0 0 16.72 [1977PAU] T= 298.15 K,I=n/a 18.8 [1979PAT/OBR] T= 298.15 K, 1=0 0 17.46 [1980MAN/DEU] T = 298.15 K, 1=0 0 18.78 [198OSCH] T = 298.15 K, 1=0 0 16.72 [1982PAU] T= 298.15 K,I=n/a 19.13 [1983LAN] T= 298.15 K,I=n/a 16.97 [1983SAN/BAR] -. T= 298.15 K,I=n/a 19.00 [1985GAL] T= 298.15 K,I=n/a 16.08 [1986FLE/JOH] T= 298.15 K,I=n/a 19.05 [1988PHI/HAL] T= 298.15 K,I=n/a 16.24 [1989SMI/MAR1 T= 298.15 K, 1=0.3 0.3 Table 6.17: continued 2+ 2 + log Kso: 10Pb + 6CO3 - + 7H2O <=> Pbw(CO3)6(0H)60(plumbonacrite) + 8H 8.76 [1989SMI/MAR1 T= 298.15 K, 1=0 0_ 2+ 2 log Kso: 2Pb + CO3 - + H2O t* PbCO3PbO(s) 0.81 [1952LAT] T= 298.15 K,I=n/a -17.28 [1977BAR/KNA] T= 298.15 K,I=n/a 0.81 [1977PAU] T= 298.15 K, I=n/a -17.28 [1979KUB/ALC] T= 298.15 K,I=n/a 0.48 [1980BEN/TEA] T= 298.15 K,I=n/a 0.81 [1982PAU] T= 298.15 K,l=n/a 0.49 [1982WAG/EVA] T= 298.15 K, I=n/a 0.78 [1985BAB/MAT] T= 298.15 K, I=n/a 0.81 [1985GAL] T= 298.15 K,I=n/a 2+ 2 + log Kso: 3Pb + CO3 ' + 2H2O <=> PbCO3(PbO)2(s) + 4H -10.91 [1952LAT] T= 298.15 K, I=n/a -10.91 [1977PAU] T= 298.15 K,I=n/a -10.99 [1982PAU] T= 298.15 K,I=n/a -11.04 [1985BAB/MAT] T=298.15K, I=n/a -11.09 [1985GAL1 T= 298.15 K,I=n/a 2+ 2 log Kso: 2Pb + CO3 - + 2CI- c=> Pb2CO3Cl2(phosgenite) 19.87 [1971NAU/RYZ] T= 298.15 K,I=n/a 19.80 [1980BEN/TEA] T= 298.15 K, I=n/a 19.80 [1980MAN/DEU] T=298.15K, 1=0 19.80 [1982WAG/EVA] T= 298.15 K,I=n/a 19.95 [1988PHI/HAL] T= 298.15 K, I=n/a 1 calculated with linear free energy relation from the Afi° value of Pb2+ by [1984SVE] 2 2+ + calculated with a log Pu of - 28.02 (Pb + 3 H2O <=> Pb(OH)3- + 3H ) 183 JNC TN8400 99-011 6. 7 Lead nitrate system Lead forms weak complexes with nitrate. Several measurements of the formation constants can be found in the literature. Experimental data for the formation of lead nitrate complexes used in this report for extrapolation to I = 0 are given in Table 6.18. Further experimental results reported in the literature are shown in Table 6.19 and log (3 values proposed by authors of previous reviews are collected in Table 6.20. Table 6.18: Experimentally determined equilibrium data compiled for the lead nitrate system, 2+ 2m n according to the equilibrium: mPb + nNO3" <=> Pbm(NO3)n - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: kin = kinetic measurements, pol = polarography, pot = potentiometry, sol = solubility measurements, and sp = spectrophotometry. Reference Comments KM) Medium Method log Pm.n 2+ + log pu: Pb + NOf d PbNO3 0.45 [1953HER/SMI] T= 298.15 K, 1=2 2 NaClO4 pol 0.52 [1953HER/SMI] T= 298.15 K, 1=2 2 NaC104 pot 0.20 [1953HER/SMI] T= 298.15 K, 1=2 2 NaC104 sp 2 0.25 1. [1955BIG/PAR] T= 298.15 K, 1=0.5 0.5 HC1O4 sp 0.31 1.2 [1955BIG/PAR] T= 298.15 K, 1=1 1 HC1O4 sp 0.36 1.2 [1955BIG/PAR] T= 298.15 K, 1=2 2 HC1O4 sp 3 0.82 [1956BAL/DAV] T= 298.15 K, 1=0.051 0.051 NaC104 sp 3 0.77 [1956BAL/DAV] T= 298.15 K, 1=0.066 0.066 NaClO4 sp 0.70 3 [1956BAL/DAV] T= 298.15 K, 1=0.079 0.079 NaClO4 sp 3 O.7O [1956BAL/DAV] T= 298.15 K, 1=0.092 0.092 NaC104 sp 3 0.33 [1956BAL/DAV] T= 298.15 K, 1=0.624 0.624 NaC104 sp 0.23 3 [1956BAL/DAV] T= 298.15 K, 1=0.943 0.943 NaClO4 sp 3 0.23 [1956BAL/DAV] T= 298.15 K, 1=1.168 1.168 NaClO4 sp 3 0.19 [1956BAL/DAV] T= 298.15 K, 1=1.455 1.455 NaClO4 sp 4 0.15 [1965HUG2] T= 298.15 K, 1=2 2 NaClO4 pot 4 0.26 [1965HUG3] T= 298.15 K, 1=2 2 NaClO4 pol 0.53 [1972FED/ROB] T= 298.15 K, 1=0.5 0.5 LiClO4 pol 0.36 [1972FED/ROB] T= 298.15 K, 1=0.75 0.75 LiClO4 pol 0.43 [1972FED/ROB] T= 298.15 K, 1=1 1 LiClO4 pol 0.40 [1972FED/ROB] T= 298.15 K, 1=2 2 LiClO4 pol 0.45 [1972FED/ROB] T= 298.15 K, 1=3 3 LiClO4 pol 0.52 [1972FED/ROB] T= 298.15 K, 1=3 3 LiClO4 pol 184 JNC TN8400 99-011 Table 6.18: continued 5 0.53 [1972FED/ROB] T= 298.15 K, 1=3 3 LiClO4 pot 5 0.57 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 sp 5 0.46 [1972FED/ROB] T= 298.15 K,I=3 3 LiC104 sol 0.62 [1973HUT/HIG] T= 298.15 K, 1=1 1 NaClO4 kin 2+ log P12: Pb + 2NO3-<^> Pb(NO3)2° 4 0.39 [1965HUG2] T= 298.15 K, 1=2 2 NaC104 pot 4 0.38 [1965HUG3] T= 298.15 K, 1=2 2 NaC104 pol 0.43 [1972FED/ROB] T= 298.15 K, 1=0.5 0.5 LiC104 pol 0.48 [1972FED/ROB] T= 298.15 K, 1=0.75 0.75 LiClO4 pol 0.43 [1972FED/ROB] T= 298.15 K, 1=1 1 LiC104 pol 0.23 [1972FED/ROB] T= 298.15 K, 1=2 2 LiC104 pol 0.41 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 pol 0.45 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 pol 5 0.48 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 pot 5 0.48 [1972FED/ROB] T= 298.15 K, 1=3 3 LiClO4 sp 5 0.30 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 sol 2+ log p]3: Pb + 3NOf <=> Pb(NO3)f -0.05 [1972FED/ROB] T= 298.15 K, 1=2 2 LiClO4 pol 0.08 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 pol 0.26 [1972FED/ROB] T= 298.15 K, 1=3 3 LiClO4 pol 5 0.30 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 pot 5 0.18 [1972FED/ROB] T= 298.15 K, 1=3 3 LiC104 sol 2+ log K*°so: Pb + NO3 + H2O <=> PbOHNO3(s) + 6 -3.55 [1945PED] T= 291 K, 1=0.3 0.3 Ba(NO,)7 sol 50-600 mM Pb corrected to 1=0 with the Davies equation by the respective authors original values corrected with Daviesby [1956BAL/DAV]; corrected back in this report 0.1-100 mMPb put together by [1972FED/ROB] from earlier measurements of Fedorov and Coworkers. 5-400 mM Pb 185 JNC TN8400 99-011 6.7.1 Lead nitrate complexes Extrapolation of the measurements compiled in Table 6.18 is shown in Figures 6.24, 6.25 and 6.26: PbNO3+ log P°u = 1.06 Pb(NO3)2° log P°i.2= 1-48 3NO3 Pb(NO3)3_0 log P°i,3 = 0.76 2+ 2 For the reaction Pb + 4NO3" O Pb(NO3)4 " no log §°\£ is proposed in this report as the correction with the SIT term of values measured at I = 3 to I = 0 is not possible without making 2 uncertain assumptions. The complex Pb(NO3)4 " is important only in very concentrated nitrate 2 solutions, which makes the extrapolation to 1=0 difficult and the occurrence of Pb(NO3)4 " not very likely under environmental conditions. Pb2+ + NO PbNCV 2+ + Figure 6.24: Plot of log (3U + 4 D vs. Im for the reaction : Pb + NO3- PbNO3 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.09; log 1.06. Calculated from data compiled in Table 6.18. 186 JNC TN8400 99-011 2+ Pb + 2NO3" <:* Pb(NO3)2° ; - • • .-.:.•.• .-:•„'«.:;y., {".--•!•;••<•;:•.-.; V •'•—..i.-.v .:•:•.:!•:>!.:•• . :::•:: : [> [ •. .•; . • . : "::••.•-;'•::;: ; \~:..^,: '^r-i'~ --;••;• -'':'::'.i:"t:- ::\' •'-. 4.5 - '.' -.'.'•'".'''•1"!^':;.;:'[h1'::'!::!^ii;i:i!:i;:!i*iii;ii;Diii-!':-"^'-;--:r:"^;i^^i:'i ;; 1 : : : : : : "••••. ::."-.•; /••.•:";• .I'- ':"!;'; -' !=": .:;;i"i -:-;*-v' ••"^•• " •"••:-.•:: i.-iv 4- ••'• ..-•:••;•;.: ' •:•-•"' "-: :-'•. -:^;^^ =-:j:":v::it ?:.•;•?•«::<:j••!-.<;.'/• :•": ^; Q 3.5- •?: -f S'L? S? ?:•;";: 3-.:?R;'«S;'::.TP:H| : : : 1 + 3 . ::; - •'••• • "•.. •::•:.]•::• ^ " -"'" " 'i '^^^;: V:.•:.;•:':•:.::'..::; CM 1 : A | :i : | ^2.5 v'T!~^. ": :i''\/;;'L Sf;;--:;i::!r«^i*i:r:A : i " :t!;i;S- : : 1 1 1 .: ' . •••• ••:•;;:•!;•!«;=: irvi",--.:-:--' . ;-.:/;-'':- !--:? l fl'-:Si;V;y:--:.t': ::i:-;fi:r7.;:.•:•,;:• g) 2 :«*.;; :..:.».:A::::-.':-:V;:;:^:::;;,;P.r 1.5 • : : : 1 l -^- 1.p1 S:;i;S'll:B|:;;l;:s;!w 0.5 0 - 1 2 lm, molal 2+ Figure 6.25: Plot of log p1>2 + 6 D vs. Im for the reaction : Pb + 2NO3- o Pb(NO3)2° at 25 °C. The straight line shows the result of the linear regression: Ae = -0.09; log P°i,2 = 1-48. Calculated from data compiled in Table 6.18. 2+ Pb + 3NO3- <=> Pb(NO3)3- 2+ Figure 6.26: Plot of log p]>3 + 6 D vs. Im for the reaction : Pb + 3NO3" <=> Pb(NO3)3" at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.22; log P°j 3 = 0.76. Calculated from data compiled in Table 6.18. 187 JNC TN8400 99-011 6. 7.2 Pb(NO3)2(s) and PbOHNO3(cr) The high solubility of Pb(NO3)2(s) in water is well established and is given in the review of [1980CLE/JOH] as 1.80 mol/kg H2O (in agreement with e.g. [1935AKEATUR] and [1960KAZ]). Assuming the predominance of Pb(NC>3)20 and 1=0, this value can be corrected for the formation of Pb nitrate complexes, which gives a tentative log K*so of -0.28. In solution with pH > 3, however, [1945PED] has shown the precipitation of PbOHNO3(cr) with X-ray diffraction. The only experimentally determined value is by [1945PED] in 0.3 M nitrate medium. Extrapolation to 1=0 with the SIT model (assuming Ae ~ 0) results in: + NO3-+H2O PbOHNO3(cr) + H log K*°so= - 2.94 6. 7.3 Additional data compiled for the lead nitrate system Table 6.19: Additional experimentally determined data for the lead nitrate system, according to the 2+ 2m n equilibrium: mPb + nNO3" <=> Pbm(NO3)n ' . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 6.13: 'Comments on selected references'. Method: con = conductivity measurements, pol = polarography, pot - potentiometry, sol = solubility measurements, and sp = spectrophotometry. log Pm,n Reference Comments KM) Medium Method 2+ + log pu: Pb' + NO3 <=> PbNO3 1.19 l [1930RIG/DAV] T= 298.15 K, 1=0.001-0,02 0 NaCl con 2 0.52 [1953HER/SMI] T= 298.15 K, 1=6 6 NaC104 sp 3 1.19 [1955NAN] T= 298.15 K, 1=0.001 0 NaC104 con 3 1.15 [1956BAL/DAV] T= 298.15 K, 1=0.9-1.5 0 NaClO4 sp 4 1.11 F1972FED/ROB1 T= 298.15 K, 1=0 0 LiC104 pol 2+ log j5u: Pb + 2N0f <=> Pb(NO3)2° 4 1.40 [1972FED/ROB1 T= 298.15 K, 1=0 LiClO4 pol 2+ log pu: Pb + 3NOf <=> Pb(NO3)f -2.30 5 [1965HUG31 T= 298.15 K, 1=2 NaCIO, pol 2+ 2 log j3l4: Pb + 4NOf <=> Pb(NO3)4 - -0.30 [1972FED/ROB] T= 298.15 K, 1=3 LiClO4 pol 6 0.11 [1972FED/ROB] T= 298.15 K, 1=3 LiClO4 pot -0.52 6 [1972FED/ROB] T= 298.15 K, 1=3 LiClO. sol JNC TN8400 99-011 Table 6.19: continued 2+ log K*°so: Pb + 2NO3 <=> Pb(NO3)2(s) -0.28 7 [1935AKE/TUR] T=298.15 K, I=dil 0 water sol -0.28 7 [1960KAZ] T=298.15 K, I=dil 0 water sol^ 1 corrected to 1=0 with a simplified Debye-Hiickel equation 2 not used because of high ionic strength 3 corrected to 1=0 with Davies equation. 4 corrected to 1=0 with Debye-Huckel equation 5 tentative value 6 put together by [1972FED/ROB] from earlier measurements of Fedorov and Coworkers. 7 after correction for the formation of Pb nitrate complexes using the constants calculated in Section 6.7.1. Table 6.20: Thermodynamic data for the lead nitrate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log Pm n Reference Comments I (M) 2 log hi: Pb * + NOf <=> PbNO3* 1.17 [1976SMI/MAR] y 298.15 K, 1=0 0 0.25 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.33 [1976SMI/MAR] T= 298.15 K, 1=1 1 0.40 [1976SMI/MAR] T= 298.15 K, 1=2 2 0.51 [1976SMI/MAR] T= 298.15 K,I=3 3 1.18 [1980BEN/TEA] T= 298.15 K, 1=0 0 1.15 [1982WAG/EVA] T= 298.15 K, 1=0 0 1.20 [1987BROAVAN] T= 298.15 K, 1=0 0 2+ log (lu: Pb + 2NO3- <=> Pb(NO3)2° 1.40 [1976SMI/MAR] T= 298.15 K,I=0 0 0.40 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.40 [1976SMI/MAR] T= 298.15 K,I=1 1 0.40 [1976SMI/MAR] T= 298.15 K, 1=2 2 0.40 [1976SMI/MAR] T= 298.15 K, 1=3 3 0.83 [1987BROAVAN1 T= 298.15 K, 1=0 0 2 log pu: Pb * + 3NOf <=> Pb(NO3)f 0.10 [1976SMI/MAR] T= 298.15 K, 1=2 2 0.20 [1976SMI/MAR] T= 298.15 K, 1=3 3 -0.86 ri987BRO/WAN] T= 298.15 K, 1=0 0_ 2 2 log $lA: Pb * + 4NOf <=> Pb(NO3)4 - -0.30 [1976SMI/MAR] T= 298.15 K, 1=3 3 -3.76 [1987BROAVAN] T= 298.15 K, 1=0 0 189 JNC TN8400 99-011 Table 6.20: continued to 2 log K so: Pb * + 2NO/ <=> Pb(NO3)2(s) 1.09 [1952LAT] T= 298.15 K, I=n/a 1.09 [1985GAL] T= 298.15 K, I=n/a -0.28 ' [1980CLE/JOH1 T= 298.15 K, I=n/a 2+ + log K*°so: Pb + NO/ + H2O <=> PbOHNO3(s) + H -12.01 [1952LAT] T= 298.15 K, I=n/a -12.03 fl985GAL1 T= 298.15 K, I=n/a I after correction for the formation of Pb nitrate complexes using the constants calculated in Section 6.7. 190 JNC TN8400 99-011 6.8 Lead phosphate system Lead forms complexes with phosphate. Several measurements of the formation constants can be found in the literature. Experimental data for the formation of lead phosphate complexes and solids used in this report are given in Table 6.21. Further experimental results reported in the literature are shown in Table 6.22 and thermodynamic values proposed by authors of previous reviews are compiled in Table 6.23. Table 6.21: Experimentally determined equilibrium data compiled for the lead phosphate 2+ 3 2m 3n system, according to the equilibrium: mPb + nPO4 " <=> Pbm(PO4)n - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: sol = solubility measurements. Reference Comments I (M) Medium Method Pm,n 2+ 3 + log plu: Pb + PO4 - + H <=>PbHPO4° 1 15.45 [1972NRI] T= 298.15 K, I=dil 0 H,PO4 sol 2+ 3 + log fiI21: Pb + PO4 ' + 2H ] 21.05 [1972NRI] T= 298.15-K, I=dil 0 H,PO4 sol 2+ 3 log K*so-Pb + PO4 - + H+ <=>PbHPO4(secondary lead phosphate) 23.78 1 [1972NRI] T= 298.15 K, 1=0.1 0^ KOH sol 2+ 3 log K*so: 3Pb + 2PO4 ' <=> Pb3(PO4)2(tertiary lead phosphate) 44.40 ! [ 1972NRI] T= 298.15 K, 1=0.1 0 KOH sol 2+ 3 log K*so: Pb + 2PO4 - + 4H+ <=> Pb(H2PO4)2(primary lead orthophosphate) 48.94 J [1973NRI2] T= 298.15 K, 1=0.1 () NaCl sol 2+ 3 + log K*S0:4Pb + 2PO4 - + H2O <^ Pb4(PO4)2O(tetraplumbite orthophosphate) + 2H 37.09 ' [1972NRI] T= 298.15 K, 1=0.1 0 KOH sol 191 JNC TN8400 99-011 Table 6.21: continued 2+ 3 + log K*S0:5Pb + 3PO4 - + H2O <=> Pb5(PO4hOH(hydroxypyromorphite)+ H 62.80 * [1972NRI] T= 298.15 K, 1=0.10KOHsol 2+ 3 log K*so: 5Pb + 3PO4 - + Cl- <=> Pb5(PO4)3Cl(s) (chloropyromorphite)+ H+ 84.4 ! [1973NRI2] T= 298.15 K, 1=0.1 0 NaCl sol 2+ 3 log K*S0:5Pb + 3PO4 - + F- <=>Pb5(PO4)3F(s) (fluoropyromorphite)+ H+ 71.60 * [1973NRI1] T= 298.15 K, 1=0.1 0 NaF sol 1 extrapolated to 1=0 with Davies equation by (1972NRI], [1973NRI1] and [1973NRI2]. + 6.8.1 PbH2PO4 and PbHPO4° Nriagu [1972NRI] made measurements in dilute phosphoric acid solution and determined the constants for the formation of the aqueous species PbH2PO4+ and PbHPO4°. These values are the only experimentally determined values available and a correction to 1=0 with the Davies equation has been made by [1972NRI]. 2+ 3 Pb + PO4 " + H+ <=> PbHPO4° log p\u = 15.45 2+ 3 Pb + PO4 - + 2H+ <=» PbH2PO4+ log p°i'2,i = 21.05 6.8.2 PbHPO4(cr) and Pb3(PO4)2(s) The log K*so for PbHPO4(secondary lead phosphate) and for Pb3(PO4)2(tertiary lead phosphate) given in [1972NRI] (Table 6.20) agree very well with the values determined at 310 Kby [1932JOW/PRI] and by [1969AWA/ELH] in 1 M phosphate solution (compiled in Table 6.21). As only Nriagu determined log K*so values for several different lead phosphates, it is recommended to use the consistent dataset given by Nriagu: 2+ 3 Pb + PO4 - + H+ <=> PbHPO4(secondary lead phosphate) log K*°So = 23.78 2+ 3 3Pb + 2PO4 - 4=> Pb3(PO4)2(tertiary lead phosphate) log K*°So = 44.40 [1972NRI] states that secondary and tertiary lead phosphate, however, are not stable in the presence of water at ambient temperatures but will transform to hydroxypyromorphite or other lead phosphates. 192 JNC TN8400 99-011 6.8.3 Pb(H2PO4)2(s), Pb4(PO4)2O(s) and plumbogummite(PbAl3(PO4)2(OH)5(cr)) The following values have been determined by [1972NRI] and [1973NRI2]: 2+ 3 4Pb + 2PO4 -+ H2O & Pb4(PO4)2O(s) + 2H+ log K*°so = 37.09 tetraplumbite orthophosphate 2+ 3 Pb + 2PO4 " + 4H+ <=> Pb(H2PO4)2(s) log K*°so = 48.94 primary lead orthophosphate Again it is recommended to use the internally consistent dataset given by Nriagu. The log K*°so of 99.3 given by [198NRI] (Table 6.23) for plumbogummite (PbAl3(PO4)2(OH)5(cr)), however, is only an estimate. 6.8.4 Pyromorphites: Pb5(PO4)3Cl(s), Pb5(PO4)3F(s), and Pb5(PO4)3OH(s) Again, not many authors have determined log K*so values for pyromorphites and it is recommended to use the internally consistent dataset given by Nriagu (Table 6.21): 2+ 3 5Pb + 3PO4 -+ H2O <=> Pb5(PO4)3OH(cr) + H+ log K*°so = 62.8 hydroxypyromorphite 2+ 3 5Pb + 3PO4 -+ Cl- <=> Pb5(PO4)3Cl(cr) log K*°s0 = 84.4 chloropyromorphite 2+ 3 5Pb + 3PO4 " + F- « Pb5(PO4)3F(cr) log K*°so = 71.6 fluoropyromorphite The log K*so of 84.4 for chloropyromorphite (Pb5(PO4)3Cl(cr)) determined in [1973NRI2] after an aging time of 48 h agrees well with the value determined at 310 °K by [1932JOW/PRI] after a few hours (Table 6.22). [1964BAK] determined after an equilibration time of several months a much lower log K*so of 34.5 for chloropyromorphite (Table 6.22). An examination of his data leads to the conclusion that [1964BAK] assumed all phosphate to be present as 3 PO4 " (in a Na2HPO4 solution). Thus, the value given by [1964BAK] was not used in this report. 6.8.5 Pb10(PO4)6(OH)2(hydroxylapatite) For lead hydroxylapatite (Pbi0(PO4)6(OH)2(cr)), [1976RAO] reported a slightly higher solubility than for calcium hydroxylapatite which leads to the conclusion that the precipitation of lead hydroxylapatite can not be expected in the presence of calcium. However, the formation of mixed lead-calcium hydroxylapatite has also been observed [1976RAO]. 193 JNC TN8400 99-011 6.8.6 Additional data compiled for the lead phosphate system Table 6.22: Additional experimentally determined data for the lead phosphate system, according to the 2+ 3 2m 3n equilibrium: mPb + nPO4 " <=> Pbm(PC>4)n " . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text in Section 6.8. Method: pot = potentiometry, sol = solubility measurements. Reference Comments KM) Medium Method 2+ } log K*S0: Pb + P04 - + //+<=> PbHPO4(s) 22.28 ' [1929MIL/JOW] T=310K, I=dil 0 PbCl2> Na2HPO4 sol 22.20 ' [1929MIL/JOW] T= 298.15 K, I=dil 0 PbCl2, Na2HPO4 sol 23.71 2 [1932JOW/PRI] T= 310 K, 1=0.02 0 NaCl sol 23.75 [1969AWA/ELH1 T= 303 K, I=dil 0 H,PO4, NaH,PO, pot 2+ 3 log K*S0:3Pb + 2PO4 - <=> Pb3(PO4)2(s 42.0 ' [1929MIL/JOW] T=31OK,I=dil 0 PbCl2, Na2HPO4 sol 42.1 ' [1929MIL/JOW] T= 298.15 K,I=dil 0 PbCl2, Na2HPO4 sol 43.53 2 [1932JOW/PRI] T= 310 K, 1=0.02 0 NaCl sol 42.03 ri969AWA/ELH1 T= 303 K, I=dil 0 H,PO4, NaH,P04 pot 2+ 3 + log K*S0:5Pb + 3PO4 ' + Cl- <=> Pb5(PO4)3Cl(s) (chloropyromorphite)+ H 79.12 2 [1932JOW/PRI] T= 310 K, 1=0.16 0 NaCl sol 3 34.50 H964BAK] T= 298.15 K, 1=0.008-0.028 0 PbCl,, Na,HP04 sol 1 extrapolated to 1=0 with Debye-Huckel equation; later criticized by [1932JOW/PRI] 2 extrapolated to 1=0 with Debye-Hiickel equation 3 3 J given by [1964BAK] assuming equilibria with PO4 ' in a Na2HPO4 solution 194 JNC TN8400 99-011 Table 6.23: Thermodynamic data for the lead phosphate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log |3m n Reference Comments I (M) 2 + log P,.u: Pb H <=> PbHPO4° 15.45 [1976SMI/MAR] T= 298.15 K, 1=0 0 15.24 [1983LAN] T= 298.15 K, 1=0 0 15.45 [1984VIE/TAR] T= 298.15 K, 1=0 0 23.28 [1985GAL] T= 298.15 K, I=n/a 16.59 [1987BRO/WAN1 T= 298.15 K, 1=0 2+ + 2 log Pi.2.2- Pb + 2P0J- + 2H <=> Pb(HPO4)2 32.52 ri987BROAVANl T= 298.15 K, 1=0 2+ 3 + log PUJ: Pb + 3PO4 - + 3H <=> Pb(HPO4)/- 48.02 [1987BRO/WAN] T= 298.15 K, 1=0 2+ + 6 log P,,4.4: Pb + 4PO/- + 4H & Pb(HPO4)4 ' 63.21 n987BRO/WANl T= 298.15 K, 1=0 2+ 3 + + log Pu.i: Pb + PO4 - + 2H PbH2PO4 21.05 [1976SMI/MAR] T= 298. 15 K, 1=0 0 19.90 [1983LAN] 298. 15 K, 1=0 0 21.05 [1984VIE/TAR] T= 298. 15 K, 1=0 0 20.89 ri987BROAVANl 298. 15 K, 1=0 0 2+ + log P,.4.2: Pb + 2PO/- + 4H <=> Pb (H2PO4)2° 41.09 ri987BRO/WANl T= 298.15 K, 1=0 2+ + log P,.6j- Pb + 3PO/- + 6H <=> Pb (H2PO4)3- 60.85 f!987BROAVAN1 T= 298.15 K, 1=0 : p 2+ 3 + 2 log P,,8.< b + 4PO4 ' + 8H <=> Pb (H2PO4)4 - 80.27 [1987BROAVAN1 T= 298.15 K, 1=0 195 JNC TN8400 99-011 Table 6.23: continued 2+ 3 + log K*S0: Pb + PO4 - + H <=> PbHPO>4( secondary lead phosphate) 23.41 [1952LAT] T= 298.15 K, 1=0 0 26.82 [1971NAU/RYZ] T= 298.15 K,I=n/a 23.78 [1976SMI/MAR] T= 298.15 K, 1=0 0 23.78 [1980CLE/JOH] T=298.15K, I=n/a 26.78 [1982WAG/EVA] T= 298.15 K,I=n/a 23.78 [1984VIE/TAR] T= 298.15 K,I=n/a 23.28 [1985GAL1 T= 298.15 K,I=n/a 2+ 3 log K*S0:3Pb + 2PO4 ' <=> Pb3(PO4)2(tertiary lead phosphate) 54 [1952LAT] T= 298.15 K, 1=0 0 43.53 [1976SMI/MAR] T= 298.15 K, 1=0 0 42.08 [1977TAR/VIE] T= 298.15 K, I=n/a 44.4 [1980CLE/JOH] T= 298.15 K, I=n/a 44.40 [1984VIE/TAR] T= 298.15 K,I=n/a 43.41 [1985GAL] T= 298.15 K,I=n/a 42.18 [1988PHI/HAL1 T= 298.15 K, I=n/a 2+ 3+ 3 + log K*so.- Pb + 3Al + 2PO4 ' 5H2O <^> PbAl3(PO4)2(OH)5(plumbogummite) + 5H 99.30 ' [1984NRI2] T= 298.15 K, I=n/a 2+ + log K*S0:5Pb + 3PO/- + H2O « Pb5(PO4)3OH(hydroxypyromorphite) + H 62.80 [1980CLE/JOH] T= 298.15 K, I=n/a 62.80 [1984VIE/TAR1 T= 298.15 K, I=n/a 2+ 3 + log K*sa-5Pb + 3PO4 - + Ci <=> Pb5(P04)3Cl(chloropyromor,phite)+ H 9.11 [1971NAU/RYZ] T= 298.15 K, I=n/a 79.12 [1972NRI] T= 298.15 K,I=n/a 84.45 [1976HEM] T= 298.15 K,I=n/a 84.4 [1980CLE/JOH] T= 298.15 K,I=n/a 84.40 ' [1984NRI2] T= 298.15 K,I=n/a 84.40 [1984VIE/TAR1 T= 298.15 K, I=n/a 2+ 3 + log K*S0:5Pb + 3PO4 - + F- <=> Pb5(P04)3F(fluoropyromorphite)+ H 71.6 ri980CLE/JOH] T=298.15 K, I=n/a [1984NRI2] claims that chloropyromorphites and plumbogummite or mixtures of them precipitate in soils 196 JNC TN8400 99-011 6.9 Lead sulfate system Lead(II) forms complexes with sulfates in aqueous solutions and several authors determined stability constants (Table 6.24). Also, a number of solid lead sulfate compounds are known to exist. The data used in this report for the calculations of the formation constants of lead(II) sulfate complexes are compiled in Table 6.24 and shown in Figure 6.27 and Figure 6.28. Additional data for the lead(II) sulfate system that were not selected for the calculation of log (3° values in this report are compiled in Table 6.25 and 6.26. Table 6.24: Experimentally determined equilibrium data compiled for the lead sulfate system, 2+ 2 2m 2n according to the equilibrium: mPb + nSO4 " <=> Pbm(SO4)n - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: con = conductivity, el = electrophoresis, pol = polarography, pot = potentiometry, sol = solubility measurements Reference Comments KM) Medium Method log Pm,n 2+ 2 log fa ,: Pb + SO4 -& PbSO4° 2.07 [1969DRY/TVA] T=298 K, 1=0.2 0.2 NaClO4/Na2SO,! SOl 2.70 1 [1970GAR/NAN] T=298 K, I=dil 0 NaC104 pot 2.75 l [1970GAR/NAN] T=298 K, I=dil 0 no con 0.70 [1971BON1] T=298 K, 1=3 3 NaC104 pol 0.74 [1972BON] T=298 K, 1=3 3 NaC104 pol 1.05 [1982ROH] T=296 K, 1=0.7 0.7 NaC104 el 2 1.20 [1989NYHAVIK] T=298.15 K, 1=1 1 NaClO4 pol 2+ 2 2- log Pj 2: Pb + 2SO4 -<^Pb(SO4)2 1.98 [1971BON1] T=298 K, 1=3 3 NaC104 pol 2.00 [1972BON] T=298 K, 1=3 3 NaClO4 pol 1.18 [1982ROH] T=296 K, 1=0.7 0.7 NaClO4 el 197 JNC TN8400 99-011 2+ 2 log Kso: Pb + SO4 -<=> PbSO4(anglesite) 7.64 [1942KOL/PER] T= 298 K, I=dil 5E-04 no sol 3 7.70 [1959JAG] T= 298 K, I=dil 0 HNO3 sol 6.20 [1961RAM/STE] T= 298 K, 1=1 1 NaC104 sol 7.03 [1969DRY/IVA] T= 298 K, 1=0.2 0.2 NaClO4/Na2SO4 sol 7.76 4 [1970LIT/NAN] T= 298 K, I=dil 0 no pot 1 extrapolated to 1=0 with Davies equation, measured at 1=0.001 M 2 0.0001-0.1 raM Pb; mean of 6 measurements 3 measured at 1=0.001 M, extrapolated to 1=0 by [ 1959JAG] as a function of 1°5 4 measured at 1=0.001 M, extrapolated to 1=0 with Davies equation 6.9.1 Lead sulfate complexes Several values of stability constants for PbSC>40 are reported in literature (Table 6.24), these data show a good agreement. Extrapolation of these measurements to 1=0 with the SIT method is shown in Figure 6.27: PbSO4° log (3°u= 2.82 2 2 2SO4 - Pb(SO4)2 - log p°lt2= 2.37 2 For Pb(SO4) " the log (3^2 value of 2.37 (Figure 6.28) is calculated from only two independent measurements ([1971BON1, 1972BON], 1=3 and [1982ROH], 1=0.7). 2+ 2 Pb + SO4 " <:* PbSO4° - : : 4.5 - 5 _\-Y ; V; •;. V*;/. '\__7V.-.- „->,» 4 - Q 3.5 • 00 o 3 • 0 ) . 2.5 • , - o; - COL D) JO 2 • 1.5 • 1 - 0.5 n - ^——t ^- 1 2 lm> molal 2+ 2 Figure 6.27: Plot of log p1(1 + 8 D vs. Im for the reaction : Pb + SO4 - <=> PbSO4° at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.01; log (3° \t\ = 2.82. Calculated from data compiled in Table 6.24. 198 JNC TN8400 99-011 2+ 2 2 Pb + 2SO4 " <=> Pb(SO4): " a • 4 5 • 4 • + 3 • j? 2 1.5 ; : E ! : - [ z "1 • Tb:;! J;=i I! i > • ^: I ^ • ^ H~f:; L ^: Ti - - i'; M [J: =;-1" ;1 J n ^ - M^Tr 1:1! I :fii :=^* 1^ ^l^ ';!: M i IWiii ^ I= ^1 ^U1 Ij 1= !^ i^; J^!^ ^i': i; - ^ -f! /i \-^V-CfW^^V<^\jAf^' •-• 1 0.5 • 0 '•"•: • '''•'•'":::'r':;:':-:;*::p:}r'• •':j:'::"':|: 1;-!j'":-w:if 'H• •''-':•vt!:V:^T'l^ySff •>•-:-r':- lm, molal 2+ 2 2 Figure 6.28: Plot of log pi,2 + 8 D vs. Im for the reaction : Pb + 2SO4 " <=> Pb(SO4)2 - at 25 °C. The straight line shows the result of the linear regression: Ae = -0.43; log P°i>2 = 2.37. Calculated from data compiled in Table 6.24. 6.9.2 PbSO4(anglesite) A number of solid lead sulfate species are reported in different compilations (Table 6.26). Most of these solids form only at higher temperatures, however [1989KEL, 1995MAR/MAC]. [1995MAR/MAC] showed that only anglesite (PbSO^cr)) precipitates from aqueous solutions at room temperature. From several solubility measurements, a log K*so of 7.81 can be calculated (Figure 6.29), a value which is in excellent agreement with the value of 7.80 given by [1995MAR/MAC]. PbSO4(s) log K*°so= 7.81 199 JNC TN8400 99-011 2+ 2 Pb + SO4 - «^ PbS04(s) in I U ^ . •• v •" 9.5 , • •- 9 y = 0;09x + 7.81 ' J;' oo 8.5 + 8 ., o -" t 7 • 6.5 6 5.5 1_- fi \—• 0.5 1 1.5 lmi molal 2+ 2 Figure 6.29: Plot of log K*So + 8 D vs. Im for the reaction : Pb + 2SO4 - <=> PbSO4(s) at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.09; log K*°so = 7.81. Calculated from data compiled in Table 6.24. 6.9.3 Hinsdalite: PbAl3PO4SO4(OH)6(cr) The values given in different compilations for hinsdalite (PbAl3PO4SO4(OH)6(cr)) (compiled in Table 6.26), spread over a wide range and are based either on an estimate [1984NRI2] or on the measurements by [1964BAK]. [1964BAK] (Table 6.25) determined a log K*So of 2.11 for hinsdalite using an equilibration time of several months. However, a closer examination of his data led to the conclusion that [1964BAK] assumed all phosphate to be present as PO43- (at a pH of 3). This value is therefore not used in this report and no log K*so for hinsdalite can be proposed in this report. 200 JNC TN8400 99-011 Table 6.25: Additional experimentally determined data for the lead sulfate system, according to the 2+ 2 2m 2n equilibrium: mPb + nSO4 " <=> Pbm(SO4)n " . These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility measurements, SC>2= pSC>2 measurements at high temperatures. log Pm,n Reference Comments Medium Method 2+ 2 log Pu: Pb + SO4 -<=> PbSO4° 2.62 ' [1969DRY/TVA] T= 298 K, 1=0 0 NaClO4/Na?SO4 sol 2+ 2 log Kso: Pb + SO4 -<=> PbSO4(s, anglesite) 7.82 ' [1969DRY/TVA] T= 298 K, 1=0 0 NaC104/Na2S04 sol 7.24 [198OWAL/SIN] T= 298 K, I=dil H2SO4 sol 15.40 2 [1989KEL] T= 298.15 K,I=n/a n/a SO, 7.80 3 [1993MAC/PAG] T = 298.15 K, 1=0 0 water sol 7.80 4 [1995MAR/MAC1 T = 298.15 K, 1=0 0 water sol 2+ 2 + log Kso: 2Pb + SO4 -+ H2O <=> Pb2OSO4(s, lanarkite) +2H 5.83 2 [1989KEL] T= 298.15 K, I=n/a n/a SO, 2+ + log Kso: 3Pb + SO/-+ 2H2O <=> Pb3O2SO4(s) + 4H -7.22 2 [1989KEL] T= 298.15 K, I=n/a n/a SO, 2+ 2 + log Kso: 5Pb + SO4 -+ 4H2O <=> Pb5O4SO4(s) + 8H -29.61 2 [1989KEL] T= 298.15 K, I=n/a' n/a SO, 2+ 3 2 3+ + log Kso: Pb + PO4 - + SO4 - 3Al + 6H2O <=> PbAl3PO4SO4(OH)6(s, hinsdalite) + 6H . 5 2.11 H964BAK] T= 298 K, 1=0.008-0.028 PbCl,, Na,HPO4 sol ' not indicated how extrapolation to 1=0 was made 2 measurements at high temperatures; this value for PbSO4 may refer to a high temperature modification 3 [1993MAC/PAG] compare lead solubility calculated with thermodynamic data with experimental lead solubility 4 [1995MAR/MAC] compare lead solubility calculated with thermodynamic data with experimental lead solubility 5 3 calculated by [1964BAK] assuming equilibria with PO4 " at a pH of 3 201 JNC TN8400 99-011 Table 6.26: Thermodynamic data for the lead sulfate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log Pm,n Reference Comments Medium 2+ log Pu: Pb PbSO4° 2.7 [1972ZIR/YAM] T= 298.15 K, 1=0 0 2.62 [1976HEM] T= 298.15 K, I=n/a 2.75 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.74 [1976SMI/MAR] T= 298.15 K, 1=3 3 2.75 [1981TURAVHI] T= 298.15 K, 1=0 0 2.86 [1983LAN] T= 298.15 K, 1=0 0 2.58 [1985BAB/MAT] T= 298.15 K, 1=0 0 2.28 [1987BRO/WAN] T= 298.15 K, 1=0 0 1.29 [1988BYRyKUM] T=291 K, 1=0.7 0.7 sea water 2.65 [1988PHI/HAL1 T= 298.15 K, 1=0 0 2.69 [1989SMI/MAR] T=298.15K, 1=0 0 2.07 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 0.74 [1989SMI/MAR1 T= 298.15 K, 1=3 3 2+ 2 2 log P12: Pb + 2SO4 -<=> Pb(SO4)2 - 3.47 [1972ZIR/YAM] T= 298.15 K, 1=0 0 3.47 [1976HEM] T= 298.15 K,I=n/a 1.99 [1976SMI/MAR] T= 298.15 K, 1=3 3 4.51 [1981TURAVHI] T= 298.15 K, 1=0 ^ 0 3.68 [1987BRO/WAN] T=298.15K, 1=0 ' 0 2.48 fl988BYR/KUMl T=291 K, 1=0.7 0.7 sea water 2+ log Pi,.,: Pb + 3SO/'<=> Pb(SO4)/- 4.45 [1987BRO/WAN] T= 298.15 K, 1=0 0 2+ 2 log p1A: Pb + 4SO4 -& Pb(SO4)/- 4.70 ri987BROAVAN1 T= 298.15 K, 1=0 + log Ks0: FV + SO/-<=> PbSO4(anglesite) 7.50 [1952LAT] T= 298.15 K,I=n/a 7.87 [1968ROBAVAL] T= 298.15 K,I=n/a 7.75 [1969HEL] T= 298.15 K, 1=0 0 7.75 [1971NAU/RYZ] T= 298.15 K,I=n/a 7.87 [1973BAR/KNA] T= 298.15 K,I=n/a 7.79 [1976SMI/MAR] T= 298.15 K, 1=0 0 6.20 [1976SMI/MAR1 T= 298.15 K, 1=1 1 202 JNC TN8400 99-011 Table 6.26: continued 7.50 [1977PAU] T= 298.15 K, I=n/a 7.88 [1978COD] T= 298.15 K, I=n/a 7.82 [1978ROB/HEM2] T= 298.15 K, I=n/a 7.86 [1979KUB/ALC] T= 298.15 K, I=n/a 7.60 [1980CLE/JOH] T= 298.15 K, 1=0 0 7.83 [1980BEN/TEA] T= 298.15 K, I=n/a 7.85 [1981STU/MOR] T= 298.15 K, I=n/a 7.50 [1982PAU] T= 298.15 K, I=n/a 7.83 [1982WAG/EVA] T= 298.15 K, I=n/a 7.87 [1983LAN] T= 298.15 K, I=n/a 7.70 [1984NRI2] T= 298.15 K, I=n/a 7.50 [1985BAB/MAT] T= 298.15 K, I=n/a 8.02 [1985GAL] T= 298.15 K, I=n/a 7.85 [1988PHI/HAL] T= 298.15 K, I=n/a 7.85 [1989COX/WAG] T= 298.15 K, I=n/a 7.79 [1989SMI/MAR] T= 298.15 K, 1=0 0 7.03 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 6.20 [1989SMI/MAR1 T= 298.15 K, 1=1 1_ 2+ + log Kso: 2Pb + SO/-+ H2O <=> Pb2OSO4(lanarkite) + 2H 9.34 [1952LAT] T= 298.15 K, I=n/a 0.44 [1971NAU/RYZ] T= 298.15 K, I=n/a 3.22 [1973BAR/KNA] T= 298.15 K, I=n/a 0.42 [1977BAR/KNA] T= 298.15 K, I=n/a 0.37 [1980BEN/TEA] T= 298.15 K, I=n/a 0.37 [1983LAN] T= 298.15 K, I=n/a 9.34 [1985GAL] T= 298.15 K, I=n/a 0.50 [1988PHI/HAL] T= 298.15 K, I=n/a 2+ 2 + log Kso: 3Pb + SO4 -+ 2H2O <=> Pb3O2SO4(s) + 4H ' -1.72 [1973BAR/KNA] T= 298.15 K, I=n/a -10.70 [1977BAR/KNA] T= 298.15 K, I=n/a -10.79 [1980BEN/TEA] T= 298.15 K, I=n/a -10.79 [1982WAG/EVA] T= 298.15 K, I=n/a 2+ 2 + log Kso: 4Pb + SO4 '+ 3H2O «• Pb4O3SO4(s) + 6H -21.91 [1977BAR/KNA] T= 298.15 K, I=n/a -22.02 [1980BENATEA] T= 298.15 K, I=n/a -22.01 [1982WAG/EVA1 T= 298.15 K, I=n/a 2+ + lag Kso: 5Pb + SO/ + 4H2O <=> Pb5O4SO4(s) + 8H -38.24 [1973BAR/KNA] T= 298.15 K, I=n/a 2+ 3 2 3+ + log Kso: Pb + PO4 - + SO4 ' + 3Al + 6H2O <=> PbAl3PO4SO4(OH)6(hinsdalite) + 6H 1.63 [1971NAU/RYZ] T= 298.15 K, I=n/a 15.10 [1984NRI2] T= 298.15 K, I=n/a 203 JNC TN8400 99-011 6.10 Lead sulfide system Lead forms complexes with sulfide in aqueous solutions, see Section 6.10.1. PbS(s), galena, has a very low solubility and often controls lead solubility in reducing environments. The data used for the calculations of the formation constants of lead(II) sulfide complexes are compiled in Table 6.27. Additional data for the lead(II) sulfide system that are not chosen for the calculation of log (3° values in this report are compiled in Table 6.28 and 6.29. Table 6.27: Experimentally determined equilibrium data compiled for the lead sulfide system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 6.13: 'Comments on selected references'. Method: col = colorimetric analysis, sol = solubility measurements. cr R Reference Comments I (M) Medium Method 6 Fm,n + log Kia: PbS(galena) + HS' + H <=> Pb(HS)2° 0.13 l [1953HEM] T= 298.15 K, I=dil 0 n/a col 0.21 ! [1979GIO/BAR] T= 303 K, 1=0.5-3 0 n/a sol_ log KJi3: PbS(galena)+ 2HS' + H+ <=>Pb(HS)3- 1.43 i [1953HEM] T= 298.15 K, I=dil 0 n/a col 1.41 l [1979GIO/BAR] T= 303 K, 1=0.5-3 0 n/a sol_ 2+ log K*so: Pb + HS-<=>PbS(s, galena) + H+ 11.79 [1984UHL/HEL] T= 298. 15 K, 1=0. 16 0. 16 H2S sol 11.95 [1984UHL/HEL] T= 298. 15 K, 1=0. 02 0.02 H2S sol 11.91 [1984UHL/HEL] T= 298. 15 K, 1=0. 03 0.03 H?s sol extrapolated by [the respective authors to 1=0 with extended Debye-Huckel equation 204 JNC TN8400 99-011 6.10.1 Lead sulfide complexes Determination of stability constants for Pb(HS)2° and Pb(HS)3~ (with regard to PbS(s)) are reported by [1953HEM] and [1979GIO/BAR]. [1984UHL/HEL] recalculated the values of [1953HEM] and [1979GIO/BAR]. However, from the comments given by [1984UHL/HEL], their calculations are not traceable and are therefore not used in this report. [1953HEM] and [1979GIO/BAR] corrected their results with the Debye-Hiickel equation to 1=0 (Table 6.27). As the values are similar, the mean may be used: log K°]2 = 0.17 for the reaction + + PbS(galena) + HS" + H <^ Pb(HS)2°, and log K°u =1.42 for PbS(galena) + 2HS" + H <=> Pb(HS)3-). Using the log K*°So of 12.17 calculated in Section 6.10.2, the following constants can be calculated: 2+ O Pb + 2HS- « Pb(HS)2° log (3 1)2 = 12.34 2 Pb + + 3HS- o Pb(HS)3- log p°u = 13.59 6.10.2 Galena (PbS) Data reported in the literature for the equilibrium Pb2+ + 2HS" <=> PbS(galena) + 2H+ are often based either on calculations from AS and AH values measured at high temperature or on the selection made by [1953HEM] (Table 6.28 and 6.29). Experimentally determined values are scarce. The solubility product of galena (PbS) with respect to Pb(HS)2° and Pb(HS)3- can be determined quite precisely [1953HEM, 1956KIV/RIN]). The complex formation between lead and sulfide, however, is so strong, that direct determination of the equilibria Pb2+ + 2HS' <=> PbS(galena) + 2H+ is difficult. [1956KIV/RIN] tried to determine the Pb2+ concentration directly via competitive complex formation with chloride ions. This method, however, is prone to errors as the complex formation with chloride is much weaker than the complex formation of Pb with HS\ The log K*so of galena was determined more precise by [1984UHL/HEL] the with help of competitive complex formation of Pb with EDTA and TRIS. Extrapolation of the values [1984UHL/HEL] to I = 0 with the SIT equation is shown in Figure 6.30: 2+ Pb + HS- <=> PbS(galena) + H+ log K*°so= 12.17 205 JNC TN8400 99 -Oil Pb2+ + HS- <=> PbS(galena) + H+ 15 14.5 14 13.5 •; 0.2 0.4 0.6 L molal 2+ Figure 6.30: Plot of log K*So + 4 D vs. Im for the reaction : Pb + HS" <=> PbS + H+ at 25 °C. The straight line shows the result of the linear regression: As = - 0.82; log K*°so = 12.17. Calculated from data compiled in Table 6.27. 6.10.3 Additional data compiled for the lead sulfide system Table 6.28: Additional experimentally determined data for the.lead sulfide system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: emf = emf measurements, pol = polarography, sol = solubility measurements. log Pm.n Reference Comments KM) Medium Method + log K'SQ: Pb2+ + HS'<=> PbS(galena) + H 14.94 1 [1956KTV/RIN] T= 298.15 K, 1=0.8-1.1 0 HC1 pol 15.66 2 [1977SHA] T= 298.15 K, 1=0 0 n/a emf 15.66 2 [1981SHA/MIS] T= 298.15 K, 1=0 0 n/a emf 3 12.42 [1984UHL/HEL] T= 298.15 K, 1=0 0 H2S sol 11.78 3 [1984UHIVHEL1 T= 298.15 K, 1=0.7 0.7 H,S sol 1 extrapolated to 1=0 by [1956KIV/RIN] with Davies equation; calculated in this report using a log K(H2S/HS) = 6.99, log KH(H2S(g)/H2S(aq)) = 1.02 measured at 633-698 K, AH, AS also determined. extrapolated by [1984UHL/HEL] to 1=0 and 1=0.7 with the Davies equation 206 JNC TN8400 99-011 Table 6.29: Thermodynamic data for the lead sulfide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K Reference Comments I (M) log Klt2: PbS + HS- + H+ <=> Pb(HS)2° -0.59 [1983LAN] T= 298.15 K, I=n/a 0 0.3 ' [1984UHL/HEL] T= 298.15 K, 1=0 0 -0.05 2 [1984UHL/HEL1 T= 298.15 K, 1=0 0 + log Ku: PbS + 2HS- + H <=> Pb(HS)f 1.42 [1983LAN] T= 298.15 K, I=n/a 1.55 2 ri984UHL/HELl T= 298.15 K, 1=0 2+ + log K*so: Pb + HS-<=> PbS(galena) + H 14.12 [1952LAT] T= 298.15 K, I=n/a 14.58 [1953HEM] T= 298.15 K, I=dil 14.06 [1964HIR] T= 298.15 K, I=n/a 14.67 [1969HEL] T= 298.15 K, I=n/a 15.19 [1971NAU/RYZ] T= 298.15 K, I=n/a 15.19 [1973BAR/KNA] T= 298.15 K, I=n/a 14.82 [1974MIL] T= 298.15 K, I=n/a 13.60 3 [1976SMI/MAR] T= 298.15 K, 1=0 0 14.82 [1977BAR/KNA] T= 298.15 K, I=n/a 14.70 [1978ROB/HEM2] T= 298.15 K, I=n/a 14.84 [1979KUB/ALC] T= 298.15 K, I=n/a 14.81 [1980BEN/TEA] T= 298.15 K,J=n/a 15.70 4 [1980CLE/JOH] T=298.15 K, I=n/a 15.16 [1981STU/MOR] T= 298.15 K, I=n/a 15.16 [1982WAG/EVA] T= 298.15 K, I=n/a 14.81 [1983LAN] T= 298.15 K, I=n/a 15.16 [1983SAN/BAR] T= 298.15 K, I=n/a 14.84 [1985CHA/DAV] T= 298.15 K, I=n/a 15.19 [1985GAL] T= 298.15 K, I=n/a 14.73 [1985MUL] T= 298.15 K, I=n/a 13.51 5 [1986MYE] T= 298.15 K, I=n/a 14.77 [1988PHI/HAL1 T= 298.15 K, I=n/a 1 recalculation of measurements by [1953HEM] 2 recalculation of measurements by [1979GIO/BAR]; the constants used by [1984UHL/HEL] do not agree in all cases with the original references 3 calculated with pK(HS7S2>13.9 as given in [1976SMI/MAR] 4 corrections made are not traceable 5 calculations of [1986MYE] are based on the values from [1976SMI/MAR] and new values determined by [1986MYE] for the protonation of S2" to HS". 207 JNC TN8400 99-011 6.11 Redox equilibria 6.11.1 Pb2+/Pb(cr) Based on a critical review of available experimentally determined data the COD AT A [1989COX/WAG] Key Values recommended for Pb2+ (aq) an AfG° value of -24.20 kJ/mol resulting in: Pb2++2e- <=> Pb(cr) log K° = -4.25 E° =-0.126 V 6.11.2 Pb2+/Pb4+ Solid and dissolved Pb(IV) species exist only under very oxidizing conditions [1976BAE/MES, 1985GAL]. From the data given by [1952LAT and 1985GAL] a tentative log K value of 57.23 can be calculated for the redox equilibrium Pb2+ <=> Pb4+ + 2e" (Table 6.31). However, no direct measurements are available. 6.11:3 PbO2andPb3O4 [1985GAL] reviewed available experimental data for PbO2 and Pb3C>4. He gives for the formation of PbO2 from Pb(cr) a E° of 1.690 V and states that PbO2 is thermodynamically unstable under acidic conditions, where this E° value is more positive than that of the oxidation of H2O to O2. From the data given in [1985GAL] the following tentative formation constants can be calculated: 2+ 3Pb + 4H2O o Pb3O4(s) + 8H++2e- log K*°so = - 70.98 2+ Pb + 2H2O <=> PbO2(s) + 4H++2e- log K*°so = - 48.98. 208 JNC TN8400 99-011 6.11.4 Data compiled for the lead redox system Table 6.30: Experimentally determined data for the lead redox system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 6.13: 'Comments on selected references'. Method: pot = potentiometry. log K Reference Comments I (M) Medium Method log K: Pb2+ + 2e <=> Pb(0)+ -4.27 [1929CAR] T= 298.15 K, 1=0.001-0.15 0 water pot -14.20 ' [1945LIN] T=298.15,1=0.5 0.5 KC1 pot -4.18 2 [1971VAS/GLA] T= 298.15 K, 1=0 0 n/a pot 1 probably formation of PbCl complexes, hydrolysis not controlled 2 no experimental details reported Table 6.31: Thermodynamic data for the lead redox system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK Reference Comments I CM) log K: Pb2+ Hh 2e- « Pb(0) -4.26 [1952LAT] T= 298.15 K, 1=0 0 -4.28 [1969HEL] T= 298.15 K, 1=0 0 -4.26 [1977PAU] T= 298.15 K, 1=0 0 -4.20 [1978COD] T= 298.15 K, 1=0 0 -4.27 [1978ROB/HEM2] T= 298.15 K, 1=0 0 -4.28 [198OBEN/TEA] T= 298.15 K, 1=0 0 -4.21 [1981HEL/KIR] T= 298.15 K, 1=0 0 -4.27 [1981STU/MOR] T= 298.15 K, 1=0 0 -4.26 [1982PAU] T= 298.15 K,I=0 0 -4.28 [1982WAG/EVA] T= 298.15 K, 1=0 0 -4.28 [1983LAN] T= 298.15 K, 1=0 0 -4.28 [1983SAN/BAR] T= 298.15 K, 1=0 0 -4.27 [1984VIE/TAR] T= 298.15 K, 1=0 0 -4.28 [1985BAB/MAT] T= 298.15 K, 1=0 0 -4.26 [1985GAL] T= 298.15 K, 1=0 0 -4.24 [1985RAI/RYA] T= 298.15 K,I=n/a -4.21 [1988PHI/HAL] T= 298.15 K, 1=0 0 -4.25 [1989COXAVAG1 T= 298.15 K, 1=0 0 log K: Pb2+ <^ Pb4+ + 2e- 57.23 [1952LAT] T= 298.15 K, 1=0 0 -48.76 [1983LAN] T= 298.15 K, 1=0 0 57.23 [1985GAL] T= 298.15 K, 1=0 0 209 JNC TN8400 99-011 Table 6.31: continued 2+ + log K*so: 3Pb + 4H2O <=> Pb3O4(s) + 8H + 2c -66.38 [1929MIL] T= 298.15 K, I=n/a -70.73 [1952LAT] T= 298.15 K, I=n/a -70.81 [1954COU] T= 298.15 K,I=n/a -70.84 [1963WIC/BLO] T= 298.15 K, I=n/a -70.99 [1973BAR/KNA] . T= 298.15 K,I=n/a -73.52 [1977BAR/KNA] T= 298.15 K,I=n/a -73.57 [1978ROB/HEM2] T= 298.15 K, I=n/a -73.52 [1979KUB/ALC] T= 298.15 K,I=n/a -70.60 [1979PAT/OBR] T = 298.15 K, 1=0 0 -73.58 [198OBEN/TEA] T= 298.15 K,I=n/a -73.58 [1980SCH] T = 298.15 K, 1=0 0 -73.59 [1981STU/MOR] T= 298.15 K,I=n/a -73.53 [1982PAN] T= 298.15 K,I=n/a -73.59 [1982WAG/EVA] T= 298.15 K, I=n/a -73.60 [1983LAN] T= 298.15 K,I=n/a -73.60 [1983S AN/BAR] T= 298.15 K, I=n/a -70.73 [1985BAB/MAT] T= 298.15 K,I=n/a -73.52 [1985CHA/DAV] T= 298.15 K,I=n/a -70.98 [1985GAL] T= 298.15 K, I=n/a -73.58 [1988PHI/HAL1 T= 298.15 K,I=n/a 2+ + log K*so: Pb + 2H2O <=> PbO2(s) + 4H• +2c -49.04 [1929MEL] T= 298.15 K, I=n/a -49.14 [1952LAT] T= 298.15 K,I=n/a -49.01 [1954COU] T= 298.15 K,I=n/a -49.01 [1963WIC/BLO] T= 298.15 K,I=n/a -49.23 [1971NAU/RYZ] T= 298.15 K,I=n/a -50.12 [1973BAR/KNA] T= 298.15 K,I=n/a -49.60 [1977BAR/KNA] T= 298.15 K,I=n/a -49.62 [1978ROB/HEM2] T= 298.15 K,I=n/a -49.60 [1979KUB/ALC] T= 298.15 K,I=n/a' -49.02 [1979PAT/OBR] T = 298.15 K, 1=0 0 -49.62 [1980BEN/TEA] T= 298.15 K,I=n/a -49.24 [1980SCH] T = 298.15 K, 1=0 0 -49.25 [1981STU/MOR] T= 298.15 K,I=n/a -49.09 [1982PAN] T= 298.15 K, I=n/a -49.26 [1982WAG/EVA] T= 298.15 K,I=n/a -49.26 [1983LAN] T= 298.15 K,I=n/a -49.26 [1983S AN/BAR] T= 298.15 K,I=n/a -48.98 [1985BAB/MAT] T= 298.15 K,I=n/a -49.60 [1985CHA/DAV] T= 298.15 K, I=n/a -48.98 [1985GAL] T= 298.15 K, I=n/a -49.62 ri988PHI/HALl T= 298.15 K,I=n/a 2+ + log K*so: Pb + 4H2O <=> Pb(OH)4(s) + 4H + 2c 35.13 [1976TAR/GAR] T= 298.15 K, I=n/a 2 + + log K*so: 2Pb + 3H2O <=> Pb2O3(s) + 6H + 2c -61.00 [1985GAL1 T= 298.15 K, I=n/a 2+ log K*so: Pb + 1.57H2O <=> PbO!57(s) +3.14H++ 1.14c -32.47 H985GAL1 T= 298.15 K, I=n/a 210 JNC TN8400 99-011 6.12 Lead(IV) The hydrolysis of Pb(OH)4° in solution is not known with reliability: Based on different reports, [1976BAE/MES] give the following estimates: 4+ Pb(OH)4° + 4H+ <=> Pb +4H2O log p°li0 = -4.0 + Pb(OH)4° + H <=> Pb(OH)3+ + H2O log (3\3 = -0.8 2 + Pb(OH)4°+2H2O o Pb(OH)6 - + 2H log [3\6 =-28.3 Pb(OH)4° o PbO2 + 2H2O log K*°so = 4.0 No newer experimental data have been found and no values for the hydrolysis of Pb(IV) are recommended here. Table 6.32: Experimentally determined data for the hydrolysis of lead(IV). These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility measurements. log K Reference Comments I (M) Medium Method 2 logpI:6.Pb(OH)4°+2H2O <=> Pb(OH)6 ' + 2H+ -28.06 [1969CHA] T= 298.15 K, I=n/a NaOH sol 0 log K*S0Pb(OH)4 <=> PbO2 + 2H2 <-4.00 [1969CHA1 T= 298.15 K, I=n/a NaOH sol Table 6.33: Thermodynamic data for the hydrolysis of lead(IV) taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K Reference Comments I (M) + 4+ log p,i0.- Pb(OH)4° + 4H » Pb + 4H2O -4.00 [1963FEI/SCH] T= 298.15 K, 1=0 0 -4.00 ' [1976BAE/MES] T= 298. 15 K, I=n/a -4.26 [1979PAT/OBR] T = 298.15 K,1=0 0 + + log Pu: Pb(OH)4° + H <=> Pb(OH)3 + H2O -0.80 ' [1976BAE/MES] T= 298.15 K, I=n/a 211 JNC TN8400 99-011 Table 6.33: continued 0 2 + bgP,,6:Pb(OH)4 +2H2O <=> Pb(OH)6 - + 2H -28.5 [1963FEI/SCH] T= 298. 15 K,1=0 0 -28.30 ' [1976BAE/MES] T= 298. 15 K, I=n/a -27.32 2 [1979PAT/OBR1 T = 298.15 K, 1=0 0 log K*S0Pb(OH)4° <=> PbO2 + 2H2 -4.00 ' [1976BAEMES1 T= 298.15 K, I=n/a 1 reported by [1976BAE/MES] from different sources 2 data originally from Pourbaix 212 JNC TN8400 99 - Oil 6.13 Comments on selected references [1922APP/REI]: [1922APP/REI] determined the solubility of litharge and massicot (PbO(cr)) in 1 M NaOH. From these measurements a log K*so can be calculated using a log |31>3 for Pb(OH)3- of-28.32 (1=1; Section 6.1.3) and a log Kw of -13.79 (1=1). [1922APP/REI] also determined the redox potential of the electrode Pb/PbO(cr) from which [1922APP/REI] calculated the solubility of litharge and massicot. [1928RAN/SPE]: [1928RAN/SPE] determined the solubility of litharge and massicot (PbO(cr)) in NaOH. From these measurements a log K*so can be calculated using a log pli3 for Pb(OH)3- of-21.91 (1=0; Section 6.1.3) and a log Kw of -14.0 (1=0). [1939GAR/VEL]: [1939GAR/VEL] measured Pb(II) solubility in dilute alkaline solutions (Table 6.34). Extrapolation of their measurements to 1=0 with the SIT model gave a b log K S3 of -1.36 (see Figure 6.31) for the reaction PbO(red) + H2O + OH" <=> Pb(OH)3-, instead of -1.34 as calculated by [1939GAR/VEL]. [1939GAR/VEL] calculated also a log K3 of 10.96 value for the reaction Pb(OH)3" + H+ <=> Pb(OH)2° + H2O. The values calculated from [1939GAR/VEL] can be converted to log (3° 12 and log P°ii3 values using a log K°So value of -12.68 for PbO(s, red) (cf. Chapter 6.2.1: PbO(litharge)): 2+ Pb +2H2O Pb(OH)2° + 2H+ log p°1>2 = -17.08 2+ + Pb +3H2O Pb(OH)3- + 3H log J3\3 = -28.04 Further formation constants for Pb(OH)2° and Pb(OH)3- are discussed in Section 6.1.2 and 6.1.3 of this report. Table 6.34: Experimentally determined equilibrium data compiled for the for the reaction PbO(red) + H2O + OH' <=> Pb(OH)3\ Method: sol = solubility measurements. log Kbs3 Reference Comments KM) Medium Method b log K S3: PbO(red) + H2O + OH- <=> Pb(OH)f -1.31 [1939GAR/VEL] T= 298.15 K, 1=0.0 0.0018 NaOH sol -1.42 [1939GAR/VEL] T= 298.15 K, 1=0.0 0.005 NaOH sol -1.25 [1939GAR/VEL] T= 298.15 K, 1=0.01 0.006 NaOH sol -1.32 [1939GAR/VEL] T= 298.15 K, 1=0.01 0.007 NaOH sol 213 JNC TN8400 99-011 Table 6.34: continued -1.35 [1939GAR/VEL] T= 298.15 K, 1=0.01 0.01 NaOH sol -1.36 [1939GAR/VEL] T= 298.15 K, 1=0.01 0.011 NaOH sol -1.33 [1939GAR/VEL] T= 298.15 K, 1=0.01 0.013 NaOH sol -1.28 [1939GAR/VEL] T— 298.15 K, 1=0.02 0.016 NaOH sol -1.32 [1939GAR/VEL] T= 298.15 K, 1=0.02 0.02 NaOH sol -1.42 [1939GAR/VEL] TT— 298.15 K, 1=0.03 0.027 NaOH sol -1.47 [1939GAR/VEL] T= 298.15 K, 1=0.03 0.03 NaOH sol -1.40 [1939GAR/VEL] T= 298.15 K, 1=0.04 0.035 NaOH sol -1.40 [1939GAR/VEL] T= 298.15 K, 1=0.05 0.05 NaOH sol -1.37 [1939GAR/VEL] T= 298.15 K, 1=0.06 0.06 NaOH sol -1.40 [1939GAR/VEL] T= 298.15 K, 1=0.1 0.1 NaOH sol -1.33 [1939GAR/VEL] T= 298.15 K, 1=0.11 0.11 NaOH sol -1.35 [1939GAR/VEL] T= 298.15 K, 1=0.15 0.15 NaOH sol -1.40 [1939GAR/VEL] T= 298.15 K, 1=0.17 0.17 NaOH sol -1.41 [1939GAR/VEL] T= 298.15 K, 1=0.2 0.2 NaOH sol -1.31 [1939GAR/VEL] T= 298.15 K, 1=0.3 0.3 NaOH sol -1.29 [1939GAR/VEL] T= 298.15 K, 1=0.4 0.4 NaOH sol -1.30 [1939GAR/VEL] T= 298.15 K, 1=0.45 0.45 NaOH sol -1.28 [1939GAR/VEL] T= 298.15 K, 1=0.65 0.65 NaOH sol -1.33 [1939GAR/VEL] T= 298.15 K, 1=1.2 1.2 NaOH sol red PbO(s)+H2O+OH" o Pb(OH)3- 05 Lmolal 1 1.5 b Figure 6.31: Plot of log K S3 + 0D vs. Im for the reaction PbO(s) + H2O + OH' <=> Pb(OH)3- at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.05, b log K °S3= - 1.36. Calculated from data compiled in Table 6.34. 214 JNC TN8400 99-011 [1952LAT]: Careful review of the early literature. Solubility of lead chloride calculated from polarographic measurements by [1922GER]. [1960CAR/OLI]: cf. [1960OLI1]; 0.25-10 mM Pb; pKw = 14.18 (1=3) and 13.76 (1=0.3). [1960OLI1]: Olin and co-workers [1960OLI1, 1960OLI2, 1960CAR/OLI, 1961OLI, and 1962PAJ/OLI] determined lead hydrolysis in 0.3 and 3 M perchlorate medium. Their careful experimental work included large range of pH values and lead concentrations. They put up a consistent system of hydrolyzed lead species including Pb2OH3+, 4+ 2 4 Pb4(OH)4 , Pb3(OH)4 +, and Pb6(OH)8 + which was later slightly enlarged by Pb3(OH)5+ [ 1980SYL/BRO]. Pb concentration = 1.25- 80 mM. [1960OLI2]: cf. [1960OLI1]; the value determined at I = 4.5 is corrected in [1962PAJ/OLI]. Pb concentration = 500 - 1490 mM. [1961OLI]: cf. [1960OLI1]; same values as reported by [1960CAR/OLI], [1960OLI1], and [1960OLI2]; pKw = 14.18 (1=3) and 13.76 (1=0.3). Pb concentration = 1 -80 mM. [1962PAJ/OLI]: cf. [1960OLI1]; Pb concentration = 5 - 1450 mM. [1964BAK]: [1964BAK] determined a log K*so for chloropyromorphite and hinsdalite using a equilibration time of several months. A closer examination of his data leads to the conclusion that he probably assumed all phosphate to be present as PO43- (in a Na2HPO4 solution in the case of chloropyromorphite and at a pH of 3 in the case of hinsdalite). These values are therefore not used in this report. [1971NAU/RYZ]: small solubility of laurionite probably due to a transposing of digits (AfG° given as -480.3 instead -408.3) of value which already originally was too low [1987BRU]. [1972NRI]: Nriagu [1972NRI] made measurements in dilute phosphoric acid solution and + determined the constants for the formation of soluble PbH2PO4 and PbHPO4° as well as the solubility of different solid lead phosphates. In 0.1 M solutions he also determined the stability of chloro-, fluoro- and bromopyromorphite [1973NRI1, 1973NRI2]. All constants were corrected to 1=0 with the Davies equation. In [1984NRI] a log K*so for plumbogummite and hinsdalite is estimated. [1972ZIR/YAM]: calculated log P values from electronegativity. [1973BIL/STU]: same values as [1976BIL/HUS]. [1973NRI1]: cf. [1972NRI]. 215 JNC TN8400 99-011 [1973NRI2]: cf. [1972NRI]. [1976BAE/MES]: Careful and extensive review of existing literature concerning hydrolysis of lead. [1976BAE/MES] extrapolated the selected values to I = 0. [1976BIL/HUS]: Bilinski and co-workers [1976BIL/HUS, 1973BIL/STU] determined lead carbonate complex formation constants in 0.1 M KNO3 medium (pK^, used by [1976BIL/HUS] is 13.96) in presence of 0.001-0.01 mM Pb potentiometrically and measured also the solubility of lead carbonate complexes. They also recalculated the data of [1965BAR/BAR] in 1 M NaC104. The log (3 value given in 0.7 NaClO4 is based on a private communication from M. Sipos (310"9 M Pb). The first and second lead hydrolysis constants determined by [1976BIL/HUS] were not chosen in this report as [1976BIL/HUS] neglected the formation of the polynuclear species Pb3(OH)4 at a Pb concentration of 10~5 M. [1976SMI/MAR]: Calculated mean of values at a given ionic strength: The values are selected after a critical review, however, no comments to the individual paper selected or not selected are made by Smith and Martell. Several compilations have been published: [1982SMI/MAR] and [1989SMI/MAR]. [1979PAT/OBR]: [1979PAT/OBR] selected thermodynamic data for lead carbonate and hydroxide complexes and solids and then tested their dataset with experimental data from different sources. [1980CLE/JOH]: [1980CLE/JOH] made a careful and extensive review of data concerning the solubility of sparingly soluble lead salts. Based on this review, log K*so for many salts are recommended. In cases were no newer data could be found, the data recommended by [1980CLE/JOH] are used. [1980KAW/ISH]: [1980KAW/ISH] determined with a potentiometric method the hydrolysis of Pb(II) in rather concentrated solutions (Pb(II) =1-80 mM). [1980KAW/ISH] + themselves classified their log P values for PbOH and Pb(OH)2 as doubtful, as only + minor concentration of PbOH and Pb(OH)2 are present under the experimental conditions . p^ in 3 M LiClO4, as determined in [1980KAW/ISH] = 13.87. [1980MAN/DEU]: selected data (from the literature) for Pb oxide/hydroxide, chloride, sulfate and carbonate complexes and solids have been corrected to 1=0 by [1980MAN/DEU]. [1980SCH]: [1980SCH] selected formation constants from the literature and verified his dataset with experimental data from different sources in systems containing considerable amounts of carbonate; pK^, used by [1980SCH] is 13.99. 216 JNC TN8400 99-011 [1980SYL/BRO]: Their determination of lead hydrolysis was carried out in 0.1 M KNO3 at 0.1-2 raM Pb concentration. As these values were determined in nitrate solutions they were not chosen in this report for extrapolation to I = 0. [1980SYL/BRO] introduced an additional lead species (Pb3(OH)5+) and recalculated the values determined by Olin and co-workers at I = 0.3 and 3, including the new species Pb3(OH)5+. [1980SYL/BRO] also showed that the additional consideration of Pb3(OH)5+ had hardly an influence on the other log P values. [1980WAL/SIN]: [1980WAL/SIN] determined the solubility of cerrusite, anglesite and lead hydroxide/oxide as well as constants for lead hydrolysis in a wastewater containing predominantly H2SO4. The values given are the mean calculated from measurements in wastewater. As also additional anionic components were present in the wastewater used which were not considered in their calculations, the calculated log K*so values are too large and the complex formation constants too small. Values from [1980WAL/SIN] were not chosen for the calculation. Nevertheless, they can be used for comparison. [1982BIL/SCH]: cf. [1976BIL/HUS]; pK^ used = 13.96 (for 1=0.1) and 13.76 (for 1=0.3). [1982SMI/MAR]: Calculated mean of values at a given ionic strength. See also [1976SMI/MAR] and [1989SMI/MAR]. [1983SCH/GAR]: Thermodynamic data selected in a review, dataset tested with experimental data from different sources. The constant for PbHCO3+ could not be verified under the conditions of the experiments. The value given for hydrocerrusite is an approximate value. [1983SCH/GAR] criticizes the older value for hydrocerrusite as too high compared to solubility measurements. [1984NRI]: cf. [1972NRI]. [1984TAY/LOP]: [1984TAY/LOP] determined with X-ray analysis the conditions under which massicot (or litharge), plumbonacrite, hydrocerrusite and cerrusite were predominantly formed in aqueous carbonate solutions (I, pCO2, pH were varied) and extrapolated their results to 1=0. They could observe the formation of all these substances at room temperature. Calculations carried out by [1984TAY/LOP] demonstrated that at a total C concentration of 10"4 M, the solubilities of plumbonacrite, cerrusite and hydrocerrusite are similar, especially near neutral pH values. [1984UHL/HEL]: [1984UHL/HEL] determined the log K*so of galena (PbS(cr)) with help of competitive complex formation with EDTA and TRIS. Extrapolation of their values to I =0 with SIT yields a value of 12.17. [1984UHL/HEL] recalculated also the values given by [1953HEM] and [1979GIO/BAR] for Pb(HS)2° and Pb(HS)3-. However, from the information given in [1984UHL/HEL], their calculations are not 217 JNC TN8400 99 -Oil traceable and the constants used by [1984UHL/HEL] did not agree in all cases with the original references. [1987BRU]: in the case of lead carbonate complexes the influence of I is not taken into account. [1988BYR/KUM]: extrapolated to 1=0.7 for seawater. The values for the lead carbonate complexes are a recalculation of [1981BYR]. These values are probably less precise than measurements e.g. by [1977SIP/VAL]. [1989DOR/MAR]: These titration experiments are unfortunately not well documented. The original I (probably the measurements were carried out in diluted suspensions) was not reported. However, the data were extrapolated to 1=0 by the authors. Their value determined for the first lead hydrolysis constant is not chosen in our calculations as the formation of polynuclear lead hydroxide species can not be excluded at a Pb concentration of 0.02-0.2 mM [1989DOR/MAR]. Their values given for the formation of PbCO3° and PbHCCV complexes and for the solubility of cerrusite seem to be more accurate. [1989SMI/MAR]: Calculated mean of values at a given ionic strength. See also [1976SMI/MAR] and [1982SMI/MAR]. [1993MAC/PAG]: [1993MAC/PAG] compare lead solubility calculated with thermodynamic data with experimental lead solubility. [1995MAR/MAC]: [1995MAR/MAC] tried to verify the MINTEQA2 database with solubility measurements under different conditions and identified the precipitated solids with X-ray diffraction. They observed the precipitation of cerrusite, hydrocerrusite and anglesite at room temperature and showed that neither lanarkite (Pb2OSO4) nor crystalline Pb(OH)2 precipitates from solutions. In absence of lanarkite and Pb(OH)2 their calculations showed a reasonable agreement with the experimental observations. 218 JNC TN8400 99-011 7 Bismuth The most stable oxidation state of bismuth is +III, but the oxidation states -III and -I can be prepared in liquid nitrogen or in the gas phase. Bismuth compounds of the +V state are strong oxidizing agents and are able to oxidize water to oxygen [1985LOV/MEK, 1995WIB]. In aqueous environments, Bi(III) is a strong acid and Bi3+ starts to hydrolyze at a pH of about 1. Bismuth has a strong tendency to form polynuclear complexes [1986BAE/MES]. For the present evaluation, thermodynamic data from literature are compiled for the formation of Bi complexes and solids with hydroxide, chloride, fluoride, carbonate, nitrate, phosphate, and sulfate. Based on experimental data reported in the literature, thermodynamic data are recommended for the complex formation of bismuth with hydroxide, chloride, nitrate and for the redox equilibria Bi3+ + 3e- <=> Bi(cr). Thermodynamic data are also selected for the formation of the solids a-Bi2O3(cr), BiOCl(s), (BiO)4(OH)2CO3(cr), (BiO)2CO3(cr), and BiONO3(s). For other potentially relevant solid Bi phases, like BiPO^s) and Bi2(SO4)3(s), no experimentally determined solubility data are available. 7.1 Hydrolysis of bismuth The Bi3+ ion starts to hydrolyze at a pH of about 1 and has the tendency to form polynuclear complexes at Bi concentrations > 1(H M under acidic and neutral conditions [1976BAE/MES]. The solubility of Bi2O3 does not depend on pH in the pH range of ~ 8 - 12, indicating the presence of Bi(OH)3°. Only in alkaline solutions (pH > 12), bismuth solubility increases again. The data used for the evaluation of equilibrium constants for the hydrolysis of bismuth(III) are given in Table 7.1 and in Figure 7.1 to Figure 7.8. Additional equilibrium data that were not selected for the calculation of the bismuth(III) hydrolysis are compiled in Tables 7.2 and 7.3. Olin [1957OLI, 1959OLI, 1961OLI, 1975OLI] was one of the first investigators who determined the hydrolysis of Bi(III). Based on potentiometric measurements he reported for the formation of mononuclear BiOH2+ a log Pu value of -1.58 and for the polynuclear Bi6(OH)i26+ a log 06,12 value of 0.33 in 3 M NaClO4. His careful experimental work encompassed a large range of pH values and bismuth concentrations and he developed a consistent dataset for the hydrolysis of bismuth. [1976BAE/MES] critically reviewed these measurements and extrapolated the measurements of Olin, [1960TOB, 197 IBID, 1972DRA/NIM2] to 1=0. These measurements, as well as the studies of bismuth hydrolysis by [1972DRA/NIM1, 1982SUG/SHI, 1985SED/SIM] in perchlorate medium were used in this report for the evaluation of bismuth hydrolysis at 1=0. The measurements of [1975ANT/NEV, 1975HEI/SCH, 1982HAT/SUG, 1987SUG/SHI, 1993KRA/DEC] in nitrate, chloride or sulfate medium were not chosen in this report for the calculations of the hydrolysis constants, because bismuth forms complexes with these anions. Also the data of [1960TOB] were 219 JNC TN8400 99-011 excluded because he neglected the possible formation of other Bi hydroxide complexes 2+ 7+ (BiOH , Bi9(OH)2o ) in his data interpretation. Table 7.1: Experimentally determined equilibrium data compiled for the bismuth hydroxide 3+ 3 n + system, according to the equilibrium: mBi + nH2O <=> Bim(OH)n " + nH . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.1:2: 'Comments on selected references'. Method: extr = solvent extraction, mig = migration, pot = potentiometry, sol = solubility measurements, sp = spectrophotometry, Reference Comments Medium Method -"m,n 3+ log pu: Bi + H2O -1.43 [1971BID] T= 298.15 K, 1=0.1 0.1 NaC104 extr -1.55 [1972DRA/NIM1] T= 298.15 K, 1=1 1 H, NaClO4 sp -1.59 [1975OLI] T= 298.15 K, 1=3 3 NaC104 pot -1.84 [1982HAT/SUG] T= 298.15 K, 1=1 1 NaC104 extr -1.46 [1985SED/SIM] T= 293.15 K, 1=0.4 0.4 NaC104 extr -1.40 [1987MIL/ROE] T= 298 K, 1=0.25 0.25 NaC104 mig 3+ + log pli2: Bi + 2H2O <=> Bi(0H)2 + 2H+ -2.82 [1972DRA/NIM1]T= 298.15 K, 1=1 1 H, NaC104 sp -4.74 [1982HAT/SUG] T= 298.15 K, 1=1 1 NaClO4 extr -3.36 [1985SED/SIM] T= 293.15 K, 1=0.4 0.4 NaClO4 extr -3.57 [1987MIL/ROE] T= 298 K, 1=0.25 0.25 NaClO4 mig 3+ log p]f3: Bi + 3H2O & Bi(0H)3° + 3H+ -7.5 [1982HAT/SUG] T=298 .15 K, 1=1 1 NaClO4 extr -6.41 [1987MTL/ROE] T=298 K, 1=0.25 0.25 NaClO4 mig log Kjt4: Bi(OH)3° + H2O&.Bi(OH)4- + -12.84 [1971BID] T=298 .15 K, 1=1 1 NaC104 sol -13.07 [1987ME7ROE] T=298 K, 1=0.25 0.25 NaClO4 mig 220 JNC TN8400 99-011 Table 7.1: continued 3+ + log P6,J2-- 6Bi + 12H2O a Bi6(OH)126+ + 12H 0.26 [1972DRA/NIM1] T= 298.15 K, 1=1 1 H, NaC104 sp 0.33 [1975OLI] T= 298.15 K, 1=3 3 NaC104 pot 6+ 7+ + log K9i20:1.5 Bi6(OH)]2 + 2 H2O <^> Bi9(OH)20 •2H -3.5 [1959OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot -3.9 [1972DRA/NIM2] T= 298.15 K, 1=0.1 0.1 NH4C1O4 sp 7+ 6+ + log K9i2i: Bi9(OH)2o + H2O <=> Bi9(OH)2i + H -3.2 [1959OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot -3.2 [1972DRA7NIM2] T= 298.15 K, 1=0.1 0.1 NH4C104 sp 6+ 5+ + log K9i22: Bi9(OH)2j + H2O <=> Bi9(OH)22 + H -2.6 [1959OLI] T= 298.15 K, 1=0.1 0.1 NaC104 pot -2.8 [1972DRA/NIM2] T= 298.15 K, 1=0.1 0.1 NH4C1O4 sp 3+ 5+ + log p3A: 3Bi + 4H2O <=> Bi3(OH)4 + 4H -0.80 [1975OLI] T= 298.15 K, 1=3 NaCIO, pot 221 JNC TN8400 99 -Oil 7.1.1 Based on potentiometric measurements, [1957OLI] determined a log BJJ value of-1.58 for the 2+ 3+ 2+ + formation of mononuclear BiOH according to the reaction Bi + H2O o BiOH + H (See Table 7.2). Bidleman [1971BID] measured in 0.1 M NH4CIO4 a log Bu of -1.43. [1976BAE/MES] extrapolated in his careful review from these two values a log B0^ of -1.09 for 1=0 (Table 7.3). In the present report, the data determined experimentally by [1971BID, 1972DRA/NIM1, 1975OLI, 1982HAT/SUG, 1985SED/SIM, 1987MH7ROE] (Table 7.1) were used for evaluation of log B°ii. Extrapolation to I = 0 with the SIT term is shown in Figure 7.1. 3+ 2+ Bi + H2O <=> BiOH log = -0.92, Ae =-0.09 The data measured by [1975ANT/NEV] and [1982HAT/SUG] in nitrate medium and [1987SUG/SHI] in chloride medium (Table 7.2) are not used for the determination of B°1;1 values because both, nitrate and chloride, tend to form complexes with bismuth. 3+ 2+ + Bi +H?O<^BiOH +H 1 2 3 L molal 3+ 2+ Figure 7.1: Plot of log Bij + 4 D vs. Im for the reaction : Bi + H2O <=> BiOH + H+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.09; log B°! i = - 0.92. Calculated from data compiled in Table 7.1. 222 JNC TN8400 99- Oil + 7.1.2 Bi(OH)2 The determination of the precise value of log (3 j ^ is difficult as in most experimental system compiled in this report, Bi(0H)2+ is only a minor species and BiOH2+ or polynuclear Big(OH)i26+ dominate the speciation in the solutions (see also [1976BAE/MES]). Thus, the measured values given in Table 7.1 [1972DRA/NIM1, 1982HAT/SUG, 1985SED/SIM, 1987MIL/ROE] show a considerable spread. [1976BAE/MES] estimated in their work (due to lack of experimental data) a value of -4.0 for log (3° 1,2- Extrapolation of the data given in Table 7.1 to 1=0 gives: 3 + Bi + + 2H2O Bi(OH)2 logP°i>2 = -2.56 This value is similar to the value calculated by [1982LAP/KOL] from measurement at 75 - 300 °C and the determinations by [1975ANT/NEV] (Table 7.2). The uncertainty of log pi,2 is much larger than for the other bismuth hydroxide species. However, even at bismuth concentrations of 10"11 M, the dominating species in acidic solutions are Bi3+, BiOH2+ and + Bi(OH)3° and not the Bi(OH)2 complex [1987SUG/SHI]. It should be mentioned that the formation constant given in [1952LAT] for BiO+ (corresponding + to Bi(OH)2 ) (Table 7.3) originates from the early works of [1923SWI, 1923SMI, + 1947GRA/SIL]. [1957OLI] could not establish the presence of Bi(OH)2 but demonstrated that the polynuclear Bi6(OH)i26+ complex is stable compared to Bi(OH)2+. [1957OLI] pointed out that as polynuclear Bi6(OH)i26+ complex is so stable that the equilibrium concentration of Bi(OH)2+ is too low to be detected in his experiments. 3+ + + Bi + 2H2O & Bi(OH)2 + 2H 3+ + Figure 7.2: Plot of log p1>2 + 6 D vs. Im for the reaction : Bi + 2H2O <=> Bi(OH)2 + 2H+ at 25 °C. The straight line shows the result of the linear regression: Ae = —0.03; log P°i,2 = -2.56. Calculated from data compiled in Table 7.1. 223 JNC TN8400 99-011 7.1.3 Bi(OH)3° [1971BID] measured in 0.1 and 1 M perchlorate solution the bismuth hydrolysis with an organic extractant (dithizone). He was able to show that no (or only little) polynuclear species were present in 0.001-0.01 mM bismuth solutions in the pH range 9.9 - 11.3. [1971BID] determined in 0.1 M NaC104 a log pli3 of-9.43 for the formation of Bi(OH)3 (Table 7.2). [1976BAE/MES] extrapolated from this a log [313 value of -8.86 for 1=0. However, newer measurements in perchlorate medium by [1982HAT/SUG, 1985SED/SIM, 1987MIL/ROE] indicate a less negative log Pi 3 value. [1982HAT/SUG] assumed that the difference between their results and the result of [1972BID] may be due to the formation of polynuclear species in the experiments of [1972BID] at Bi(III) concentrations of 0.01 mM. Also the measurements of [1975ANT/NEV] in KN03 indicate a log p]i3 value near -5.5 (Table 7.2). Extrapolation of the measurements of [1982HAT/SUG] and [1987MEL/ROE] to 1=0 results in: 3 Bi + + 3H2O Bi(OH)3° log p°li3 = -5.31 1 2 lm, molal 3+ Figure 7.3: Plot of log p1>3 + 6 D vs. Im for the reaction : Bi + 3H2O «• Bi(OH)3° + 3H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = 0.91; log P°i 3 = - 5.31. Calculated from data compiled in Table 7.1. 224 JNC TN8400 99-011 7.1.4 Bi(OH)4- For Bi(OH)4-, [1973BID] measured in 0.1 M NH4CIO4 a log p]>4 of -22.26 (Table 7.2). The data reported in the literature for log (314 are quite different (see Table 7.2), while for the + reaction Bi(OH)3° + H2O <=> Bi(OH)4- + H the variation between the different experimental papers is smaller (Table 7.1), as these values do not depend on the value of log Pi 3 (cf. Section 7.1.3). Extrapolation of the log Kj>4 values determined by [1971BID] and [1987MH7ROE] gives a log Ki>4 of -13.40 for the reaction Bi(OH)3° + H2O <=> Bi(OH)4- + H+. Using a log P°i,3 of-5.31 (Section 7.13) one obtains: 3+ Bi + 4H2O Bi(OH)4- logp°1|4 = -18.71 + Bi(OH)3°+ H20 o Bi(OH)4"+H -10 T -10.5 •• -11 CM -11.5 y = 0.14x-13.40 I -12 + s? -12.5 • O -13 -13.5 -: -14- -14.5 •• -15 0 0.5 1 1.5 lmi molal Figure 7.4: Plot of log Ki,4 + 4 D vs. Im for the reaction : Bi(0H)3° + H2O «=> Bi(0H)4- + H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = - 0.14; log K°M = -13.40. Using a log p°li3 of -5.31 one obtains a log p°M of - 3+ 18.71 for the reaction Bi + 4H2O « Bi(0H)4- + 4H+. Calculated from data compiled in Table 7.1. 225 JNC TN8400 99-011 6+ 7.7.5 Bi6(OH)]2 Polynuclear Big(OH)i26+ dominates the speciation below pH 3 in solution containing more than 0.01 mM Bi. Experimental values are compiled in Table 7.1. Dragulescu and co-workers [1972DRA/NIM1, 1972DRA/NIM2] determined bismuth hydrolysis in 0.1 and 1 M perchlorate medium. Their results are comparable to the results of Olin. Extrapolation of the values determined by [1972DRA/NIM1] and [1975OLI] to 1=0 is shown in Figure 7.5 and results in: 3+ 6+ 6Bi + 12H2O Bi6(OH)12 + 12H+ log (3°6>i2 = 1-34 The value determined by [1960TOB] was not chosen for the extrapolation, as [1960TOB] did 2+ + not include the species BiOH and Bi(OH)2 in his calculations. 6Bi3++12H O<=>Bi (OH) 6++12H+ A 2 6 12 *r 3.5 3 • Q CD 2.5 + 2 CM CO* 1.5 • CO. 1 lo g 0.5 y - 6.0"4x 4-1.34 0 -0.5 .1 1 1 1 1 2 lm, molal 3+ 6+ Figure 7.5: Plot of log p6,i2 + 6 D vs. Im for the reaction : 6Bi + 12H2O <=> Bi6(OH)i2 + 12H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = - 0.04; log p°6,i2 = 1.34. Calculated from data compiled in Table 7.1. 226 JNC TN8400 99-011 7+ 6+ 5+ 7.1.6 Bi9(OH)20 , Bi9(OH)2i , and Bi9(OH)22 Based on potentiometric measurements, [1959OLI] reported values for the formation of the 7+ 6+ 5+ 6+ polynuclear Bi9(OH)20 , Bi9(OH)2i , and Bi9(OH)22 complexes from Bi6(OH)i2 above a pH of ~ 3 and at Bi concentrations of 0.25 - 4 mM. [1960TOB] proposed the existence of a Bi6(OH)i53+ complex. Spectrophotometric measurements made by [1972DRA/NIM1] are 7+ 6+ 5+ consistent with the formation of polynuclear Big(OH)2o , Bi9(OH)2i , and Bi9(OH)22 complexes as proposed by [1959OLI]. The consecutive formation constants given by [1959OLI] and [1972DRA/NEVI1] at 1=0.1 are listed in Table 7.1. The mean of these values was corrected in this report to 1=0 using the SIT approximation assuming a Ae of 0: 6+ 7+ 1.5Bi6(OH)i2 +2H2O «• Bi9(OH)20 + 2H+ log K°9>20 =-3.37 7+ 6+ Bi9(OH)20 + H2O a Bi9(OH)21 + H+ log K\2! = -1.89 6+ 5+ Bi9(OH)2i + H2O <=> Bi9(OH)22 +H+ log K\22 = -1.61 3+ These values were then converted to log (3 values that refer to Bi , using a log P°6,j2 of 1.34 (see Figure 7.5): 3+ 7+ 9 Bi + 20 H2O <=> Bi9(OH)20 + 20H+ log (39,2o = -1-36 3+ 6+ 9 Bi + 21 H2O <=> Bi9(OH)2i + 21H+ log (39>21 = -3.25 3+ 5+ 9 Bi + 22 H2O <=* Bi9(OH)22 + 22H+ log fj9>22 = -4.86 5+ 7.1.7 Bi3(OH)4 [1975OLI] calculated a tentative log (33i4 value of -0.80 in 3 M NaClO4 solution for the 3+ 5+ reaction 3Bi + 4H2O <=> Bi3(OH)4 + 4H+. As this is the only determination of this constant it is extrapolated to 1=0 with SIT, log p\4 - 2D = P°3?4 - Ae Im, assuming a Ae of 5+ 3+ 0.18 (from Ae(Al3(OH)4 ) = 1.30, Ae(Fe ) = 0.56 and Ae(H+) = 0.14; [1992GRE/FUG]): 3+ 5+ 3 Bi + 4 H2O <=> Bi3(OH)4 + 4H+ log (3°3>4 = -0.80 3+ 7.1.8 Bi6(OH)!5 [1960TOB] proposed, based on titration experiments, the existence of a Bi6(OH)]53+ complex 7+ 6+ 5+ instead of Bi9(OH)20 , Bi9(OH)2i , and Bi9(OH)22 as proposed by [1959OLI] (see Table 7.2). Spectrophotometric measurements made by [1972DRA/NIM1], however, are consistent 7+ 6+ 5+ with the formation of polynuclear Bi9(OH)2o , Bi9(OH)2j , and Bi9(OH)22 complexes as proposed by [1959OLI]. Thus, the existence of a Bi6(OH)i53+ complex instead of 7+ 6+ 5+ Bi9(OH)2o , Bi9(OH)2] , and Bi9(OH)22 complexes seems rather unlikely. 227 JNC TN8400 99-011 7.1.9 Additional equilibrium data compiled for the bismuth hydroxide system Table 7.2: Additional experimentally determined equilibrium data compiled for the bismuth hydroxide 3+ 3-n + system, according to the equilibrium: mBi + nH2O <=> Bim(OH)n + nH . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references see text and Section 7.12. Method: extr = solvent extraction, pot = potentiometry, mig = migration, sol = solubility measurements, sp = spectrophotometry, tit = titration (pH). log Pm,n Reference Comments KM, Medium Method 3 24 + log pu: Bi + •f //26> <=> fi/Otf " + H -1.58 ' [1957OLI] T= 298.15 K, 1=3 3 NaClO4 pot -1.58 ' [1961OLI] T= 298.15 K, 1=0.1,3 3 NaC104 pot -1.552 [1975ANT/NEV] T= 298.15 K, 1=0.1 0.1 KNO3 sp -1.50 2 [1975ANT/NEV] T= 298.15 K, 1=0.3 0.3 KNO3 sp -1.41 2 [1975ANT/NEV] T= 298.15 K, 1=0.5 0.5 KNO3 sp -1.342 [1975ANT/NEV] T= 298.15 K, 1=1 1 KNO3 sp 2 -1.67 [1982HAT/SUG] T= 298.15 K, 1=1 1 NaNO3 extr -2.72 2 n987SUG/SHIl T= 298.15 K, 1=1 1 NaCl extr 3+ + + log j3u: Bi -v 2H2O « Bi(0H)2 + 2H -3.52 2 [1975ANT/NEV] T= 298.15 K, 1=0.1 0.1 KNO3 sp -3.45 2 [1975ANT/NEV] T= 298.15 K, 1=0.3 0.3 KNO3 sp -3.26 2 [1975ANT/NEV] T= 298.15 K, 1=0.5 0.5 KNO3 sp -3.10 2 [1975ANT/NEV] T= 298.15 K, 1=1 1 KNO3 sp -0.15 2 [1975HEI/SCH] T= 298.15 K, 1=1.5 1.5 H2SO4 pot 2 -4.51 [1982HAT/SUG] T= 298.15 K, 1=1 1 NaNO3 extr 3 -2.29 [1982LAP/KOL] T=298.15K, I=dil NaOH, HC1O4 sol -9.72 [1987SUG/SHI] T= 298.15 K, 1=1 1 NaCl extr -4.00 2.4 [1993KRA/DEC] T= 29610.5 K, 1=1 1 NO3- sol 4 -4.10 [1993KRA/DEC] T= 29610.5 K, 1=1 1 cio4- sol 3+ + log Pu: Bi + 3H2O <=> Bi(OH)3° + 3H -9.43 [1971BID] T= 298.15 K, 1=0.1 0.1 NH4CIO4 sol -5.94 2 [1975ANT/NEV] T= 298.15 K, 1=0.1 0.1 KNO3 sp -5.93 2 [1975ANT/NEV] T= 298.15 K, 1=0.3 0.3 KNO3 sp -5.58 2 [1975ANT/NEV] T= 298.15 K, 1=0.5 0.5 KNO3 sp -5.29 2 [1975ANT/NEV] T= 298.15 K, 1=1 1 KNO3 sp 2 -7.6 [1982HAT/SUG] T= 298.15 K, 1=1 1 NaNO3 extr 3 -9.02 [1982LAP/KOL] T= 298.15 K,I=dil NaOH, HC1O4 sol 5 -5.59 [1985SED/SIM] T= 293.15 K, 1=0.4 0.4 NaC104 extr -10.7 2 [1987SUG/SHI] T= 298.15 K, 1=1 1 NaCl extr 4 -9.90 [1993KRA/DEC] T= 29610.5 K, 1=1 1 cio4- sol -10.00 2.4 [1993KRA/DEC] T= 29610.5 K, 1=1 1 NO," sol 228 JNC TN8400 99-011 Table 7.2: continued 3+ + log PL4: Bi + 4H2O <=> Bi(OH)4- + 4H 6 -19.48 [1987MDL/ROE] T= 298 K, 1=0.25 0.25 NaClO4 mig 6 -22.26 [197 IBID] T= 298.15 K, 1=0.1 0.1 NH4CIO4 sol 4 -21.50 [1993KRA/DEC] T= 296±0.5 K, 1=1 1 C1O4" sol -21.50 2- 4 [1993KRA/DEC1 T= 296+0.5 K, 1=1 1 NCV sol + log K1A: Bi(OH)3°+ H2O<=>Bi(OH)4-+ H 4 -11.60 [1993KRA/DEC] T= 298.15 K, 1=1 1 cio4- sol -11.70 2 |"1993KRA/DEC1 T= 298.15 K, 1=1 1 NO,~ sol 3+ 6+ + 2-- 6Bi + 12H2O <=> Bi6(OH)]2 + 12H 0.33 ' [1957OLI] T= 298.15 K, 1=3 3 NaClO4 pot 0.33 ' [1961OLI] T= 298.15 K, 1=3 3 NaClO4 pot -0.53 [1960TOB1 T= 298.15 K, 1=1 1 NaClO4 tit 6+ 7+ + log K9,20:1.5 Bi6(OH),2 + 2 H2O <=> Bi9(OH)20 + 2H -3.5 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot 7+ 6+ + log K9i21: Bi9(OH)20 + H2O <=> Bi9(OH)21 + H -3.2 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO4 pot 6+ 5+ + log K9a2: Bi9(OH)2l + H2O <=> Bi9(OH)22 + H -2.6 ' [1961OLI] T= 298.15 K, 1=0.1 0.1 NaClO. pot 3+ 3+ + log P6jJ5: 6Bi + 15H2O Bi6(OH)15 + 15H -8.63 [1960TOB] T= 298.15 K, 1=1 NaClO, tit 1 same values as [1959OLI] formation of nitrate, chloride or sulfate complexes with electrolyte possible, extrapolated from measurements at 75-300 °C. 4 Bi: 0.01 - 1000 mM. Formation of polynuclear species probable 5 Bi: <0.1 mM. Formation of polynuclear species possible 6 log K14 values are given in Table 7.1 7 existence of Bi6(OH)15questionable, see Section 7.1.8 229 JNC TN8400 99-011 Table 7.3: Thermodynamic data for the bismuth hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. l°g P Reference Comments I (M) Medium 3+ 2+ + log Pu: Bi + H2O <=> BiOH + H -1.09 [1976BAE/MES] T= 298.15 K, 1=0 0 -1.10 [1976SMI/MAR] T= 298.15 K, 1=0 0 -1.09 [1981BAE/MES] T= 298.15 K, 1=0 0 -1.4 [1982WAG/EVA] T= 298.15 K, 1=0 0 -0.91 ' [1985BAB/MAT] T= 298.15 K, 1=0 0 -1.58 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -1.58 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -1.10 [1989SMI/MAR] T= 298.15 K, 1=0 0 -1.40 [1993KRA/DEC] T= 296±0.5 K, 1=1 1 NO3" -1.40 n993KRA/DECl T= 296±0.5 K, 1=1 1 C1O, 3 + + log fih2: Bi * 2H2O Bi(OH)2 + 2H -0.1 [1952LAT] T= 298 K, I=diluted 0 -4.00 [1976BAE/MES] T= 298.15 K, 1=0 0 -1.4 ri982WAG/EVA1 T= 298.15 K, 1=0 0 3+ + log pu: Bi + 3H2O <=> Bi(OH)3° + 3H -8.86 [1976BAE/MES] T= 298.15 K, 1=0 0 -8.90 [1976SMI/MAR] T= 298.15 K, 1=0 0 -8.86 [1981BAE/MES] T= 298.15 K, 1=0 0 -9.00 [1989SMI/MAR] T= 298.15 K, 1=0 0 3+ + log PK4: Bi -(- 4H2O « Bi(OH)4~ + 4H -21.80 [1976BAE/MES] T= 298.15 K, 1=0 0 -21.80 [1976SMI/MAR] T= 298.15 K, 1=0 0 -21.20 [1989SMI/MAR1 T= 298.15 K, 1=0 0 3 f 6+ + logP61,-6Bi - + 12H2O <=> Bi6(OH)12 + 12H -0.53 [1976SMI/MAR] T= 298.15 K, 1=1 1 0.32 [1982WAG/EVA] T= 298.15 K, 1=0 0 -0.31 l [1985LOV/MEK] T= 298.15 K, 1=0 0 0.33 ' [1985LOV/MEK] T= 298.15 K, 1=0 0 -0.18 [1989SMI/MAR] T= 298.15 K, 1=1 1 0.30 [1993KRA/DEC] T= 296+0.5 K, 1=1 1 NO3- 0.30 [1993KRA/DEC] T= 29610.5 K, 1=1 1 cicv 230 JNC TN8400 99-011 Table 7.3: continued 3+ 7+ + log Pi>.20: 9Bi + 20H2O <=> Bi9(OH)20 + 20H -2.61 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1 -3.00 [1982WAG/EVA] T= 298.15 K, 1=0 0 -3.95 1 [1985LOV/MEK] T= 298.15 K, 1=0 0 -3.01 1 [1985LOV/MEK] T= 298.15 K, 1=0 0 -2.61 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 3+ 6+ + log A>,2/.- 9Bi + 21H2O <=> Bi9(OH)21 + 21H -5.79 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1 -6.21 [1982WAG/EVA] T= 298.15 K, 1=0 0 -7.16 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -6.21 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -5.79 [1989SMI/MAR] T= 298.15 K, 1=0.1 0.1 3+ 5+ + log P9,2i- 9Bi + 22H2O <=> Bi9(OH)22 + 22H -8.47 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1 -8.84 [1982WAG/EVA] T= 298.15 K, 1=0 0 -9.76 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -8.81 2 [1985LOV/MEK] T= 298.15 K, 1=0 0 -8.47 ri989SMI/MAR1 T= 298.15 K, 1=0.1 01 3+ 3+ log p6J5: 6Bi + 15H2O <=> Bi6(OH)!5 + 15H+ -8.6 ri982WAG/EVAT T= 298.15 K, 1=0 0_ 1 calculated with a Afi° of -95.55 kJ/mol for Bi2+ (Section 7.11). 2 numbers given by [1985LOV/MEK] in the text and in the table are different 231 JNC TN8400 99-011 7.2 Solid bismuth-oxide/hydroxide 7.2.1 a-Bi2O3(cr) The low temperature modification of solid bismuth oxide is the monoclinic a-Bi2C>3 which is stable up to 715° C where it undergoes a polymorphic transformation to cubic P-Bi2C>3 (or 5- Bi2O3) [1943SCH/RIT, 1981KAL/BOR, 1995WIB]. Unfortunately, in some papers and reviews no description of the exact nature of the Bi2O3 used is given. However, [1943SCH/RIT] and [197IBID] state explicitly that they used a-Bi2O3. [1943SCH/RIT] reported solubility measurements in 0.5 to 2.5 M NaOH which they interpret in terms of one 3 equilibrium V2 a-Bi2O3(cr) + /2 H2O + OH- <=> Bi(OH)4- (Table 7.4). [1971BID] made measurements in 1 M NaClO4 and reported a log Kbs4 = -4.39, which are in close agreement to the data given by [1943SCH/RIT] (see Table 7.4 and Figure 7.6). b Extrapolation of the data of [1943SCH/RIT] to I = 0 with the SIT term gives a log K °S4 = -4.28 as shown in Figure 7.6. The extrapolation of the data of [197 IBID] to I = 0 with a Ae of 0.03 (Figure 7.6) resulted in a log Kb° 34= -4.39. From the mean of these two values (log b K °S4 = -4.33) and log p°])4 = -18.71, a log Kso*° can be calculated: 3+ 2Bi +3H2O <=> a-Bi2O3(cr) + 6H+ log K*°so = -0-76 The value of this constant strongly depends from log P°];4 value. Thus the values given in different compilations can differ strongly (Table 7.6). Based on log K*°so = -0.76 and a AfG° 3+ of 95.55 kJ/mol for Bi (Section 7.11), a AfG° of -515.99 kJ/mol is obtained for oc-Bi203(cr). Table 7.4: Experimentally determined equilibrium data compiled for the dissolution of bismuth oxide. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: sol = solubility measurements. log KSo Reference Comments KM, Medium Method - log KSo: 0.5 a-Bi2O3(cr)+ OH~ <=> Bi(OH)4 -4.27 i [1943SCH/RIT] T= 298.15 K, 1=0.5 0.5 NaOH sol -4.28 1 [1943SCH/RIT] T= 298.15 K, 1=0.99 1 NaOH sol -4.31 1 [1943SCH/RIT] T= 298.15 K, 1=1.41 1.4 NaOH sol -4.30 ] [1943SCH/RIT] T= 298.15 K, 1=1.97 1.97 NaOH sol -4.29 » [1943SCH/RIT] T= 298.15 K, 1=2.46 2.46 NaOH sol -4.39 [197 IBID] T= 298.15 K, 1=1 1 NaCIO, sol Table III of [1943SCH/RIT] contains an error in the exponent of K). 232 JNC TN8400 99-011 0.5 a-Bi2^?03 -2 -2.5 -3 O -3.5 [1943SCH/RIT] + -4 y = -0.03x - 4.28 o o CO -4.33 JO -5 -5.5 {1971 BID] -6 -6.5 -7 0 12 3 !mi molal b Figure 7.6: Plot of log K S4 + 0 D vs. Im for the reaction 0.5 a-Bi2O3(cr) + OH- <=> b Bi(OH)4-at 25 °C, calculated from data compiled in Table 7.4 (K S4 denotes the solubility product of a-Bi2O3(cr) in equilibrium with the hydroxide OH~ concentration. Details of notation are given in Table 2.2 of this report). The straight line shows the result of the linear regression of the experimental data of b [1943SCH/RIT]: Ae = 0.03; log K °S4 = - 4.28. Extrapolation of the data of b b [1971BID] to 1=0 gives a log K °S4 = -4.39, giving a mean log K °S4 of-4.33. 7.2.2 Precipitated Bi(OH)3(s) [1993KRA/DEC] give for freshly precipitated Bi(OH)3(s) a log K*S3 of -4.70 for the reaction Bi(OH)3 <=> Bi(OH)3(s) (Table 7.5), indicating a slightly higher solubility of Bi(OH)3(s) in comparison to well crystallized oc-Bi203(cr). 7.2.3 Additional data compiled for solid bismuth oxides/hydroxides Data calculated from the AfG° values given in different compilations are difficult to interpret, as already the ArG° values reported in different compilations for the redox equilibria between the 3+ 3+ Bi ion and Bi(cr) differ by 10 kJ/mol. The use of different ArG° values for Bi results in a difference of 3.2 log units in the calculated log K*°so values. Thus, log K*°So values calculated from AfG° values are quite uncertain. Solubility products derived from emf measurements at higher temperature (see Table 7.5) are in the range of log K*so of -4 to -6 (using a ArG° value of 95.55 kJ/mol or a log P of-16.74 for the redox equilibria between Bi3+ and Bi(cr); see also Section 7.11: Bi3+/Bi(cr)). 233 JNC TN8400 99-011 Table 7.5: Additional experimentally determined equilibrium data compiled for the precipitation of bismuth hydroxide/oxide. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 7.12: 'Comments on selected references'. Method: cal = calculated from AH and AS values, emf = emf measurements at high temperature, sol = solubility. logK•so Reference Comments KM) Medium Method i+ log K*so: 2Bi + 3H2O <=> a-Bi2O3(cr)+ 6m -3.78 '' [1978CAH/VER] T= 298.15 K, I=n/a n/a emf -4.10 3' [1981GOR/GAV] T= 298.15 K, I=n/a n/a cal -5.90 4 [1982LAP/KOL1 T= 298.15 K, I=diluted water sol 3+ + log K*so: 2Bi + 3H2O -5.08 5 [1943SCH/RIT1 T= 298.15 K, 1=1.41 1.4 NaOH sol log K*S3: Bi(OH)3° <=> 0.5a-Bi2O3(cr) + 1.5H2O 10.80 6 [1971BID] T= 298.15 K, 1=0.5-2.5 NaOH cal 10.68 fl971BID] T= 298.15 K,I=dil water sol log K*S4: Bi(OH)s o 0.5a-Bi2O3(cr) + 1.5H2O + OH' -4.32 6 [197IBID] T= 298.15 K, 1=0.5-2.5 NaOH cal log K*S3: Bi(OH)3° « Bi(OH)3(s) -4.70 7 [1993KRA/DEC] T= 296±0.5 K, 1=1 sol 7 -4.80 ri993KRA/DEC1 T= 296±0.5 K, 1=1 cio4 sol 1 extrapolated from measurements at 1000 K 2 calculated with a AjG° of -95.55 kJ/mol for Bi2+ (Section 7.11). 3 AG values calculated by [1981GOR/GAV] from measurements of heat capacity and data taken from literature 4 extrapolated from measurements at higher temperatures 5 calculated with a log P,4 =-20.66 (1=1.41) (See Section 7.1.4: Bi(OH)4") 5 calculated by [1971BID] based of the data of [1943SCH/RIT] 7 freshly precipitated Bi(OH)3 (t=30 min.) 234 JNC TN8400 99 -Oil Table 7.6: Thermodynamic data for the precipitation of bismuth hydroxide/oxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log Kso Reference Comments I (M) Medium 3+ + log Kso: 2Bi + 3H2O <=> a-Bi2O3(cr) + 6H -4.44 '-2 H995RIS/HAL1 T= 298.15 K, I=n/a 3+ + log Kso: 2Bi + 3H2O <=> Bi2O3(cr) + 6H -4.16 2 [1952LAT] T=298.15K, I=n/a -4.30 2 [1954COU] T= 298.15 K,I=n/a -4.16 2 [1963WIC/BLO] T= 298.15 K, I=n/a -9.17 [1968ROBAVAL] T= 298.15 K,I=n/a -4.67 2 [1971NAU/RYZ] T= 298.15 K,I=n/a -4.68 2 [1973BAR/KNA] T= 298.15 K, I=n/a -6.92 [1976BAE/MES] T= 298.15 K, 1=0 0 0.18 3 [1976SMI/MAR] T= 298.15 K, 1=0 0 -4.08 3 [1976SMI/MAR] T= 298.15 K, 1=1 1 -5.26 2 [1977BAR/KNA] T= 298.15 K,I=n/a -9.2 [1978ROB/HEM2] T= 298.15 K,I=n/a -5.26 2 [1979KUB/ALC] T= 298.15 K,I=n/a -6.92 [1981BAE/MES] T= 298.15 K, I=n7a -4.71 2 [1982PAN] T= 298.15 K, I=n/a -9.13 [1982WAG/EVA] T=298.15K, I=n/a -4.66 2 [1984VE/TAR] T= 298.15 K, I=n/a -4.67 2 [1985LOV/MEK] T= 298.15 K, I=n/a 3+ + log Kso: Bi + 2H2O <=> BiOOH(s) + 3H -1.562 [1952LAT] T=298.15 K, I=n/a -4.1 [1982WAG/EVA] T= 298.15 K, I=n/a 3+ + log Kso: Bi + 3H2O <=> Bi(OH)3(s) + 3H -7.49 2 [1952LAT] T= 298 K, I=diluted -5.81 2 [1971NAU/RYZ1 T= 298.15 K, I=n/a 1 extrapolated from measurements at higher temperatures 2 calculated with a AfG° of -95.55 kJ/mol for Bi2+ (Section 7.11). 3 calculated with a log Pu =-5.31 (1=0) and -7.44 (1=1) (See Section 7.1.3: Bi(OH)3°) 235 JNC TN8400 99-011 7.3 Bismuth chloride system 7.3.1 Bismuth chloride complexes Bismuth ions form complexes with chloride. Several authors determined stability constants for the formation of chloride complexes in chloride and perchlorate medium. The log P° values given in Table 7.7 are extrapolated to 1=0 by using the SIT equation (see Figures 7.7 - 7.12): Bi3+H hCl- BiCl2+ log 1.1 =3.65, Ae = -0.01 3+ + Bi , h2Cl- <=> BiCl2 log 1.2 = 5.85, Ae = -0.16 3+ Bi H h3Cl- <=> BiCl3°(aq) log 1>3 = 7.62, Ae = -0.19 3+ Bi H h4Ch <=> BiCl4- log 1,4 = 9.06, Ae = -0.12 3+ 2 Bi ^ h5Cl- <=> BiCl5 - log 1*5 = 8.33, AE = -0.31 3+ Bi Hh 6 el- <=> BiCl63- log p°ii6 = 7.64, Ae = -0.12 Based on the data available, only tentative values can be given for the formation of BiCls2" and BiCl63". These complexes however, will only become important in concentrated chloride solution (chloride concentration > 1 M). Table 7.7: Experimental equilibrium data compiled for the bismuth(III) chloride system, 3+ 3 m according to the equilibria Bi + mCl" <=> BiClm - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: cat = cation exchange, sol = solubility, sp = spectrophotometry, pol = polarography and pot = potentiometry. Reference Comments I (M) Medium Method log p\,n 3+ 2 log pltl: Bi + CI- <=> BiCl " 2.36 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaClO4 sol 1.91 [1959DES/PAN] T= 298.15 K, 1=6 6 NaC104, HC1O4 pol 1.96 [1959DES/PAN] T= 298.15 K, 1=4 4 NaC104, HC1O4 pol 2.09 [1959DES/PAN] T= 298.15 K, 1=3 3 NaC104, HC1O4 pol 2.00 [1959DES/PAN] T= 298.15 K, 1=2.5 2.5 NaC104, HC1O4 pol 2.08 [1959DES/PAN] T= 298.15 K, 1=2 2 NaClO4, HC1O4 pol 2.18 [1959DES/PAN] T= 298.15 K, 1=1 1 NaClO4, HC1O4 pol 2.20 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiClO4 pot 3.00 [1969JOH] T= 298 K, 1=4 4 NaClO4 sol 2.16 [1970BONAVAU] T= 303.15 K, 1=2 2 NaClO4 pol 2.35 [1970KAN] T= 298.15 K, 1=5 5 NaCIO, SP 236 JNC TN8400 99-011 Table 7.7: continued 2.82 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 2.71 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 2.53 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 2.36 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 3+ + log fil,2-- Bi + 2CI- <=> BiCl2 3.50 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaClO4 pot 3.80 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaC104 sol 4.58 [1959DES/PAN] T= 298.15 K, 1=6 6 NaC104, HC1O4 pol 4.54 [1959DES/PAN] T= 298.15 K, 1=4 4 NaC104, HC1O4 pol 3.90 [1959DES/PAN] T= 298.15 K, 1=3 3 NaC104, HC1O4 pol 4.04 [1959DES/PAN] T= 298.15 K, 1=2.5 2.5 NaC104, HC1O4 pol 4.22 [1959DES/PAN] T= 298.15 K, 1=2 2 NaC104, HC1O4 pol 3.74 [1959DES/PAN] T= 298.15 K, 1=1 1 NaClO4, HC1O4 pol 3.50 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiG104 pot 4.30 [1969JOH] T= 298 K, 1=4 4 NaC104 sol 3.82 [1970BONAVAU] T= 303.15 K, 1=2 2 NaC104 pol 4.40 [1970KAN] T= 298.15 K, 1=5 5 NaC104 sp 4.44 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 4.04 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 4.66 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 3.61 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 3+ 0 log P1.3: Bi + 3CI- <=> B1CI3 5.35 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaC104 pot 5.90 [1959DES/PAN] T= 298.15 K, 1=6 6 NaC104, HC1O4 pol 6.11 [1959DES/PAN] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 5.40 [1959DES/PAN] T= 298.15 K, 1=3 3 NaC104, HC1O4 pol 5.30 [1959DES/PAN] T= 298.15 K, 1=2.5 2.5 NaC104, HC1O4 pol 5.71 [1959DES/PAN] T= 298.15 K, 1=2 2 NaClO4, HC1O4 pol 4.87 [1959DES/PAN] T= 298.15 K, 1=1 1 NaC104, HC1O4 pol 5.80 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiC104 pot 6.70 [1969JOH] T= 298 K, 1=4 4 NaClO4 sol 5.60 [1970BONAVAU] T= 303.15 K, 1=2 2 NaClO4 pol 5.45 [1970KAN] T= 298.15 K, 1=5 5 NaClO4 sp 5.45 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 5.18 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 6.32 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 4.95 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 237 JNC TN8400 99-011 Table 7.7: continued 3+ log pli4: Bi + 4CI- & BiCl4- 6.10 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaC104 pot 7.69 [1959DES/PAN] T= 298.15 K, 1=6 6 NaC104, HC1O4 pol 6.91 [1959DES/PAN] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 6.87 [1959DES/PAN] T= 298.15 K, 1=3 3 NaC104, HC1O4 pol 7.47 [1959DES/PAN] T= 298.15 K, 1=2.5 2.5 NaC104, HC1O4 pol 7.18 [1959DES/PAN] T= 298.15 K, 1=2 2 NaC104, HC1O4 pol 6.90 [1959DES/PAN] T= 298.15 K, 1=1 1 NaC104, HC1O4 pol 6.75 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiC104 pot 6.90 [1969JOH] T= 298 K, 1=4 4 NaC104 sol 6.90 [1970BON/WAU] T= 303.15 K, 1=2 2 NaC104 pol 6.65 [1970KAN] T= 298.15 K, 1=5 5 NaC104 sp 6.23 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 6.41 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 7.93 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 3+ 2 log P15: Bi + 5Ct <=* BiCl5 - 9.29 [1959DES/PAN] T= 298.15 K, 1=6 6 NaC104, HC1O4 pol 8.49 [1959DES/PAN] T= 298.15 K, 1=4 4 NaC104) HC1O4 pol 7.68 [1959DES/PAN] T= 298.15 K, 1=3 3 NaC104, HC1O4 pol 8.04 [1959DES/PAN] T= 298.15 K, 1=2.5 2.5 NaC104, HC1O4 pol 6.75 [1959DES/PAN] T= 298.15 K, 1=2 2 NaC104, HC1O4 pol 6.65 [1959DES/PAN] T= 298.15 K, 1=1 1 NaC104, HC1O4 pol 6.72 [1961AHR/GRE] T= 298.15 K, 1=2 2 NaC104 pot 7.30 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiC104 pot 8.60 [1969JOH] T= 298 K, 1=4 4 NaC104 sol 7.29 [1970KAN] T= 298.15 K, 1=5 5 NaC104 sp 6.11 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 5.95 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 8.18 [1974FED/KAL] T=298K, 1=3 3 HC1O4 pot 3+ 3 log Pi,6- Bi + 6CI- <=> BiCl6 ~ 7.70 [1959DES/PAN] T= 298.15 K, 1=6 6 NaClO4, HC1O4 pol 7.54 [1959DES/PAN] T= 298.15 K, 1=4 4 NaClO4, HC1O4 pol 6.56 [1961AHR/GRE] T= 298.15 K, 1=2 2 NaC104 pot 7.36 [1963MIR/KUL] T= 298.15 K, 1=3 3 LiC104 pot 8.40 [1969JOH] T= 298 K, 1=4 4 NaC104 sol 7.06 [1970KAN] T= 298.15 K, 1=5 5 NaClO4 sp 6.68 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 6.00 [1974FED/KAL] T= 298 K, 1=3 3 HCIO; pot 238 JNC TN8400 99-011 Bi3+ + Cl- o BiCI2+ 4 6 lm, molal 3+ 2+ Figure 7.7: Plot of log (3U + 6 D vs. Im for the reaction : Bi + Cl~ <^> BiCl at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.01; log P0^] = 3.65. Calculated from data compiled in Table 7.7. 5.5 • 4 6 10 lm, molal 3+ + Figure 7.8: Plot of log pi,2 + 10 D vs. Im for the reaction : Bi + 2Ch <=> BiCl2 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.16; log (3° 1,2 = 5.85. Calculated from data compiled in Table 7.7. 239 JNC TN8400 99 -Oil 3+ Bi +3Ch <=> BiCI3° 10 -i 9.5 lli-cpiiiSi^Eiiill-iS 9 misnp) CM 8.5 - •§i liliiii'iSiii- + 8 "7.5 lift O) 7 S3 O — 6.5 •Illillilt lilis||||il:|i|||j 6 ^ t; : : : ;:: ! :j:r !:: : i: ?; - '- ; p : i i '"^' ' f ':W^-:'vf 4 6 10 lm, molal 3+ Figure 7.9: Plot of log pii3 + 12 D vs. Im for the reaction : Bi + 3Ch <=> BiCl3° at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.19; log (3° 13 = 7.62. Calculated from data compiled in Table 7.7. Bi3+ + 4Ch<=> BiCLf 11 - •.;••:;.: •• .,::=:«?;:;*,;.T-.»:i:f:!y:jrk-ty.;:~y,'.:_•?,? 10.5 iK;.;::^sili!i#;;3i;Sl;blai 10 -"S::i;^li?l;iliiBiii Q ; CM 9.5 - •J^ ^:*lji@iBiiiit 9 ;||||||||||||||:|;| 8.5 rn i : 8 IKfe'K.: i I;:' "::;v::::.:'.;:..;: \,v;-^i";.:,;..;-. 7.5 7 • 6.5 R 2 4 6 8 10 U, molal 3+ Figure 7.10: Plot of log p]>4 + 12 D vs. Im for the reaction : Bi + 4C1~ <^=> BiCl4- at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.12; log p°] 4 = 9.06. Calculated from data compiled in Table 7.7. 240 JNC TN8400 99-011 2 5CI•<=> BiCI5 - • o c. 11 5 o 11 Q 8 o 10 5 - o o + 10 o "1 9 5 - CO. o D) 9 o _g 8 5 = 0 31x + 8.33 8 oo y 7.5 7 —1 \ 1- , molar 10 3+ 2 Figure 7.11: Plot of log pi>5 + 10 D vs. Im for the reaction : Bi + 5C1~ o BiCl5 - at 25 °C. The straight line shows the tentative result of the linear regression: Ae = -0.31; P°i>5 = 8.33. Calculated from data compiled in Table 7.7. BiCL3" 4 6 lm, molal 3+ 3 Figure 7.12: Plot of log pli6 + 6 D vs. Im for the reaction : Bi + 6C1~ <=> BiCl6 - at 25 °C. The straight line shows the tentative result of the linear regression: Ae = - 0.12; log P°i,6 = 7.64. Calculated from data compiled in Table 7.7. 241 JNC TN8400 99 -Oil 7.3.2 BiOCl(s) and Bi(OH)2Cl(s) Thermodynamic data for precipitated BiOCl(s) have been determined by different authors (Table 7.8). No direct determination of the crystalline structure of the precipitated BiOCl(s) are given in the different experimental reports. Table 7.8: Experimental data for the precipitation of BiOCl(s) and Bi(OH)2Cl(s). These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: sol = solubility, pot = potentiometry. log Kso Reference Comments I (M) Medium Method 3+ log Kso: Bi + H2O + Or <=> BiOO(cr) + 2H+ 8.42 i [1918NOY/CHO] T= 298.15 K, 1=0.001-0.5 0 HC1 pot 7.26 2 [1933FEI] T= 298.15 K, 1=0.25 0.25 HC1 sol 7.26 2 [1933FEI] T= 298.15 K, 1=0.54 0.54 HC1 sol 7.15 2 [1933FEI] T= 298.15 K, 1=0.6 0.6 HC1 sol 7.10 2 [1933FEI] 298.15 K, 1=0.9 0.9 HC1 sol 7.01 [1957AHR/GRE] T= 298.15 K, 1=2 2 NaClO4 pot 7.39 [1969JOH] T= 298 K, 1=4 4 NaClO4 sol 1 estimated from potentiometric data assuming a log (3 (Bi(cr)/Bi3*) = 16.74(see Section 7.11: BP/Bi(cr)) 2 2+ + corrected in this report for the formation of BiCl , BiCl2 , BiCl3°, BiCI4" using the constants calculated in Section 7.3.1. Original values given in Table 7.9. [1933FEI] determined the solubility of freshly precipitated Bi(OH)2Cl(s). The data compiled in Table 7.9 seem to indicate a higher solubility of Bi(OH)2Cl(s) than of BiOCl(s). [1933FEI], however, did not correct his measurements for the formation of bismuth chloride complexes, thus resulting in a too high solubility. A correction of the measurements of [1933FEI] for the formation of bismuth chloride complexes gives the same solubility product as reported by [1918NOY/CHO, 1957AHR/GRE, 1969JOH] in Table 7.8. The potentiometric data of [ 1918NOY/CHO] for the formation of BiOCl(s) from Bi(cr) were extrapolated in this report to I = 0 (See [1918NOY/CHO] in Section 7.12: Comments on selected references) and give a log K*°s of -8.32 for the reaction Bi(cr) + Cl~ + H2O <=> BiOCl(s) + 2H+ + 3e-. Correction with a log (3° = 16.74 (Section 7.11) for the reaction Bi(cr) 3+ 3+ <^> Bi + 3e- gives a K*So of 8.42 for the reaction Bi + H2O+ Cl- <=> BiOCl(s) + 2H+ as given in Table 7.8. Extrapolation of the data compiled in Table 7.8 to I = 0 (Figure 7.13) gives: H2O+C1- *=> BiOCl(s) + 2H+ log K*°so = 8.47 242 JNC TN8400 99-011 + Bi3++ BiOCI(s)+2H 10 9.5 9 g s.s 0 2 3 4 lm, molal 3+ Figure 7.13: Plot of log K*so + 8D vs. Im for the reaction Bi + H2O+ Or <=> BiOCl(s) + 2H+ at 25 °C. The straight line shows the result of the linear regression: Ae = -0.18; log K*s° = 8.47. The value at 1=0 is taken from Figure 7.25. Calculated from data compiled in Table 7.8. 7.3.3 BiCl3(s) Thermodynamic data are compiled for the precipitation of BiCl3 in Table 7.10. The solubility product of ~ 103 indicates that under environmental conditions BiOCl will be the stable solid phase and BiCl3 will not precipitate except in very acidic solutions (pH < 0). 7.3.4 Additional equilibrium data compiled for the bismuth chloride system Table 7.9: Additional, experimentally determined equilibrium data compiled for the bismuth(III) chloride system and the precipitation of BiOCl(s) and Bi(OH)2Cl(s). These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in Section 7.12: 'Comments on selected references'. Method: sol = solubility, sp = spectrophotometry, pol = polarography and pot = potentiometry. log pllir Reference Comments I (M) Medium Method 3+ 2 log Pu: Bi + Cl- <=> BiCl ' 2.43 ' [1957NEW/HUM] T= 298.15 K, 1=1 HC1O4 sp 2 2.44 [1969CAR] T= 298.15 K, 1=0.1-1 HC1 pol 3 3.70 [1974FED/KAL1 T= 298 K, 1=0 HC1O, pot 243 JNC TN8400 99 - Oil Table 7.9: continued 3 + log Pi.2.' Bi + + 2Cl~ <=> BiCl2 4.46 ' [1957NEW/HUM] T= 298.15 K, 1=5 1 HC1O4 sp 3.10 2 [1969CAR] T= 298.15 K, 1=0.1-1 HC1 pol 3 5.50 [1974FED/KAL] T= 298 K, 1=0 0 HC1O4 pot 3+ 0 log pu: Bi + 3Cl~ <=> BiCl3 5.76 » [1957NEW/HUM] T= 298.15 K, 1=5 5 HC1O4 sp 3.74 2 [1969CAR] T= 298.15 K, 1=0.1-1 HC1 pol 3 6.90 ri974FED/KAL1 T= 298 K, 1=0 0 HC1O4 pot 3+ log P1A: Bi + 4CI- <=> BiClf l 5.42 [1953BAB/GOL] T= 298.15 K, 1=2,3 or 3 3 KNO3 pot ! 6.24 [1957NEW/HUM] T= 298.15 K, 1=5 5 HC1O4 sp 3.77 2 [1969CAR] T= 298.15 K, 1=0.1-1 HC1 pol 3 7.90 [1974FED/KAU| T= 298 K, 1=0 0 HC1O4 pot 3+ 2 log Pu: Bi + 5CI- <=> BiCl5 - 7.50 ' [1953BAB/GOL] T= 298.15 K, 1=2,3 or 3 3 KNO3 pot 6.72 ' [1957NEW/HUM] T= 298.15 K, 1=5 5 HC1O4 sp 4 7.72 [1959AHR/GRE] T= 298.15 K, 1=2 2 NaClO4 pot 3 7.00 f!974FED/KAL1 T= 298 K, 1=0 0 HC1O4 pot 3+ 3 log Pi,6: Bi + 6CI- » BiCl6 - 5 6.42 [1953BAB/GOL] T= 298.15 K, 1=2,3 or 3 3 KNO3 pot 4 7.56 [1959AHR/GRE] T= 298.15 K, 1=2 2 NaC104 pot 3 7.30 T1974FED/KAL1 T= 298 K, 1=0 0 HC1O4 pot 3+ + log Kso: Bi + 2H2O + Ct <=> Bi(0H)2Cl(s) + 2H 2.25 6 [1933FEI] T= 298.15 K, 1=0.3 0.3 HC1 sol 2.56 6 [1933FEI] T= 298.15 K, 1=0.5 0.5 HC1 sol 2.49 6 [1933FEI] T= 298.15 K, 1=0.6 0.6 HC1 sol 2.53 6 [1933FEI1 T= 298.15 K, 1=0.9 0.9 HC1 sol 3+ + log Kso: Bi + H2O + Cl- <=> BiOCl(s) + 2H 7 8.63 [1918NOY/CHO] T= 298.15 K, 1=0.001-0.5 0 HC1O4 pot 8 8.13 [1969VAS/GRE1 T= 298.15 K, 1=1-3 0 HC1O4 pot 1 as reported by [1970KAN] 2 I not constant 3 not reported how data were extrapolated to 1=0 by [1974FED/KAL] 4 corrected later by [1961AHR/GRE] 5 reported by [1957AHR7GRE], Original not available 6 formation of Bi chloride species neglected, corrected values in Table 7.8 7 value corrected and selected by [1918NOY/CHO], Calculated in this report from E° = 0.1599 V and log K (Bi(cr)/Bi3+) = 16.74 (Section 7.11). Additional values by [1918NOY/CHO] are given in section 7.27. 8 as reported by [1985LOV/MEK], extrapolated to 1=0 by [1985LOV/MEK], Calculated in this report from E° = 0.1697 V and log K (Bi(cr)/Bi?+) = 16.74 (Section 7.11) 244 JNC TN8400 99-011 Table 7.10: Thermodynamic data for the bismuth(III) chloride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log P,m Reference Comments Medium log Pij: Bi3+ + Cl- <=> BiCl2+ 2.20 [1967AHR] T= 298.15 K, 1=3 3 LiC104 2.78 [1967VAS/LOB] T=298.15K, 1=4 4 NaClO4 2.94 ' [1967VAS/LOB] T=298.15K, 1=5 5 NaC104 3.12 > [1967VAS/LOB] T= 298.15 K, 1=6 6 NaC104 3.42 ' [1967VAS/LOB] T= 298.15 K, 1=0 0 NaC104 2.36 [1976SMI/MAR] T= 293.15 K, 1=2 2 2.20 [1976SMI/MAR] T= 298.15 K, 1=3 3 2.25 [1982WAG/EVA] T= 298.15 K, 1=0 0 3.50 [1982SMI/MAR] T= 298.15 K, 1=0 0 2.40 [1982SMI/MAR] T= 298.15 K, 1=1 1 2.30 [1982SMI/MAR] T= 298.15 K, 1=4 4 3.42 [1985BAB/MAT] T= 298.15 K, 1=0 0 2.33 [1985LOV/MEK1 T= 298.15 K, 1=0 0 l 3+ + og Pi.2: Bi + 2Cl~ <=> BiCl2 3.50 [1976SMI/MAR] T= 293.15 K, 1=2 2 3.50 [1976SMI/MAR] T= 298.15 K, 1=3 3 4.51 [1982WAG/EVA] T= 298.15 K, 1=0 0 5.50 [1982SMI/MAR] T= 298.15 K, 1=0 0 4.00 [1982SMI/MAR] T= 298.15 K, 1=1 1 4.20 [1982SMI/MAR] T= 298.15 K, 1=4 4 4.99 [1985BAB/MAT] T= 298.15 K, 1=0 0 3.45 [1985LOV/MEK1 T= 298.15 K, 1=0 0 3+ 0 log p,j: Bi + 3CI- t=> BiCl3 5.40 [1976SMI/MAR] T= 293.15 K, 1=2 2 5.80 [1976SMI/MAR] T= 298.15 K, 1=3 3 7.10 [1982SMI/MAR] T= 298.15 K, 1=0 0 5.20 [1982SMI/MAR] T= 298.15 K, 1=1 1 6.00 [1982SMI/MAR] T= 298.15 K, 1=4 4 6.11 [1985BAB/MAT] T= 298.15 K, 1=0 0 5.28 [1985LOV/MEK1 T= 298.15 K, 1=0 0 3+ log J3,i4: Bi + 4CI- t=> BiClf 8.48 2 [1952LAT] T= 298.15 K, 1=0 0 6.10 [1976SMI/MAR] T= 293.15 K, 1=2 2 6.80 [1976SMI/MAR] T= 298.15 K, 1=3 3 6.91 [1982WAG/EVA] T= 298.15 K,I=0 0 8.10 [1982SMI/MAR] T= 298.15 K, 1=0 0 6.40 [1982SMI/MAR] T= 298.15 K, 1=1 1 245 JNC TN8400 99-011 Table 7.10: continued 7.30 [1982SMI/MAR] T= 298.15 K, 1=4 4 8.12 [1985BAB/MAT] T= 298.15 K, 1=0 0 6.04 [1985LOV/MEK1 T= 298.15 K, 1=0 0 3+ 2 log p]:5: Bi + 5CI- <=> BiCl5 ~ 7.69 [1960FRI/SAR] T= 298.15 K, I=n/a 6.70 [1976SMI/MAR] T= 293.15 K, 1=2 2 7.30 [1976SMI/MAR] T= 298.15 K, 1=3 3 8.30 [1982SMI/MAR] T= 298.15 K, 1=4 4 6.65 [1985LOV/MEK] T= 298.15 K, 1=0 0 3+ 3 log pl6: Bi + 6CI- <=> BiCl6 ~ 6.60 [1976SMI/MAR] T= 293.15 K, 1=2 2 7.40 [1976SMI/MAR] T= 298.15 K, 1=3 3 7.40 [1982WAG/EVA] T= 298.15 K, 1=0 0 7.90 [1982SMI/MAR] T= 298.15 K, 1=4 4 6.50 fl985L0V/MEK1 T= 298.15 K, 1=0 0_ 3+ log Kso: Bi + 3Cl~ <=> BiCl3(s) 3.65 2 [1952LAT] T= 298.15 K, I=n/a 3.47 2 [1963WIC/BLO] T= 298.15 K, I=n/a 2.63 2 [1977BAR/KNA] 298.15 K, I=n/a 2.63 2 [1979KUB/ALC] T= 298.15 K, I=n/a 0.73 [1982WAG/EVA] T= 298.15 K, I=n/a 2.23 [1985LOV/MEK] 298.15 K, I=n/a 3+ + log Kso: Bi + 2H2O + Ct <=> Bi(OH)2Cl(s) + 2H ZS [1982WAG/EVA1 T= 298.15 K, I=n/a 3+ + log Kso: Bi + H2O + Ct <=> BiOCl(s) + 2H 8.6 2 [1952LAT] T= 298.15 K,I=n/a 8.14 2 [1971NAU/RYZ] T= 298.15 K,I=n/a 8.45 2 [1977BAR/KNA] T= 298.15 K,I=n/a 8.43 2 [1979KUB/ALC] T= 298.15 K, I=n/a -7.80 3 [1976SMI/MAR] T= 293.15 K, 1=0 0 -6.47 3 [1976SMI/MAR] T= 293.15 K, 1=1 1 -6.75 3 [1976SMI/MAR] T= 293.15 K, 1=3 3 6.40 [1982WAG/EVA] T= 298.15 K,I=n/a 7.49 ri985LOV/MEKl T= 298.15 K,I=n/a 1 the values given in [1967VAS/LOB] are extrapolated to I = 4, 5 and 6 from data compiled in [1964SIL/MAR] 2 calculated with a A,G° of -95.55 kj/mol for Bi2+ (Section 7.11). 3 the data compiled in [1976SM1/MAR] seem to have the wrong sign. 246 JNC TN8400 99-011 7.4 Bismuth perchlorate 2+ 7.4.1 BiClO4 or [Bi 2+ [1993KRA/DEC] reported a log pi;i of 3.5 at I = 1 for the formation of BiC104 . However, this constant for BiC1042+ seems to be much too large compared to the constants of the bismuth complexes with nitrate and chloride. [1993KRA/DEC] present in their report no direct experimental proof for the predominance of the species BiC1042+ (instead of Bi3+) at pH < 2. [1990SUG/ONO] deduced from cation-exchange measurements the dominance of the 3+ Bi(H2O)6 - C1O4- ion pair in 1 M NaC104 at a pH of 1. Additionally, [1982HAT/SUG] and [1987SUG/SHI] showed with hydrolysis experiments, that the presence of both nitrate and perchlorate had only a small influence while the influence of chloride (and thus also the complex formation with chloride) is more important. Therefore, the value calculated by [1993KRA/DEC] is not recommended in this report. 7.4.2 BiOClO4(precip) In 1 M perchlorate medium, BiOClO4(s) precipitates below a pH of 6 [1993KRA/DEC]. At I = 1, [1993KRA/DEC] calculate a log Ks0* for BiOClO4(precip) of 0.87. Extrapolation to 1=0 assuming a Ae of 0 results in a tentative value log Kso*° of -0.78 for the reaction Bi3+ + H2O + CIO4- <=> BiOClO4(precip) + 2H+. Table 7.11: Experimentally determined equilibrium data compiled for the bismuth(III) perchlorate system and the precipitation of Bi0C104(s). These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting this reference is given in the text. Method: sol = solubility. log P,m Reference Comments KM) Medium Method + 2 + log pu: Bi3 + ClOf <=> BiClO4 3.5 [1993KRA/DEC] T= 296±0.5 K, 1=1 1 cicv sol 3+ + log K*so: Bi + H2O + ClOf <=> BiOClO4 (precip) + 2H 0.87 ri993KRA/DEC] T= 296+0.5 K, 1=1 1 cicv sol 247 JNC TN8400 99-011 7.5 Bismuth fluoride system In the presence of fluoride, bismuth fluoride complexes can be formed. Some constants were determined by [1967LOM/VAN] and [1969BON] in 1.9 M and 2 M perchloric acid (Table 7.12). Unfortunately, there are not enough data available for the extrapolation to I = 0. From the thermodynamic data given in the compilations of [1977BAR/KNA] and [1979KUB/ALC] the formation of quite insoluble BiF3(s) is expected. Direct measurements of the solubility, or details about the formation of this solid species, have not been found in the literature, however. No thermodynamic data for the formation of bismuth fluoride complexes or solids are recommended in this report. Table 7.12: Experimentally determined equilibrium data compiled for the bismuth(III) fluoride system. These data are not chosen in this report for the evaluation of recommended log (3 values. Reasons for not selecting these references are given in Section 7.12: 'Comments on selected references'. Method: cat = cation exchange, pol = polarography log (3l.m Reference Comments I(M) Medium Method log J5 + HF <=> BiF2+ +H+ 1 1.43 [1967LOM/VAN] T= 298,1=1.89 1.9 HC1O4 cat 2 1.41 [1969BON] T= 303 K, 1=2 2 NaBr, HC1O4 pol + log ft7.2-" + 2HF <=> BiF2 H-2H 1 1.76 [1967LOM/VAN] T= 298,1=1.89 1.9 HC1O4 cat 2 0.30 [1969BON1 T= 303 K, 1=2 2 NaBr, HC1O4 pol 1 3+ logfi Bi + 3HF <=> BiF3° -3H-4 2 2.70 [1969BON1 T= 303 K, 1=2 2 NaBr, HC1O4 pol >pH<0 2 pH < 0.3 248 JNC TN8400 99 -Oil Table 7.13: Thermodynamic data for the bismuth(III) fluoride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log P,.n Reference Comments I (M) Medium 3+ 2+ + log pu: Bi + HF <=> BiF + H 1.70 [1980BON/HEF] T= 298.15 K, 1=1.89 1.9 HC1O4 1.42 [1976SMI/MAR1 T= 298 K, 1=2 2 3+ + + log Pli2:Bi + 2HF <=> BiF2 + 2H 2.30 [1980BON/HEF1 T= 298.15 K, 1=1.89 1.9 HC1O, 3+ + log Pu: Bi + 3HF <^> BiF3° + 3H 4.70 [T980BON/HEF1 T= 303 K, 1=2 2 NaBr, HC1O4 3+ log Kso: Bi + 3F~ <=> BiF3(s) 15.4 ' [1963WIC/BLO] T= 298.15 K, I=n/a 14.7 ! [1977BAR/KNA] T= 298.15 K, I=n/a 15.8 ' ri979KUB/ALC1 T= 298.15 K, I=n/a 1 2+ calculated with a AfG° of -95.55 kJ/mol for Bi (Section 7.11) 249 JNC TN8400 99-011 7.6 Bismuth carbonate system No data concerning dissolved bismuth carbonate complexes have been found in the literature. [1984TAY/SUN] determined with X-ray analysis the conditions under which (BiO)4(OH)2CO3(cr) and (BiO)2CO3(cr) were predominantly formed in aqueous carbonate solutions (I, pCO2, pH were varied) and extrapolated their results to 1=0 with the Davies equation. They observed the formation of these solids at room temperature. From the AfG° values given in their report, the following log K*°so values can be calculated (Table 7.14): 2 + CO3 -+6H2O (BiO)4(OH)2CO3(cr) + 10 H log K*°so = 8.68 2 CO3 -+2H2O (BiO)2CO3(cr) + 4 H+ log K*°so = 14.27 Table 7.14: Experimentally determined equilibrium data compiled for the bismuth(III) carbonate system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: cat = cation exchange, sol = solubility log Reference Comments I (M) Medium Method 3+ 2 log KSo*: 4Bi + CO3 ~ + 6H2O <=> (BiO)4(OH)2CO3(s) + 10 H+ 8.68 ' [1984TAY/SUN] T= 298 K, 1=0.3-1 0 KOH sol 3+ 2 + log Kso*-- 2Bi + CO3 ~ + 2H2O <=> (BiO)2CO3(s) + 4 H 14.27 > [1984TAY/SUN] T= 298 K, 1=0.3-1 0 KOH sol 1 0 [1984TAY/SUN] gives (based on a selected A,G° (Bi2O3) of-493.5 kJ/mol), A^ of-1678 kJ/mol and -945 kJ/mol for (BiO)4(OH)2CO3(s) and (BiO)2CO3(s), respectively. This values are corrected for difference in A(G° of Bi2O3 between [1984TAY/SUN] and this report (AfG° = -515.99 kJ/mol, Section 7.2.1). 250 JNC TN8400 99-011 7. 7 Bismuth nitrate system 7. 7.1 Bismuth nitrate complexes Bismuth forms weak complexes with nitrate. Several authors determined the equilibrium constants for bismuth nitrate complexes. The existence of bismuth nitrate complexes in acid solutions is well established by anion exchange results and Raman spectra ([1954NEL/KRA], [1968OER/PLA]). Extrapolation of the measurements of [1967KAP/NAP], [1971 FED/KALI], [1990SUG/ONO1] and [1993KRA/DEC] as given in Table 7.15 results in the following formation constants (Figures 7.14 - 7.17): 2+ BiNO3 log P°,,, = 1.97 Bi(NO3)2+ log p°li2 = 2.95 3+ Bi + 3NO3- Bi(NO3)3° log (3°1,3 = 3.62 Bi(NO3)4- log (3\4 = 3.09 2 3 The formation constants for Bi(NO3)s ~ and Bi(NO3)6 ~ have only been determined at high ionic strength and show a large difference (see Table 7.17). Thus, no constants for the 2 3 formation of Bi(NO3)s ~ and Bi(NO3)6 ~ are selected in this report. It can be stated that the complex formation of bismuth with nitrate is much weaker than with chloride. Table 7.15: Experimental equilibrium data compiled for the bismuth(III) nitrate system, 3+ 3 m according to the equilibria Bi + mN03- <=> Bi (NO3)m - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: cat = cation exchange, pot = potentiometry, sol = solubility. log P, Reference Comments I (M) Medium Method 3+ 2+ log J5U: Bi + NO 3- t=> BiNO3 0.96 [1967KAP/NAB] T= 298.15 K, 1=1 1 NaClO4 cat 0.72 [1971FED/KAL1] T= 298.15 K, 1=0.5 0.5 LiClO4 pot 0..81 [1971FED/KAL1] T= 298.15 K, 1=1 1 LiClO4 pot 0.72 [1971 FED/KALI] T= 298.15 K, 1=2 2 LiClO4 pot 0.72 [1971FED/KAL1] T= 298.15 K, 1=3 3 LiClO4 pot 0.92 [1971 FED/KALI] T= 298.15 K, 1=4 4 LiClO4 pot 0.73 [1971 FED/KALI] T= 298.15 K, 1=3 3 LiClO4 pot 0.74 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 1.20 [1993KR A/DEC] T= 296+0.5 K, 1=1 1 NO," sol 25; JNC TN8400 99-011 Table 7.15: continued 3+ log Bi + 2NO3- <=> Bi(NO3)2 1.58 [1967KAP/NAB] T= 298.15 K, 1=1 1 NaC104 cat 0.94 [1971FED/KAL1] T= 298.15 K, 1=0.5 0.5 LiC104 pot 0.90 [1971FED/KAL1] T= 298.15 K, 1=1 1 LiC104 pot 0.98 [1971FED/KAL1] T= 298.15 K, 1=2 2 LiC104 pot 0.96 [1971FED/KAL1] T= 298.15 K, 1=3 3 LiClO4 pot 1.23 [1971 FED/KALI] T= 298.15 K, 1=4 4 LiC104 pot 1.16 [1971 FED/KALI] T= 298.15 K, 1=3 3 LiC104 pot 1.22 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 3+ log Pl,ij: Bi + 3NO3~ & Bi(NO3)3° 1.93 [1967KAP/NAB] T= 298.15 K, 1=1 1 NaClO4 cat 0.72 [1971 FED/KALI] T= 298.15 K, 1=1 1 LiC104 pot 0.20 [1971FED/KAL1] T= 298.15 K, 1=2 2 LiC104 pot 0.11 [1971FED/KAL1] T= 298.15 K, 1=3 3 LiC104 pot 1.08 [1971FED/KAL1] T= 298.15 K, 1=4 4 LiC104 pot 0.88 [1971 FED/KALI] T= 298.15 K, 1=3 3 LiClO4 pot 1.54 [1990SUG/ONO1] T= 298.15 K, 1=1 1 HC1O4 cat 3+ l°g Pi,'t: Bi + 4NO3~ <=> Bi(NO3)4- 0.58 [1971FED/KAL1] T= 298.15 K, 1=2 2 LiC104 pot -0.22 [1971FED/KAL1] T= 298.15 K, 1=3 3 LiClO4 pot 0.40 [1971FED/KAL1] T= 298.15 K, 1=4 4 LiC104 pot 0.54 [1971 FED/KALI] T= 298.15 K, 1=3 3 LiC104 pot 252 JNC TN8400 99-011 2 3 lm, molal 3+ 2+ Figure 7.14: Plot of log pu + 6D vs. Im for the reaction Bi + NO3- <=>BiNO3 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.07; log (31,1° = 1.97. Calculated from data compiled in Table 7.15. 5 i::!;:!af!::!g|;;-i:;i;^ 4.5 '• :•••". ".yr-y: y: -y/^i^r- '-'• '•-'• '•-•'••'h'.:.-;::-.rJ^.''-''-'-^['[ •'. \\ : :;.-Y-. •; "."••••;},': •:.-]•:--l:%]-^^'^.:\^{"\:\v--: •-.iifgiijiflnhfgsjjiif; : : ; 4 ••:\-:^:-::;^ -"-.-:-..v^::K™;ii^:..^j :;;:;ria:| ;j;:i:;:::i:;i::ij;iv:E:;- lillillillli Q ;?S;C;f -ff^Qw^ illlillsS o 3.5 • !li3!S3|S||p£ i : ,|i^:;sjj!i:|i.- j:^.i;:i; i^iii:- ".:.h;;,^:;j.s-^;:;i;:i;i:.-::v:. 3 : : i: : Bi:li'!:^b: .j::i::' ;.l ': :';:-h'^.'-: + mmmmmmm i l l j i ; : CM -yiK^'O-;Sht0^i:i"f^fvmbX:K Hsi '!i Wh ' : :i!:-V '\ ,-T 2.5 • : : ' -•..'. '•: . i: -i -.-'•'•':: •"-< ':• .'•::.'0 ~ '' ''i\ :-"-'-T-Js- :•..:•'• •">£[•• 1 1 1 CO. ;•• •:;;!-;;.:!::: rii'j^vi'-iiji": f:-:i;i"?^ •-.".*;--:i*j=; I•;•;!!;-:;i i:;::;: ••'ISljifslJjtiiilSr : ? ':' en 2 jo "-: :ii!;:;.: if iO.::!= J;y?;w jss;! J;::HSl:i;;:y 1.5 - Kftiij W&i- iBTW:' ' •:::::;:^p::^|^^j-:'i;v..;;-^> 0.5 • 5 :; g ::L :fr : 0 - ' l' " 2 3 lm, molal 3+ + Figure 7.15: Plot of log pu + 10D vs. Im for the reaction Bi + 2NO3- o Bi(NO3)2 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.15; log Pi,20 = 2.95. Calculated from data compiled in Table 7.15. 253 JNC TN8400 99 -Oil 3+ Bi +3NO3- oBi(N03)3° 4.5 i!l!iilt|iilJll!i : 4 i|S|l|fi;;;:Jl|f-: ::i||:ii;s|:p pjjillligll «^&:S-&»;:-o;ri-:::--:i;-:f?&iiiiffiW^ 3.5 i ; 1 1 : 1 - ::•-• ""•••• :" iv' 1--" *• =. ,!.':J'"".;;::: r;: ;;-: .^..-•';::••.;•; • : - . : : : : : : =:. "' ;•: ;-:-.[-~ =" '-' -~-. ;••••••;.:"/• • :V ••.:...v : »:."-::;. •-•::• ..'•:-:•*' •: " 3 |f;::|;^l||ll|i|i|:: r : 1 : ! : ; : 11 : : --- F-'^': -j::i>:ii; ii-;=:J - O^' •-•::.[ .: i'=7jii^-iKJ ".:-' : + 12 D :.;-,;;;;i^';-:.:..|;:..;:.;:.';.|: :;"vi.";';;; ': ==-E:-;=§::... -.-\r-.-\ •:•. ] ••,;;• CO : 2.5 CO. |||i|l|Il|f|l||i|;;| 2 lo g f i : i; : : 1.5 lfll||3j||:llillllllgiii;^:^™' "U*UO•''A- AQv^LM^QA T OiQ&u CO ^^'> - -• 1 -lli;!fflliIfiliSsl|iSSi§ 0.5 • n - ='••:-•• - . .:.|:':'.-).|r.:;v,-;--..-:v...;:v.'/f:-^-r:'- 1 2 3 lm, molal 3+ Figure 7.16: Plot of log p1>3 + 12D vs. Im for the reaction Bi + 3NO3- <=> Bi(NO3)3° at 25 °C. The straight line shows the result of the linear regression: Ae = 0.03; log (3°it3 = 3.62. Calculated from data compiled in Table 7.15. 3+ Bi +4NO3"^Bi(NO3)4" o • 4.5 iil|6il;|i!l|S;llllSI: 4 1 : 1 : ; i; : -1: '-i^r T ii-ii":i:fi->-r-5:iivi; ; i;:iiuin;ii!iH;:;^:^V^KL o.-' + 3 ^2.5 iSfii}i|}|liipiiif|ii|: D) 2 l!lff-|P|:p3 ;:;!-:::f:||SSB:;f : •"..•:L-.K:iir.pi'.:t:.ii!1.i:-.:; jiji.:. ;:•.•:; :. 'i. !.';ji'j1"vv;j::''-';:: •;". 1.5 • ,>""i iij !. i-' -i.:: ••:-••••'•• ••'."• . ':'-l:-:::-:-'.;:-v:.•;:':;•;•'••; ,- »•- ""\Ji\J- 1 -jC- *T" »J • Uw ''"•••• 0.5 - ; : : : ;; : :v n - "' '*' "' i' ' ' ' ' ''1 " ' ' "'' " — 2 3 lm, molal 3+ Figure 7.17: Plot of log pli4 + 12D vs. Im for the reaction Bi + 4NO3- <=> Bi(NO3)4- at 25 °C. The straight line shows the result of the linear regression: Ae = 0.01; log P°i,4 = 3.09. Calculated from data compiled in Table 7.15. 254 JNC TN8400 99-011 7.7.2 BiONO3(s) In 1 M nitrate medium BiONO3(s) precipitates below pH 6 [1993KRA/DEC]. The data given in [1951SWI/GAR] were recalculated in this report including a correction for the formation of 2+ BiNO3 complex (see [1951SWI/GAR] and Figure 7.26 in Section 7.12: Comments on selected references). Extrapolation of the data of [1951SWI/GAR] and [1993KRA/DEC] (Table 7.16) is shown in Figure 7.18 and result in: 3+ + Bi + H2O + NO3- BiONO3(s) + 2H log Ks0*° = 2.75; Ae = - 0.09 3+ + Bi +NO3-+H2O<^BiONO3(s)+2H 4.5 - 4 -• Q 3.5 | CO + 3 -: o G 2.5 - 1.5 -•• 1 0.5 - 0 0.5 1 1.5 lm, molal 3+ Figure 7.18: Plot of log Ks0* + 8D vs. Im for the reaction Bi + H2O + NO3~ <=> BiONO3(s) + 2H+ at 25 °C. The straight line shows the result of the 'linear regression': Ae = -0.09; log KSo*° = 2.75. The value at 1=0 results from Figure 7.29. Calculated from data compiled in Table 7.16. 255 JNC TN8400 99-011 Table 7.16: Experimental equilibrium data compiled for the precipitation of BiONC>3(s). These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: sol = solubility. Reference Comments I (M) Medium Method log KSo 3+ + log Kso: Bi + H20 + N03- + 2H l 2.75 [1951SWI/GAR] T= 298.15 K, 1=0.1-0.025 0 HNO3 sol 1.20 [1993KRA/DEC] T= 296+0.5 K, 1=1 1 NO," sol 1 2+ recalculated considering the formation of BiNO3 complexes; See Figure 7.26 ([1955SWI/GAR]) in the Section 7.12: Comments on selected references) 7. 7.3 Additional equilibrium constants compiled for the bismuth(III) nitrate system Table 7.17: Additional, experimentally determined equilibrium data compiled for the bismuth(III) nitrate 3+ 3 m system, according to the equilibria Bi + mNO3~ <=> Bi (NO3)m - . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text. Method: cat = cation exchange, sol = solubility, pot = potentiometry. log P,.n Reference Comments I (M) Medium Method 3+ 2+ log f}u: Bi + NOf <=> BiNO3 1.74 ' [1971FED/KAL1] T=298.15K, 1=0 0 LiClO4 pot 2 1.26 [1972BON] T= 298.15 K, I=n/a HNO3 sol 3 1.74 [1974FED/KAL] T= 298 K, 1=0 0 HC1O4 pot 3 0.72 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 3 0.81 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 3 0.72 ri974FED/KAU] T= 298 K, 1=3 3 HC1O4 pot 3+ log Pu: Bi + 2NO3~ 2.55 > [1971FED/KAL1] T= 298.15 K, 1=0 0 LiClO4 pot 3 2.55 [1974FED/KAL] T= 298 K, 1=0 0 HC1O4 pot 3 0.95 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 3 0.90 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 3 0.96 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 3+ log PJ • Bi + 3NO3~ « Bi(NO3)3° 3 0.72 [1974FED/KAL] T=298K, 1=1 HC1O4 pot 3 0.11 [1974FED/KAL1 T= 298 K, 1=3 HC1O4 pot 256 JNC TN8400 99 - Oil Table 7.17: continued 3+ log p]i4: Bi + 4NOf <=> Bi(NO3)4 4 1.99 [1967KAP/NAB] T= 298.15 K, 1=1 1 3 NaNO3 cat -0.22 3 [1974FED/KAL1 T= 298 K, 1=3 3 HC1O. pot 3+ .2- log Pu: Bi + 5NOf <=> Bi(NO3)5 4 3 1.81 [1967KAP/NAB] T= 298.15 K, 1=2 2 NaNO3 cat -0.10 T1971 FED/KALI] T= 298.15 K, 1=3 3 LiClO. pot 3+ .3- log Bi + 6NOf<=> Bi(NO3)6 4 3 1.25 [1967KAP/NAB] T= 298.15 K, 1=3 3 NaNO3 cat -0.40 [1971FED/KAL11 T= 298.15 K, 1=3 3 LiClO4 pot 3 + + log Kso: Bi + H2O + NOf <=> BiONO3(s) + 2H 2.55 5 [1951SWI/GAR1 T= 298.15 K, 1=0 0 HNO, sol 1 extrapolation by [1971 FED/KALI] to 1=0 with Vasilev equation. 2 In [1972BON]. reported from Yatsimirski, 1954 (in Russian, not available). Experimental details not known. 3 same data as in [1971 FED/KALI] 4 ionic strength not constant, estimated 5 original value (not including complex formation of bismuth with nitrate ; recalculated value in Table 7.16). Table 7.18: Thermodynamic data for the bismuth(III) nitrate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log PV Reference Comments 3+ 2+ log Bi + NOf <=> BiNO3 1.70 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.72 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.81 [1976SMI/MAR] T= 298.15 K, 1=1 1 0.72 [1976SMI/MAR] T= 298.15 K, 1=2 2 0.72 [1976SMI/MAR] T= 298.15 K, 1=3 3 0.92 [1976SMI/MAR] T= 298.15 K, 1=4 4 3+ + log Bi + 2NOf <=> Bi(NO3)2 2.50 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.94 [1976SMI/MAR] T= 298.15 K, 1=0.5 0.5 0.90 [1976SMI/MAR] T= 298.15 K, 1=1 1 0.98 [1976SMIMAR] T= 298.15 K, 1=2 2 0.96 [1976SMI/MAR] T= 298.15 K, 1=3 3 1.23 fl976SMI/MARl T= 298.15 K, 1=4 4 257 JNC TN8400 99-011 Table 7.18: continued + log /3U: Bi-' + 3NO3~ <=> Bi(NO3)3° 0.70 [1976SMI/MAR] T= 298.15 K, 1=1 1 0.20 [1976SMI/MAR] T= 298.15 K, 1=2 2 0.10 [1976SMI/MAR] T= 298.15 K, 1=3 3 1.10 [1976SMI/MAR] T= 298.15 K, 1=4 4_ 3+ log PJA: Bi + 4NOf <=> Bi(NO3)4~ 0.60 [1976SMI/MAR] T= 298.15 K, 1=2 2 -0.20 [1976SMI/MAR] T= 298.15 K, 1=3 3 0.40 H976SMI/MAR1 T= 298.15 K, 1=4 4_ 3+ + log Ks0: Bi + H2O + NOf <=> BiONO3(s) + 2H 2.55 [1976SMI/MAR] T= 298.15 K, 1=0 0 2.64 ri982WAG/EVA] T= 298.15 K, I=n/a 258 JNC TN8400 99-011 7.8 Mixed bismuth nitrate and chloride system 7.8.1 Bismuth chloride nitrate complexes The formation of mixed bismuth chloride nitrate complexes is also described in the literature. 3m n + [1974FED/KAL] measured log p values for the formation of mixed BimCln(NO3)0( - -°) complexes (given in Table 7.19). Extrapolation of these values to 1=0 with SIT as shown in Figures 7.19 to 7.23 results in: BiCl(NO3)+ log P°u,i =5.16, Ae = -0.11 BiCl2(NO3)° log p\2'i = 6.86, Ae = -0.30 BiCl3(NO3)- log (3°!^] = 8.09, AE = -0.45 BiCl(NO3)20 log p°i'i'2 = 5.28, Ae = -0.27 BiCl2(NO3)2- log p°li2i2 = 5.75, As = -0.67 Thermodynamic formation constants for other bismuth chloride complexes (BiCl(NO3)3~, 2 2 BiCl2(NO3)3 -, and BiCl3(NO3)2 ~) were determined only at I = 3 and could therefore not be extrapolated to I = 0. These values are compiled in Table 7.20. Again, it can be observed that the influence of chloride is stronger than the influence of nitrate. Table 7.19: Experimental equilibrium data compiled for the bismuth(IH) chloride nitrate 3+ 3 m n system, according to the equilibria Bi + mCl - + nNO3~ <=> BiClm(NO3)n - - . These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: cat = cation exchange, sol = solubility, sp = spectrophotometry, pol = polarography and pot = potentiometry. Reference Comments I (M) Medium Method Pl.m,n 3+ + log pUJ: Bi + Cl- + NO3- BiCl(NO3) 3.40 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 3.36 [1974FED/KAL] T= 298 K, 1= 1 1 HC1O4 pot 3.15 [1974FED/KAL] T= 298 K, 1=3 3 HC1CX pot 3+ log ft • Bi + 2CI- +2NO3- d BiCl2(NO3)° 4.60 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 5.17 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 5.00 [1974FED/KAL] T= 298 K, 1=3 3 HC1O4 pot 259 JNC TN8400 99-011 Table 7.19: continued 3+ log p1>3J: Bi + 3CI- + NO3- <=> BiCl3(NO3)- 6.30 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 6.14 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 6.93 [1974FED/KAL] T= 298 K, 1=3 3 HCKX pot Bi3+ 0 log $UX- + CI- + 2NO3~ & BiCl(NO3)2 3.19 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 3.36 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 3.38 [1974FED/KAL] T= 298 K, 1=3 3 HC1O, pot 3+ log $ 2: Bi + 2CI- + 2NO3~ <=> BiCl2(NO3)2- 3.85 [1974FED/KAL] T= 298 K, 1=0.5 0.5 HC1O4 pot 4.30 [1974FED/KAL] T= 298 K, 1=1 1 HC1O4 pot 5.29 [1974FED/KAL] T= 298 K, 1=3 3 HC1CX pot BiCINO, 1 2 lmi molal 3+ Figure 7.19: Plot of log pu>] + 10D vs. Im for the reaction Bi + Cl" + NO3- <=> BiClNO3+ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.11; log p\u° = 5.16. Calculated from data compiled in Table 7.19. 260 JNC TN8400 99 -Oil BiCI2(NO3)° 1 2 lm, molal 3+ Figure 7.20: Plot of log pu,i + 12D vs. Im for the reaction Bi + 2C1~ + NO3- <=> BiCl2NO3 0 at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.30; log Pi^,!0 = 6.86. Calculated from data compiled in Table 7.19. BiCI3(NO3)- 10 11 1 1 \J 9.5 Illlilil '-.vi.-i'vi::- r"i:'::;^;/i: ifVX ••'". . '•-" • 9 '• i •'•'-':^': ^'.Lii.ii'i::;-::!.'. >:"-"/ " "•' Q CM 8.5 • + 8 ":7.5 • O) 7 • o 6.5 6 - 5.5 - 5 • ——-—"• ••" '' i 1 2 3 L molal 3+ Figure 7.21: Plot of log p1)3>1 + 12D vs. Im for the reaction Bi + 3C1" + NO3- <=> BiCl3(NO3)~ at 25 °C. The straight line shows the result of the linear regression: o Ae = - 0.45; log p1>3il = 8.09. Calculated from data compiled in Table 7.19. 261 JNC TN8400 99 -Oil BiCI(NO3)2 CO. O) 5 fi o 4.5 £ 4 3.5 | 3 -t- --F- 1 2 3 lm, molal 2D 3+ Figure 7.22: Plot of log p1>lj2 + 1 vs. Im for the reaction Bi + Ch + 2NO3- <=> BiCl(NO3)2° at 25 °C. The straight line shows the result of the linear regression: O Ae = - 0.27; log (3U>2 = 5.28. Calculated from data compiled in Table 7.19. BiCI2(NO3)2 y = 0.67x4-5,75 1 2 3 lm, molal 3+ Figure 7.23: Plot of log pi,2,2 + 12D vs. Im for the reaction Bi + 2C1~ + 2NO3~ <=> BiCl2(NO3)2~ at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.67; log Pi,2,2° = 5.75. Calculated from data compiled in Table 7.19. 262 JNC TN8400 99-011 7.5.2 Additional equilibrium data compiled for the bismuth(III) chloride nitrate system Table 7.20: Additional, experimentally determined equilibrium data compiled for the bismuth(III) chloride 3+ 3 m n nitrate system, according to the equilibria Bi + mCl" + nNO3~ <=> BiClm(NO3)n ~ - . These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry. l°g Pi.m.n Reference Comments I (M) Medium Method 3+ + log pUJ: Bi + Cl- + NOf <=> BiCl(NO3) 5.00 ' r 1974FED/KAL1 T= 298 K, HC1O. pot 3+ log PUJ: Bi + 2Ct- + NOf <=> BiCl2(NO3)° 6.90 ' H974FED/KAL1 T= 298 K, HC1O, pot 3+ log P,,3j: Bi + 3Cl~ + NOf <=> BiCl3(NO3)- 7.80 ' [ 1974FED/KAL1 T= 298 K, 1=0 pot l : Bi3+ °g Pi.i.2 + Cl- + 2N0f <=> BiCl(NO3)2° 5.20 ' ri974FED/KAL1 T= 298 K, 1=0 HC1O4 pot Bi3+ log Pi.2,2-- + 2CI- + 2NOf <=> BiCl2(NO3)2- 5.60 ' f!974FED/KALl T= 298 K, 1=0 HC1O,. pot Bi3+ log Pi.J.3- + Cl- + 3N0f <=> BiCl(NO3)f 2.83 [1974FED/KAL1 T= 298 K, 1=3 HC1O. pot 3+ 2 log Pu,3:Bi + 2CI- + 3N0f <=> BiCl2(NO3)3 - 4.38 F1974FED/KAL1 T= 298 K, 1=3 HC1O4 pot Bi3+ 2- log Pi.iX- + 3Cl~ + 2N0f «• BiCl3(NO3)2 6.10 [1974FED/KAL] T= 298 K, 1=3 HCIO. pot Not indicated by [1974FED/KAL] how the data were extrapolated to 1=0. 263 JNC TN8400 99-011 7.9 Bismuth phosphate system From the thermodynamic data given in the compilations of [1971NAU/RYZ] and [1984VTE/TAR] the formation of quite insoluble bismuth phosphate BiPO^s) can be expected. However, direct measurements of the solubility or details about the formation of this solid species have not been found in the literature. BiPC>4(s) is a potentially a solubility determining solid under repository conditions. However, no solubility product can be proposed in this report as experimental data are missing. Table 7.21: Thermodynamic data for the precipitation of BiPO4(s) recommended previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log K$o Reference Comments 3+ log K*so: Bi + PO/- <=> BiPO4(s) 22 .4 ' [1971NAU/RYZ] T= 298.15 K, I=n/a 22 .4 ' [1977TAR/VIE] T= 298.15 K, I=n/a 22 .9 [1984VIE/TAR1 T= 298.15 K, I=n/a calculated with a AfG° of -95.55 kj/mol for Bi2+ (Section 7.11) 264 JNC TN8400 99 - Oil 7.10 Bismuth sulfate system 7.10.1 Bismuth sulfate complexes [1971FED/KAL2] determined equilibrium constants for bismuth sulfate complexes in 3 M sulfate/perchlorate solutions at different temperatures (Table 7.22). Unfortunately, as no other values are available, these values can not be extrapolated to 1=0. 7.10.2 Bi2(SO4)3(s) The formation of Bi2(SO4)3(s) is reported in the literature (see values compiled in Table 7.23). Direct measurements of the solubility or details about the formation of this solid species have not been found in the literature. At higher temperature (800 - 1000 K) mixed precipitates with oxygen (Bi2O2SO4, a-Bi2O(SO4)2 and P-Bi2O(SO4)2 can also be formed [1984JON]. 7.10.3 Equilibrium data compiled for the bismuth sulfate system Table 7.22: Experimentally determined equilibrium data compiled for the bismuth sulfate system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references see text. Method: pot = potentiometry log Pi.r Reference Comments I (M) Medium Method 3+ + log pu: Bi + SO/~ <=> Bi(SO4) 1.98 [1971FED/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiC104 pot 3+ 2 log j3;i2: Bi + 2SO4 ~ <=> Bi(SO4)2~ 3.41 [1971FED/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot 3+ 2 3 log pu: Bi + 3SO4 ' <=> Bi(SO4)3 - 4.08 [1971FBD/KAL2] T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot 3+ 2 5 log P1A: Bi + 4SO4 - <=> Bi(SO4)4 ~ 4.34 ri971FED/KAL21 T= 298.15 K, 1=3.0 HC1O4, LiClQ, pot 3+ 2 7 log pL5: Bi + 5SO4 ~ <=> Bi(SO4)5 ~ A_J6 fl971FED/KAL21 T= 298.15 K, 1=3.0 HC1O4, LiClO4 pot 265 JNC TN8400 99-011 Table 7.23: Thermodynamic data for the bismuth sulfate system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. l°g Pi,m Reference Comments I(M) 3+ 2 + log pu: Bi + SO4 ~ <=> Bi(SO4) 1.98 [1976SMI/MAR1 T= 298.15 K, 1=3.0 3+ 2 log pl2: Bi + 2SO4 ~ <=> Bi(SO4)2- 3.41 [1976SMI/MAR1 T= 298.15 K, 1=3.0 3+ 2 3 log Pu: Bi + 3SO4 ~ o Bi(SO4)3 ~ 4.08 [1976SMI/MAR1 T= 298.15 K, 1=3.0 3+ 2 5 log pli4: Bi + 4SO4 ~ <=> Bi(SO4)4 ~ 4.34 ri976SMI/MAR] T= 298.15 K, 1=3.0 3+ 2 7 log Pu: Bi + 5SO4 ~ <=> Bi(SO4)5 ~ _4.6 [1976SMI/MAR1 T= 298.15 K, 1=3.0 3+ 2 log K*so: 2Bi + 3SO4 ~ t=> Bi2(SO4)3(s) 29.2 ' [1977BAR/KNA], T= 298.15 K, I=n/a 29.2 1 [1979KUB/ALC1, T= 298.15 K, I=n/a 1 2+ calculated with a AfG° of-95.55 kJ/mol for Bi (Section 7.11) 266 JNC TN8400 99-011 7.11 Bi3+/Bi(cr) 3+ The ArG° values for the redox equilibria between Bi ion and metallic rhombohedral Bi(cr) 3+ reported in different compilations differ by 10 kJ/mol. The use of different ArG° values for Bi can result in a difference of 1.6 log units in the calculated log (3° values (compare the values compiled in Table 7.26). Potentiometric measurements of [1969VAS/GLA] with a platinum electrode in perchlorate media are compiled in Table 7.24. Table 7.24: Experimentally determined equilibrium data compiled for the equilibrium Bi3+ + 3e- <=> Bi(cr). These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 7.12: 'Comments on selected references'. Method: pot = potentiometry logK Reference Comments Medium Method log K: Bi3+ + 3e- <=> Bi(cr) 14.86 [1969VAS/GLA] T=298.15, 1=1.2 1.2 HC1O4 pot 14.73 [1969VAS/GLA] T=298.15, 1=2 2 HC1O4 pot 14.79 [1969VAS/GLA] T=298.15, 1=2.3 2.3 HC1O4 pot 14.79 [1969VAS/GLA] T=298.15, 1=2.3 2.6 HC1O4 pot 14.62 [1969VAS/GLA] T=298.15, 1=3 3 HC1O4 pot 14.61 [1969VAS/GLA] T=298.15, 1=3.1 3.1 HC1O4 pot 14.65 [1969VAS/GLA] T=298.15, 1=3.3 3.3 HC1O4 pot 14.73 [1969VAS/GLA] T=298.15, 1=3.6 3.6 HC1O4 pot 14.62 [1969VAS/GLA] T=298.15, 1=4 4 HC1O4 pot 14.61 [1969VAS/GLA] T=298.15, 1=4.3 4.3 HC1O4 pot 14.72 [1969VAS/GLA] T=298.15, 1=4.6 4.6 HC1O4 pot Extrapolation of the potentiometric measurements of [1969VAS/GLA] (compiled in Table 7.24) to 1=0 with the SIT result in (see Figure 7.24): Bi3++ 3e- Bi(cr) logK°= 16.74 E° = 0.330 V 3+ corresponding to a AfG° of 95.55 kJ/mol for Bi . The data of [1984PIN/GAL] in 1 M HN03 and of [1945LIN] in 0.6 M NaCI medium which are compiled in Table 7.25 agree well. 267 JNC TN8400 99-011 Bi(cr) n, molal 3+ Figure 7.24: Plot of log K + 9D vs. Im for the reaction Bi + 3e- <=> Bi(cr) at 25 °C. The straight line shows the result of the linear regression: Ae = - 0.05; log K° = 16.74. Calculated from data compiled in Table 7.24. Table 7.25: Additional, experimentally determined equilibrium data compiled for the equilibrium Bi3+ + 3e" <=> Bi(cr). These data were not chosen in the present report for the evaluation of recommended stability values. Method: pot = potentiometry logK Reference Comments Medium Method log K: Bis+ + 3e- <=> Bi(cr) l 16.09 [1969VAS/GLA] T=298.15,1=0 0 HC1O4 pot 2 15.62 [1975HEI/SCH] T=298.15,1=1.5 1.5 H2SO4 pot 3 14.19 [1984PIN/GAL] T=298.15,1=1 1 HNO3 pot 14.71 4 [1945LIN1 T=298.15,1=0.6 0.6 NaCl, tartaric acid pot 1 extrapolated to 1=0 by [1969VAS/GLA] from the measurements given in Table 7.24 2 2 Bi makes strong complexes with SOt ', exact I not clear 3 complexes with NC>3~ are probable 4 complexes with Cl are probable Table 7.26: Thermodynamic data for the Bi3+/Bi(cr) redox equilibrium taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK Reference Comments log K: Bi3+ + 3e- <=> Bi(cr) 14.50 [1968 ROB AVAL] T=298.15,1=1 14.51 [1982 WAG/EVA] T=298.15,1=0 0 16.08 [1985BAB/MAT], T=298.15,1=0 0 16.09 [1985LOV/MEK] T=298.15,I=0 0 268 JNC TN8400 99-011 7.12 Comments on selected references [1918NOY/CHO]: The potentiometric data of [1918NOY/CHO] (compiled in Table 7.27) for the reaction Bi(cr) + Cl~ + H2O o BiOCl(s) + 2H+ + 3e~ can be extrapolated to I = 0 as shown in Figure 7.25, resulting in a log K*°s of -8.32. Correction for the reaction Bi(cr) <=> Bi3+ + 3e- with a log K° of 16.74 (Section 7.11) gives: H2O+Cl- BiOCl(s)+2H+ log K*°so = 8.42 Further data concerning the solubility of BiOCl(s) are discussed in Section 7.3.2: BiOCl(s) and Bi(OH)2Cl(s) and shown there in Table 7.8 and Figure 7.13. Table 7.27: Experimental data for the precipitation of BiOCl(s) determined by [1918NOY/CHO]. These data were chosen for the evaluation of recommended values in the present report. Method: pot = potentiometry log Kso Reference Comments I (M) Medium Method + log K*so: Bi(cr) + Or + H2O <=> BiOCl(s) + 2H + 3er -8.31 1 [1918NOY/CHOJ T= 298.15 K, 1=0.001 0.000984 HC1 pot -8.31 i [1918NOY/CHO] T= 298.15 K, 1=0.002 0.001 HC1 pot -8.23 1 [1918NOY/CHO] T= 298.15 K, 1=0.003 0.00318 HC1 pot -8.23 i [1918NOY/CHO] T= 298.15 K, 1=0.01 0.01 HC1 pot -8.24 i [1918NOY/CHO] T= 298.15 K, 1=0.02 0.02 HC1 pot -8.28 i [1918NOY/CHO] T= 298.15 K, 1=0.03 0.046 HC1 pot -8.29 ] [1918NOY/CHO] T= 298.15 K, 1=0.1 0.1 HC1 pot -8.30 i [1918NOY/CHO] T= 298.15 K, 1=0.23 0.23 HC1 pot -8.30 ' [1918NOY/CHO] T= 298.15 K, 1=0.5 0.5 HC1 pot values calculated in this report from the E° values and HC1 concentrations given by [1918NOY/CHO]. 269 JNC TN8400 99-011 + Bi(cr)+Cr+H2O o BiOCI(s)+3e+2H 0.2 0.4 0.6 lm> molal Figure 7.25: Plot of log K*s - ID vs. Im for the reaction Bi(cr) + Cl~ + H2O <=> BiOCl(s) + 2H+ + 3e~ at 25 °C. The straight line shows the result of the linear regression: Ae = -0.38; log K*°s = -8.32. Calculated with the data given in Table 7.27. This value can be corrected for the reaction Bi(cr) <=> Bi3+ + 3e- (log K° = 16.74; see 3+ Section 7.11) to a log K*So° of 8.42 for the reaction Bi + H2O+ Cl~ <=> BiOCl(s) +2H+. [1951SWI/GAR]: [1951SWI/GAR] determined the solubility of BiONO3(s) in HNO3. As the log K*s0 of 2.55 given by [1951SWI/GAR] for 1=0 was criticized by [1993KRA/DEC], this value was recalculated for the present report from the data given in [1951SWI/GAR] (Table 7.27 and Figure 7.26). The Bi3+ concentration was 2+ corrected in this report for the formation of BiNO3 (Using the log B j j determined in Section 7.7.1). Extrapolation of the values compiled in Table 7.28 to I = 0 as shown in Figure 7.26 result in: 3+ Bi + H2O + NO3- BiONO3(s) + 2H+ log Kso*° = 2.75 This value compares well with the value determined by [1993KRA/DEC] (see Section 7.7.2: BiONO3(s)). 270 JNC TN8400 99-011 Table 7.28: Experimental equilibrium data compiled for the precipitation of BiONO3(s). These data were chosen for the evaluation of recommended values in the present report. Method: sol = solubility. Reference Comments I (M) Medium Method log KSo 3+ log Kso: Bi + H2O + NO3- <^BiONO3(s) + 2H+ 2.18 1 [1951SWI/GAR] T= 298.15 K, 1=0.001 0.025 HNO3 sol ] 2.20 [1951SWI/GAR] T= 298.15 K, 1=0.03 0.03 HNO3 sol 2.17 1 [1951SWI/GAR] T= 298.15 K, 1=0.035 0.035 HNO3 sol 2.15 1 [1951SWI/GAR] T= 298.15 K, 1=0.04 0.04 HNO3 sol 2.13 1 [1951SWI/GAR] T= 298.15 K, 1=0.046 0.046 HNO3 sol 2.09 1 [1951SWI/GAR] T= 298.15 K, 1=0.05 0.05 HNO3 sol 2.05 1 [1951SWI/GAR] T= 298.15 K, 1=0.06 0.06 HNO3 sol 2.00 ! [1951SWI/GAR] T= 298.15 K, 1=0.07 0.07 HNO3 sol 1.94 ] [1951SWI/GAR] T= 298.15 K, 1=0.08 0.08 HNO, sol 1 2+ recalculated in this report including BiNO3 complex formation 3+ + Bi +H2O+NO3- o BiONO3(s)+2H 4-5 -- 4 -: 00 3.5 -~ g> 2 + 1.5 - 1 0.5 +. 0 0 0.02 0.04 0.06 0.08 0.1 lm, molal 3+ Figure 7.26: Plot of log K*So + 8D vs. Im for the reaction Bi + H2O + NO3- o BiONO3(s) + 2H+ at 25 °C. The straight line shows the result of the linear regression: Ae = 0.34; log K*°s = 2.75. Calculated with the data given in Table 7.28. [1957OLI]: [Bi] = 0.1 - 50 mM. Olin ([1957OLI], [1959OLI], [1961OLI], [1975OLI]) was one of the first investigators who measured the hydrolysis of Bi(III) in 0.1 and 3 M 271 JNC TN8400 99-011 perchlorate medium. Based on potentiometric measurements he reported for the formation of mononuclear BiOH2+ a log pi] value of -1.58 and for the polynuclear 6+ Bi6(OH)i2 a log p6,i2 value of 0.33 in 3 M NaClO4. His careful experimental work encompassed a large range of pH values and bismuth concentrations and he developed a consistent system for the hydrolysis of bismuth. The values determined graphically by [1957OLI] were later recalculated with a computer fitting program by [1975OLI]. [1959OLI]: Bi = 0.25 - 4 mM in 0.1 M NaC104. For further comments see [1957OLI]. 3+ [1960TOB]: [1960TOB] proposed the existence of a Bi6(OH)]5 complex (Table 7.2). Spectrophotometric measurements made by [1972DRA/NIM1] are consistent with 7+ 6+ 5+ the formation of polynuclear Bi9(OH)2o , Bi9(OH)2i , and Bi9(OH)22 complexes as proposed by [1959OLI]. From the evaluation of his data, [1960TOB] also calculated log p6,l2 value of -0.53 6+ for the formation of Bi6(OH)]2 . However, as [1960TOB] did not seem to take into 2+ + account any other complexes (e.g. BiOH or Bi(OH)2 ) his data were not chosen in this report. [1960TOB] made also an attempt to recalculate the data given in [1957OLI] and found a log (36,i2 value of +0.03 (instead of the 0.33 as given by [1957OLI]). [1960TOB] writes that the cause of the discrepancy between the values of the constants obtained from Olin's data is not known. However, the neglect of mononuclear complexes by [1960TOB] may be reason for this discrepancy. [1961OLI]: Gives the results already published in [1957OLI] and [1959OLI]. For further comments see [1957OLI]. [1971BID]: [197IBID] measured in 0.1 and 1 M perchlorate solution the bismuth hydrolysis with an organic extractant (dithizone). He was able to show that no (or only little) polynuclear species were present in 0.001-0.01 mM bismuth solutions in the pH range 9.9 - 11.3. [1971BID] used a log Kw of -13.79 in 0.1 M perchlorate solutions. [1972DRA/NIM1]: Bi = 0.4-10 mM. Dragulescu and co-workers ([1972DRA/NIM1], [1972DRA/NIM2]) determined bismuth hydrolysis in 0.1 and 1 M perchlorate medium. Their results are comparable to the results of Olin. [1972DRA/NIM2]: cf. [1972DRA/NIM1]. [1975ANT/NEV]: [1975ANT/NEV] measured bismuth hydrolysis in 0.1, 0.3, 0.5 and 1 M KNO3 spectrophotometrically with an organic ligand. As bismuth tends to form weak complexes with nitrate these values were not included in our calculations. [1975ANT/NEV] excluded in presence of an organic ligand (l-(2-pyridylazo)-2- naphthol) the formation of polynuclear Bi species in their 0.02 mM Bi solutions. 272 JNC TN8400 99 - Oil [1975HEI/SCH]: [1975HEI/SCH] determined potentiometrically in 1.5 M H2SO4 the redox potential of the Bi3+/Bi(cr) and BiO+/Bi(cr) couples. [1975HEI/SCH] tried to correct for complex formation with sulfate. However, the results are difficult to interpret and are not used for extrapolation. [1975OLI]: Recalculation of the results already published in [1957OLI] and [1959OLI]. For further comments see [1957OLI]. [1982HAT/SUG]: [1982HAT/SUG] measured in 1 M perchlorate and nitrate solutions the hydrolysis of bismuth with an organic extractant (dithizone) in ~ 0.0001 mM bismuth solutions. [1982HAT/SUG] assumed that the difference between their results and the result of [1972BID] may be due to the formation of polynuclear species. [1982LAP/KOL]: The values reported by [1982LAP/KOL] are linearly extrapolated to 25 [1987MIL/ROE]: [1987MHVROE] determined the hydrolysis of Bi(III) in 0.25 perchlorate solutions and at bismuth concentration of 10~12 M. [1987SUG/ISH]: [1987SUG/ISH] measured in 1 M chloride solutions the hydrolysis of bismuth with an organic extractant (dithizone) in = 10~n M bismuth solutions. Due to the formation of chloride complexes the determined hydrolysis constants were smaller than the ones determined earlier by the same group ([1982HAT/SUG]) indicating a quite strong complex formation of bismuth with chloride. The measurements are not used in this report for extrapolation with the SIT method. [1987SUG/ISH] give in 1 M chloride a log Kw of 13.62. [1993KRA/DEC]: [1993KRA/DEC] determined in 1 M perchlorate and nitrate solutions the hydrolysis of bismuth based on solubility measurements (oversaturation). The hydrolysis constants do not fit well with other data found in the literature as [ 1993 KRA/DEC] did not consider the possible formation of the polynuclear 7+ 6+ 5+ Bi9(OH)2o , Bi9(OH)2i , and Bi9(OH)22 complexes in their calculations. 2+ [1993KRA/DEC] reported also a log pij of 3.5 for the formation of BiC104 . However, this equilibrium constant seems to be much too large in comparison with the constants of the bismuth complexes with nitrate and chloride as given in this report. [1993KRA/DEC] present in their report no direct experimental proof for the predominance of the BiC1042+ (instead of Bi3+) at pH below 2. Additionally, [1982HAT/SUG] and [1987SUG/SHI] showed that nitrate and perchlorate had only a small influence on bismuth hydrolysis while the influence of chloride (and thus also the complex formation with chloride) is more important. In any case, the possible 273 JNC TN8400 99-011 existence of BiC1042+ has no influence on the hydrolysis constants of Bi(OH)3 and Bi(OH)4~ determined at much higher pH values. [1993KRA/DEC] also determined constants for the complex formation of Bi(III) with nitrate and the precipitation of BiONO3(s). These constants do not depend on the presence or absence of the polynuclear Bi9(OH)2o7+, Bi9(OH)2i6+, and Bi9(OH)225+ complexes and are therefore included in the present report. 274 JNC TN8400 99-011 8 Niobium A number of oxidation states ranging from -I to +V and +VII have been mentioned in the review of [1985UDU/VEN]. Thermodynamic data, however, exist only for compounds with the oxidation states +11, +IV, and +V. Only Nb(V) forms aqueous species. Aqueous alkali- 8 m metal niobate solutions contain the hexaniobate anion HmNb6Oi9( ~ )- at pH > 7 in concentrated solutions [1983POP]. Niobium also forms pentahalogenides with Cl, F, Br and I [1985UDU/VEN, 1995WIB]. 8.1 Hydrolysis of niobium(V) The solubility measurements of [1963BAB/LUK], [1992YAJ^TOB] and [1994YAJ] (see Section 8.1.1), indicate the presence of a singly, negatively charged species in the pH range 7 to 10, while they seem to exclude the presence of higher charged species. As neither [1963BAB/LUK], nor [1992YAJ/TOB and 1994YAJ] determined the predominant species under alkaline conditions, the simplest assumption is that Nb(OH)6~ is the predominant species under these conditions. For the purpose of modeling, only Nb(OH)6~ is used in this report. However, there is spectrophotometric evidence for the existence of hexaniobate anions: H3Nb6Oi95~, H2Nb6Oi96~, HNb6Oi97~, and NbgO^8-, in saturated niobium solutions in the pH range 9-14 [1964NEU, 1968SPI, 1994ETX/FER]. An overview of the data available in literature concerning polynuclear Nb(V) species is given in section 8.1.2. 8.1.1 Monomeric Nb(V) species [1963BAB/LUK] interpreted their measurements of the solubility of freshly precipitated Nb2O5(s) in terms of the monomeric species Nb(OH)4+, Nb(OH)5° and Nb(OH)6-. As this was the only determination of the solubility of Nb2C>5(s) until the very recent measurements of [1992YAJ/TOB] and [1994YAJ], these values were reported in several compilations. [1963BAB/LUK], however, did not indicate any detection limit and [1976BAE/MES] considered the values given by [1963BAB/LUK] as approximate estimates. [1963BAB/LUK, 1992YAJATOB and 1994YAJ] determined the solubility of freshly precipitated M^OsCprecip) in alkaline solutions and fitted their data using Nb(OH)6~ as dominant species (Tables 8.1 and 8.2). Their solubility measurements indicate the predominance of a singly charged species in the pH range 7 to 10. There is, however, an inconsistency between assuming the predominance of Nb(OH)6~ and the observations of [1960JAN/ERT, 1964NEU, 1968SPI, and 1994ETX/FER] (cf. Section 8.1.2) who observed the presence of the hexaniobate ion in alkaline solutions. Based on the solubility experiments of [1963BAB/LUK, 1992YAJATOB and 1994YAJ] the predominance of a higher charged polymeric species is uncertain and for the purpose of modeling, only Nb(OH)g" is assumed in this report. 275 JNC TN8400 99 -Oil [1963BAB/LUK] estimated a log KS6 of 24.5 (Table 8.2) in 1 M KNO3. More recently, [1992YAJ/TOB] and [1994YAJ] extrapolated from measurements in 0.1 M NaCl a log K°s6 of 29.2 at 1=0 (Table 8.1) for the reaction: 2Nb(OH)6- + 2H+ <=> Nb2O5(precip) + 7H2O log K 29.2 S6 Table 8.1: Experimentally determined stepwise formation constants K compiled for the niobium(V) hydroxide system. These data were chosen for the evaluation of recommended values in the present report. Additional information for the different references see Section 8.5: 'Comments on selected references'. Method: sol = solubility measurements, tit = titration (pH). Kb is the stepwise formation constant for a reaction using OH~~ as a component. log Kb Reference Comments I (M) Medium Method log K*S6: 2Nb(OH)6 Nb2O5(precip) + 7H2O 29.2 1 [1992YAJ/TOB], T= 298.15 K, 1=0.1 0 NaCl sol [1994YAJ] b 8 7 log K 6.i9.j: Nb6O19 ~+ H2O & Nb6O19H ' + OH~ -0.36 2 [1964NEU] T= 298.15 K, 1=3 3 KC1 tit -1.83 3 [1968SPI] T= 298.15 K, 1=0.5 0.5 KC1 tit -1.7 3 [1968SPI] T= 298.15 K, 1=1 1 KC1 tit -1.56 3 [1968SPI] T= 298.15 K, 1=2 2 KC1 tit -1.4 3 [1968SPI] T= 298.15 K, 1=3 3 KC1 tit -0.49 4 [1994ETX/FER] T= 298.15 K, 1=3 3 KC1 tit 7 6 log K\19i2: Nb6O]9H ~+ H2O t=> Nb6O19H2 ~ + OH~ -3.28 2 [1964NEU] T= 298.15 K, 1=3 3 KC1 tit -3.03 3 [1968SPI] T= 298.15 K, 1=0.5 0.5 KC1 tit -3.06 3 [1968SPI] T= 298.15 K, 1=1 1 KC1 tit -3.09 3 [1968SPI] T= 298.15 K, 1=2 2 KC1 tit -3.143 [1968SPI] T= 298.15 K, 1=3 3 KC1 tit -4.2 4 [1994ETX/FER] T= 298.15 K, 1=3 3 KC1 tit 276 JNC TN8400 99-011 Table 8.1: continued b 6 5 log K 6j9>3: Nb6O]9H2 - •+ H20 <^> Nb6O]9H3 ~ + OH- -4.19 3 [1968SPI] T= 298.15 K, 1=0.5 0.5 KC1 tit -4.29 3 [1968SPI] T= 298.15 K, 1=1 1 KC1 tit -4.46 3 [1968SPI] T= 298.15 K, 1=2 2 KC1 tit -4.61 3 [1968SPI] T= 298.15 K, 1=3 3 KC1 tit -4.77 4 [1994ETX/FER] T= 298.15 K, 1=3 3 KC1 tit 1 corrected to 1=0 with Davies equation by [1992YAJ/TOB, 1994YAJ] 2 Nb concentration = 50-140 mM 3 Nb concentration = 2.5-20 mM 4 Nb concentration = 2.5-25 mM 8.1.2 Polymeric Nb(V) species Aqueous alkali-metal niobate solutions (Nb concentration = 1-100 mM) appear to contain the 8 m hexaniobate anion HmNb6019( - )- at pH > 9 [1964NEU, 1968SPI, 1994ETX/FER, 1983POP]. Such solutions are prepared by fusion of Nb2Os with excess metal hydroxide or carbonate, and subsequent dissolution of the melt in water [1960JAN/ERT2]. Also direct dissolution of freshly prepared Nb2Os in water is possible [1983POP]. Ion diffusion in water [1960JAN/ERT1] and spectrophotometric measurements [1960JAN/ERT1, 1973GOI//GRA] are consistent with the hexaniobate anion being the predominant species in saturated aqueous solutions at pH > 9. There is good evidence from potentiometric [1960JAN/ERT3, 1964NEU, 1994ETX/FER, 1968SPI] and spectrophotometric studies [1973GOI//GRA, 1974GOI//SPI] that NbgOig8- is protonated below pH = 14 in solution (see Table 8.1). An overview of the data available in literature is given in the following sections. The data used for the calculations of equilibrium constants for the niobium(V) hydroxide system are given in Table 8.1 and Figures 8.1, 8.2 and 8.3. Additional data are compiled in Table 8.2 and 8.3. 8 7 6 5 8.1.2.1 Profanation at pH > 8: Nb6Oj9 ~, Nb6O19H ~, Nb6O]9H2 -, and Nb6O19H3 ~ 8 Protonation constants for Nb6O]9 - have been determined by [1964NEU], [1968SPI], and [1994ETX/FER] by potentiometric measurements in 0.1 - 3 M KC1 (see Table 8.1). [1968SPI] extrapolated his results graphically to 1=0 (Table 8.2). [1994ETX/FER] extrapolated their results and the results of [1964NEU, 1968SPI] to 1=0 with an extended Davies equation. However, for the determination of the first protonation constant (log K^i^i: Nb6O]98~ + H+ 7 <=» Nb6O19H -) [1994ETX/FER] did not include the data of [1968SPI], making their extrapolation from I = 3 to I = 0 somewhat doubtful. [1964NEU, 1968SPI, and 8 1994ETX/FER] all determined Nb6Oi9 ~ protonation as a function of OH~ concentration and then converted their constants to hydrolysis constants based on H+ with different log Kw values. [1968SPI] used for all ionic strengths a log Kw of-14.0. Therefore, in this report for 277 JNC TN8400 99-011 extrapolation to 1=0, the original stepwise formation constants Kb (which refer to the reaction with OH~) were used as given in Table 8.1. Extrapolation of the stepwise formation constants given in Table 8.1 to I = 0 is shown in Figures 8.1 -8.3: 8 7 Nb6O19 -+H+ Nb6Oi9H - log K°6,i9,i = 14.29 7 6 Nb6Oi9H - + H+ Nb6Oi9H2 - logK°6,19,2= 13.23 6 5 Nb6Oi9H2 - + H+ Nb6Oi9H3 - logK°6,19l3= 11-63 These values, however, are considered as tentative because the solubility measurements indicate the predominance of a singly charged (but not necessarily mononuclear) species in the pH range 7 to 10 (cf. Section 8.1.1). 7 <=> Nb6O19H -+OH- 3.5 i 2 -I 1.5 §&Bi^^^^^ b 8 Figure 8.1: Plot of log K 6,i9,i + 14 D vs. Im for the reaction Nb6O19 - + H2O <^> 7 Nb6Oi9H - + OH- at 25 °C. The straight line shows the result of the linear b regression: Ae = -0.73; log K °6,i9,i = 0.29; log K°6;19ii = 14.29 for the reaction 8 + 7 Nb6Oi9 - + H <^> Nb6O19H - at 25 °C. Calculated from data in Table 8.1. 278 JNC TN8400 99-011 7 Nb6O19H +H2O«Nb6O1 c ;1 : ; : : : : " , . <-v \.. 1.5 • 1 : : ]: :i i 1 • •;•-.:• •.•.":;;.•:'>:.:.:i. \. 12 U 0.5 : : C'"J ; -i••••'. •• .-l. .:ii:^.; ;; h.::::V. ' '• :;•:• = .:•"•:;:/1' v:-. - '•'..:. •:'„'."•w'yi-'^-'. 0 ! •• •::':^.. }\y:::^-'\-:-;\:_':_ -0.5 • 6,19, 2 + ... •'"':::'—-i!!';; j".'!. •;::^f1 -1 • cn ;'•••• --::v iV":-^^.!1;-!::^;"* - -1.5 - iliiii -2 -2.5 (} 1 2 3 4 lmi molal b 7 Figure 8.2: Plot of log K 6,i9i2 + 12D vs. Im for the reaction Nb6Oi9H - + H2O o 6 Nb6Oi9H2 - + OH- at 25 °C. The straight line shows the result of the linear b regression: Ae = -0.08; log K °6,i9,2 = -0.77; log K°6,i9i2 = 13.23 for the 7 6 reaction Nb6Oi9H - + H+ «• Nb6Oi9H2 - at 25 °C. Calculated from data compiled in Table 8.1. 6 & Nb6O19H2 +H2O<^Nb6O19H3 -+OH- -0.5 Q -1 o -1.5 + -2 • -2.5 - -3 O -3.5 • -4 • •4.5 - -5 - 1 2 lm, molal b 6 Figure 8.3: Plot of log K 6,i9,3 + 10D vs. Im for the reaction Nb6Oi9H2 - + H2O <=» 5 Nb6Oi9H3 - + OH" at 25 °C. The straight line shows the result of the linear b regression: Ae = -0.06; log K °6,i9,3 = -2.37; log K°6,i9,3 = 11.63 for the 6 5 reaction Nb6O19H2 - + H+ <=> Nb6Oi9H3 - at 25 °C. Calculated from data compiled in Table 8.1. 279 JNC TN8400 99-011 8 12 3 8.1.2.2 Very alkaline solutions: Nb4O]2(OH)4 ~, Nb4O16 ~ and NbO2(OH)4 ~ On the basis of Raman and UV spectroscopy, [1973GOI//GRA] and [1974GOI//SPI] suggested that in very alkaline solutions (1-12 M KOH) NbgOi98- is degraded into tetrameric and 8 12 3 monomeric species: Nb4Oi2(OH)4 -, Nb4Oi6 ~ and NbO2(OH)4 - (Table 8.2). The proportions of the latter two species depend upon niobium concentration. As I is not constant in these experiments (1-10 M KOH), extrapolation to I = 0 is not possible. These species are not important under the pH conditions of natural waters. 6 7 8 8.1.2.3 Neutral and acidic solutions: H6Nbj2036 ', H5Nb12036 ~, H4Nb]2036 ~, and Different authors proposed further polymerization of the Nb^O^^5" species in neutral and acidic solution. [1960JAN/ERT1] found evidence from diffusion measurements in water for the formation of a polymer with roughly the threefold molecular mass of NbgO^Hs5- at pH < 8. 6 7 [1968SPI] proposed the presence of dodecaniobates (HeNhnC^ ", H5Nbi2O36 ~, 8 9 H4Nb]2O36 ~, H3Nbi2C>36 ~) and determined by pH titration constants for the formation of 7 8 9 H5Nb]2O36 -, H4Nbi2O36 -, H3Nbi2O36 - in 1 M KC1 solutions (Table 8.2). However, it cannot be excluded that the solutions were oversaturated with respect to Nb2Os(s). Thus, no direct proof exists of the presence of dodecaniobates in equilibrium with 8.1.3 Additional equilibrium data compiled for the niobium(V) hydroxide system Table 8.2: Additional, experimentally determined equilibrium data compiled for the niobium(V) hydroxide system. These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility measurements, sp = spectrophotometry, and tit = titration (pH). b log K m,n0 Reference 'Comments I (M) Medium Method + log K*S6.- 2Nb(OH)6~ + 2H <=>Nb205(precip) + 7H2O 24.5 ' [1963BAB/LUK1 T= 298.15 K, 1=1 1 KNO, sol + + log Kl4 : Nb(OH)5° + H <=> Nb(OH)4 + H2O -0.6 ' [1963BAB/LUK] T= 292 K, 1=1 1 KNO, sol log KIi6 .• Nb(OH)5° + H2O » Nb(OH) -7.40 ' [1963BAB/LUK] T=292K, 1=1 1 KNO, sol 280 JNC TN8400 99-011 Table 8.2: continued b 8 7 log K 6J9J: Nb6O19 -<=>Nb6O,9H -+ OH~ -2.02 2 [1968SPI] T= 298.15 K, 1=0.1 0.1 KC1 tit -2.1 3 [1968SPI] T= 298.15 K, 1=0 0 KC1 tit 2.11 4 [1994ETX/FER1 T= 298.15 K, 1=0 0 KC1 tit h 7 log K 6J92: Nb6O!9H ~<=> Nb6O19H2'*-+ OH- -2.99 2 [1968SPI] T= 298.15 K, 1=0.1 0.1 KC1 tit -3 3 [1968SPI] T= 298.15 K, 1=0 0 KC1 tit -2.14 4 R994ETX/FER1 T= 298.15 K, 1=0 0 KC1 tit 6 log K\19J: Nb6O19H2 -<=>Nb6O19H /- + OH- -4.06 2 [1968SPI] T= 298.15 K, 1=0.1 0.1 KC1 tit -43 [1968SPI] T= 298.15 K, 1=0 0 KC1 tit -2.06 4 ri994ETX/FERl T= 298.15 K, 1=0 0 KC1 tit b 8 8 log K 4il2A: 2Nb6O,9 -+ 8OH~ <=> 3Nb4O12(OH)4 ~ s -0.61 [ 1973GOI/GRA] T= 298.15 K, 1= 1 -10 1MKC1, 1-10 MKOH sp b log K h6A: Nb4O12(0H)/-+ 40H~ + 4H2O « NbO2(OH)/- 5 -11.96 [1973GOI7GRA1 T= 298.15 K, 1=1-10 1 M KC1, 1-10 M KOH sp b 2 log K uzo: Nb4OJ2(OH)/-+ 40H~<=> Nb4O16' - + 4H2O 5 -3.23 ri973GOI/GRA1 T= 298.15 K, 1=1-10 1 MKC1, 1-10 MKOH sp 55 + 8 log KiJ2M: H3Nb601919 - + H 0.5H4Nb12O36 ~+ H2O 7.0 6 H968SPI] T= 298.15 K, 1=1 1 M KC1 tit 8 9 + log K3J2i36: H4Nb12036 ~ <=> H3Nb12036 -+ H 6 -7.8 [1968SPI] T= 298.15 K, 1=1 1 M KC1, tit 7 8 + log K4J2J6: H5Nb12036 - «- H4Nb12O36 - + H -6.34 6 [1968SPI1 T= 298.15 K, 1=1 1 1 M KC1, tit ' Detection limit not determined, determined in presence of freshly precipitated Nb2O5 (s). 2 1 probably not constant (which is rather difficult at 1=0.1 and pH 13). 3 extrapolated graphically to 1=0 by [1968SPI] 4 extrapolated I = 0 by [ 1994ETX/FER] (see also Section 8.1.1) 5 I not constant 6 oversaturation possible 281 JNC TN8400 99-011 Table 8.3: Thermodynamic data for the niobium(V) hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. b log K 619 m Reference Comments I (M) Medium b 8 7 log K 619J: Nb6O19 -<=>Nb6O19H - + OH~ -0.36 [1976BAE/MES] T= 298.15 K, 1=3 3 KC1 -0.36 fl976SMI/MARl T= 298.15 K, 1=3 3 KC1 7 log K\J9X. Nb6O19H -&Nb6O,9H2fi- + OH- -3.28 [1976BAE/MES] T= 298.15 K, 1=3 3 KC1 -3.28 [1976SMI/MAR1 T= 298.15 K, 1=3 3 KC1 + + log KlA. Nb(OH)5° + H <=> Nb(OH)4, + H2O -0.6 [1976BAE/MES] T= 292 K, 1=1 1 KNO3 -0.45 [1982WAG/EVA] T= 298.15 K, 1=1 1 -0.59 [1985CHA/DAV] T= 298.15 K,I=n/a -0.44 [1985UDU/VEN] T= 298.15 K,I=n/a -0.44 [1992PEA/BER1 T= 298.15 K,I=n/a + log KL6: Nb(OH)5° + H2O <=> Nb(OH)6~ + H -7.40 [1976BAE/MES] T= 292 K, 1=1 1 KNO3 -7.34 [1982WAG/EVA] T= 298.15 K, 1=1 1 -7.40 [1985CHA/DAV] T= 298.15 K,I=n/a -7.33 [1985UDU/VEN] T= 298.15 K,I=n/a -7.34 [1992PEA/BER1 T= 298.15 K, I=n/a 282 JNC TN8400 99-011 5.2 Solubility of solid niobium pentoxide Freshly precipitated Nb2Os(precip) has an amorphous structure. Crystallization to Nb2C>5(cr) occurs upon heating around 700 °C when all water molecules are removed [1989VAN/POU]. The data used for the calculations of the solubility of NbiOsCprecip) are given in Table 8.4. Additional data that were not chosen for the evaluation in this report are compiled in Tables 8.5 and 8.6. [1963BAJB/LUK] determined the solubility of freshly precipitated Nb2Os(precip) in 1 M KNO3. In the neutral pH range, [1963BAB/LUK] determined a minimum solubility of 1.4 x 10~5 M Nb, corresponding to a logK*so of 4.85 for the theoretical equilibrium Nb(OH)5° <=> 0.5Nb2O5(precip) + 2.5H2O (Table 8.5). However, as no detection limit is indicated in the work of [1963BAB/LUK], one may suspect that the measured concentrations below pH 7 correspond to the detection hmit and that the real concentration may be smaller. Recently, [1992YAJ/TOB] and [1994YAJ] determined in 0.1 M NaCl a minimum solubility < 1 x 10~8 M Nb, corresponding to a logK*so > 8 for the theoretical equilibrium Nb(OH)5° o 0.5Nb2C>5(precip) + 2.5H2O in both over- and undersaturation experiments (Table 8.4). From these data and the discussion concerning the hydrolysis of niobates (Section 8.1), it can be concluded that in the neutral pH range, niobium solubility is smaller than 1 x 10"8 M. Table 8.4: Experimentally determined equilibrium data for the dissolution of These data were chosen for the evaluation of recommended values in the present report. Method: sol = solubility measurements. log K*so Reference Comments I (M) Medium Method log K*S5: 2Nb(OH)5° <^Nb2O5(precip) + 5H2O > 16.0 1 [1992YAJ/TOB], T= 298.15 K, 1=0.1 0 NaCl sol [1994YAJ] 1 corrected to 1=0 with Davies equation by [1992YAJ/TOB, 1994YAJ] It is not possible, due to the lack of experimental data, to give an exact thermodynamic constant for the solubility of niobium in neutral or acidic solutions. However, an upper solubility limit is known from the experiments of [1992YAJ/TOB] and [1994YAJ] and for modeling purposes it is practical to express this upper solubility limit in the neutral and acidic pH range through an equation involving the uncharged, hypothetical Nb(OH)s° as 'dummy species'. Using the data given in [1992YAJ/TOB] and [1994YAJ] one calculates: 283 JNC TN8400 99-011 2Nb(OH)5° <=> Nb2O5(precip) + 5 H2O log K*S5= 16.0 From this also a constant for the equilibrium between the hypothetical species Nb(OH)s° and Nb(OH)6- (cf. Section 8.1.1) can be calculated + Nb(OH)5° + H2O <=> Nb(OH)6- + H log K*5,6 = -6.6 It should be stressed that the 'species' Nb(OH)$0 (and possibly Nb(OH)<5~) is an hypothetical species which allows one to calculate a solubility limit for niobium in the neutral and acidic range but does not necessarily correspond to a real species. [ 1993KUL/HAK] determined the solubility of niobium in groundwater in the pH range 6 to 13. The results of their overs aturation experiments show an increase of Nb solubility in the pH range 6 - 10 and are consistent with the results of [1963BAB/LUK, 1992YAJ/TOB and 1994YAJ]. At pH 10 - 13, [1993KUL/HAK] observed no further pH dependency of niobium solubility (« 10-3 to 10"2 M). Similarly, [1992YAJ/TOB] and [1994YAJ] also observed at pH > 10 a rather constant Nb solubility in the range of 10-4 to 10~3 M. In undersaturation experiments using Nb2O5(cr?) [1993KUL/HAK] observed after 4 months a niobium solubility in the range of <1X10~9 M (ultrafiltration experiments) to 5x10~8 M (centrifugation experiments) at pH 8 and 13. [1990PIL/STO] determined in undersaturation experiments in concrete water (pH 11.8 - 12.5) a niobium solubility of <2xlO~7 M (after 1 week) to 2x10-3 M (after 18 month). The results of [1990PIL/STO] show that the dissolution of commercially available (probably calcined) Nb2Os(cr?) is a much slower process than the dissolution of freshly prepared Nb2O5(precip). [1992YAJ/TOB] and [1994YAJ], on the other hand, found after four weeks equilibration time no difference between the solubility of Nb2C>5(cr) in undersaturation experiments and the solubility of precipitated Nb2O5(am) in oversaturation experiments. The difference between the observations of the three undersaturation studies is probably due to the different Nb2Os(cr) used. Unfortunately, no detailed information can be found regarding the nature of the solids used by [1993KUL/HAK] and [1990PIL/STO]. 284 JNC TN8400 99-011 8.2.1 Additional equilibrium data compile Table 8.5: Experimental determined equilibrium data compiled for the precipitation of Nb2O5. These data were not chosen in the present report for the evaluation of recommended stability values. Method: sol = solubility measurements. log K*so Reference Comments I (M) Medium Method log K*S5: 2Nb(OH)5° <=> Nb2O5(precip) + 5H2O 9.7 ' [1963BAB/LUK] T= 298.15 K, 1=1 1 KNO, sol 1 Detection limit not determined, freshly precipitated Nb2O5 (s). Table 8.6: Thermodynamic data for the precipitation of Nb2O5 taken from previous compilations. Precipitation of Nb2O5(precip) according to the hypothetical equilibrium (see text) 2Nb(OH)5° <=> Nb2O5(precip) + 5H2O. As pointed out in Section 2 of this report only experimental data were chosen for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log rso Reference Comments KM) Medium log K*S0: 2Nb(OH)5° <=> Nb2O5(precip) + 5H2O 9.6 [1976BAE/MES] T=292K, 1=1 1 KN03 9.66 [1982WAG/EVA] T= 298.15 K, 1=1 1 10.06 [1985UDU/VEN] T= 298.15 K,I=n/a 9.66 [1992PEA/BER] T= 298.15 K,I=n/a 285 JNC TN8400 99-011 8.3 Solid niobium phases: redox equilibria Niobium pentoxide, Nb2O5(cr), is a dense white powder and chemically relatively inert. A comprehensive review of the thermochemistry and oxidation potential of niobium has been made by [1971HIL/WOR] and [1985UDU/VEN]. Nb(cr) is insoluble in water, but reacts in presence of oxygen above 300 °C to the pentoxide Nb2O5(s). Also the solids NbO(s) and NbO2(s) are known. Both solids are thermodynarnically unstable in the presence of water where the pentoxide Nb2Os(s) is formed under reduction of water ([1985UDU/VEN], [1995WTB]). From the thermodynamic data given in [1985UDU/VEN] and other compilations equilibria between the different solid niobium phases can be calculated (Table 8.7). Table 8.7: Thermodynamic data for the redox equilibrium of Nb(cr), NbO and NbO2 with Nb2O5 taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. log Pi,0 Reference Comments l : °g Pl,0 0.5Nb2O5 + 5e~ + 5H+ <=> Nb(cr)+ 2.5H2O -51.12 [1954COU] T= 298.15 K, I=n/a -51.18 [1963WIC/BLO] T= 298.15 K, I=n/a -51.02 [1971HIL/WOR] T= 298.15 K, I=n/a -50.65 [1971NAU/RYZ] T= 298.15 K, I=n/a -50.82 [1978ROB/HEM2] T= 298.15 K, I=n/a -50.82 [1979KUB/ALC] T= 298.15 K, I=n/a -50.82 [1982PAN] T= 298.15 K, I=n/a -50.83 [1982WAG/EVA] T= 298.15 K, I=n/a -50.82 [1985CHA/DAV] T= 298.15 K, I=n/a -51.01 ri985UDU/VENl T= 298.15 K, I=n/a l °i PJ.O- 0.5Nb2Os + 3e~ + 3H+ <=> NbO+ 1.5H2O -6.66 [1952WOR] T= 298.15 K, I=n/a -4.30 [1954COU] T= 298.15 K, I=n/a -4.38 [1963WIC/BLO] T= 298.15 K, I=n/a -4.30 [1971HIL/WOR] T= 298.15 K, I=n/a -4.41 [1978ROB/HEM2] T= 298.15 K, I=n/a -4.41 [1979KUB/ALC] T= 298.15 K, I=n/a -4.19 [1982PAN] T= 298.15 K, I=n/a -4.19 [1982WAG/EVA] T= 298.15 K, I=n/a -4.41 [1985CHA/DAV] T= 298.15 K, I=n/a -4.30 [1985UDU/VEN1 T= 298.15 K, I=n/a + log A.o- 0.5Nb2O5 + e~ + H <^> NbO2+0.5H2O -25.82 [1971HILAVOR] T= 298.15 K, I=n/a -23.70 [1978ROB/HEM2] T= 298.15 K, I=n/a -23.70 [1979KUB/ALC] T= 298.15 K, I=n/a -26.01 [1982PAN] T= 298.15 K, I=n/a -26.05 [1982WAG/EVA] T= 298.15 K, I=n/a -23.70 [1985CHA/DAV] T= 298.15 K, I=n/a -26.23 [1985UDU/VEN] T= 298.15 K, I=n/a 286 JNC TN8400 99-011 5.4 Other niobium( V) complexes and compounds A few measurements concerning the formation of Nb(V) fluoride complexes are reported in the literature (Table 8.8). Only the presence of mononuclear Nb(V) is assumed in these papers. No thermodynamic values for the formation of Nb(V) fluorides are recommended in this report. The standard Gibbs energy of formation, AfG°, of the solids NbCl5(s), NbOCl3(s), NaNbO3(s), Na3Nb04(s), KNbO3(s), K3Nb04(s), and Ca(NbO3)2(s) are compiled in Table 8.9. No solubility products are calculated from these values and no thermodynamic values for the formation of these Nb(V) solids are recommended in this report. Table 8.8: Experimentally determined equilibrium data compiled for the antimony(V) system. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text in Section 8.4: Method: pot = potentiometry. log (3 Reference Comments I (M) Medium Method + log KUJ: NbOF2 + F~ <=> NbOF3° 3.78 ri972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HC1O4 pot log KUA: NbOF3° + F~ <=> NbOFf 4.30 [1972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HCI04 pot 2 log KUi5: NbOF4- + F~ <=> NbOF5 ' 2.51 [1969NEU] T= 298 K, 1=3 3.0 KC1 pot 4.51 [1972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HC1O4 pot 2 3 log K1.1.6: NbOF5 ~ + F- <=> NbOF6 ~ 4.67 [1972LAN/OSB1 T= 298 K, 1=0.5 0.50 Na, HCIO,, pot K : 3 + 2 log i.o,7 NbOF6 ~ + F~ + 2H <=> NbF7 vH~ A2O 11.41 [1972LAN/OSB1 T= 298 K, 1=0.5 0.50 Na, HCIO, pot 2 3 log Kli0_8: NbF7 ~ + F~ <=> NbF8 - 3.08 [1972LAN/OSB] T= 298 K, 1=0.5 0.50 Na, HC1O4 pot 4 log K109: NbF/- + F~ <=> NbF9 - 4.0 [1972LAN/OSB1 T= 298 K, 1=0.5 0.50 Na, HCIO, pot 287 JNC TN8400 99-011 Table 8.9: Standard Gibbs energy of formation, A{G°, of the solids NbCl5(s), NbOCl3(s), NaNbO3(s), Na3Nb04(s), KNbO3(s), K3Nb04(s), and Ca(NbO3)2(s) taken from previous compilations. Solid AfG° Reference Comments [kJ/mol] NbCl5(c) -700.7 [1963WIC/BLO] T= 298.15 K, I=n/a -682.7 [1971HEL/WOR] T= 298.15 K, I=n/a -691.8 [1971NAU/RYZ] T= 298.15 K, I=n/a -687.3 [1979KUB/ALC] T= 298.15 K, I=n/a -683.2 [1982WAG/EVA] T= 298.15 K, I=n/a -684.1 [1985CHA/DAV] T= 298.15 K, I=n/a -682.7 [1985UDU/VEN] T= 298.15 K, I=n/a NbOCl3(c) -790.9 [1971HIL/WOR] T= 298.15 K, I=n/a -784.8 [1979KUB/ALC] T= 298.15 K, I=n/a -782 [1982WAG/EVA] T= 298.15 K, I=n/a NbF5(c) -1699 [1971HILAVOR] T= 298.15 K, I=n/a -1700 [1977BAR/KNA] T= 298.15 K, I=n/a -1700 [1979KUB/ALC] T= 298.15 K, I=n/a -1699 [1982WAG/EVA] T= 298.15 K, I=n/a -1699 [1982WAG/EVA] T= 298.15 K, I=n/a NbF6(c) -1699 [1971NAU/RYZ] T= 298.15 K, I=n/a NaNbO3(c) -1242 [1971NAU/RYZ] T= 298.15 K, I=n/a -1234 [1981LIN/BES] T= 298.15 K, I=n/a -1233 [1982WAG/EVA] T= 298.15 K, I=n/a Na3Nb04(c) -1811 [1981LIN/BES] T= 298.15 K, I=n/a KNbO3(c) -1261 [1971NAU/RYZ] T= 298.15 K, I=n/a -1249 [1981LIN/BES] T= 298.15 K, I=n/a -1251 [1982WAG/EVA] T= 298.15 K, I=n/a K3NbO4(c) -1825 [1981LIN/BES] T= 298.15 K, I=n/a Ca(NbO,)?(c) -606.2 [1971NAU/RYZ] T= 298.15 K, I=n/a 288 JNC TN8400 99-011 8.5 Comments on selected references [1968SPI]: [1968SPI] determined the protonation constants of Nb6Oi98- by potentiometric measurements in 0.1 - 3 M KC1. [1968SPI] extrapolated his results graphically to 1=0. 8 [1994ETX/FER]: [1994ETX/FER] determined the protonation constants of Nb6Oi9 - by potentiometric measurements in 3 M KC1. [1994ETX/FER] extrapolated their results and the results of [1964NEU and 1968SPI] to 1=0 with an extended Davies equation. However, for the determination of the first protonation constant of NbgOig8- 8 + 7 (Nb6Oi9 - + H o Nb6O19H -) the data of [1968SPI] were not included, making the extrapolation made by [1994ETX/FER] from 1=3 to 1=0 somewhat doubtful. [1992YAJ/TOB]: see [1994YAJ]. [1994YAJ]: Recently, [1992YAJ/TOB] and [1994YAJ] determined in 0.1 M NaCl a minimum Nb(V) solubility < 1 x 10~8 M or a log K*so > 8 for the theoretical equilibrium Nb(OH)s° <=> 0.5Nb2O5(precip) + 2.5H2O. From these data it can be concluded that in the neutral pH range, niobium solubility is smaller than 1 x 10~8 M. It is not possible, due to the lack of experimental data, to give a reliable thermodynamic constant for the solubility in neutral or acidic solutions. 289 JNC TN8400 99-011 9 Palladium1 Palladium is a typical B-metal (,,soft" character). In contrast to the actinides, for example, Pd2+ forms weak complexes with fluoride but strong complexes with chloride, bromide and iodide. Pd2+ has also a strong affinity to nitrogen donors such as ammonia, ethylenediamine, pyridine and its derivatives, etc. 9.1 The redox pair Pd2+/Pd(s) Palladium is a noble metal, which tends to precipitate as Pd(s) from uncomplexed or weakly complexes Pd(II) solutions. The redox potential of the half cell Pd2+ + 2e" = Pd(s) (1) has been measured by Templeton, Watt and Garner [1943TEM/WAT] in 4 M HC1O4 solution and by Izatt, Eatough and Christensen [1967IZA/EAT]. No other experimental examinations of the redox reaction (1) have been found. The ionic strength of the latter investigation was probably 0.1 M, so the pH cannot be lower than 1, and Pd2+ must be suspected to have been present in partly hydrolyzed form, see comments in Section 9.7. Since the hydrolysis constants of these authors are doubtful, their reported redox potential of £°(1, 298.15 K) = (0.915 ± 0.005) V cannot be recommended. The measurement of Templeton, Watt and Garner [1943TEM/WAT] in 4 M HC1O4 solution is thus the only credible redox experiment available for Reaction (1). Although the larger part of their investigation was performed in hydrochloric acid medium, they did two measurement in HC1O4. The resulting mean value is £(1, 4 MHC1O4, 298.15 K) = 0.987 V. It is difficult to extrapolate this value to / = 0, because we have no reliably experimental information at lower ionic strengths. If we wanted to make an attempt to extrapolate this value to / = 0, we can convert it to log K for convenience, log K{\, 4 M HC1O4, 298.15 K) = 33.37, and by using the SIT equation, log K{2) + 4D- log K°(2) - Ae(2)7 for the reaction Pd2+ + 2 e = Pd(s) (2) with D = 0.261 (/ = 4.89 m) and As = -0.3 (from e(M2\ C1O4") « 0.3, cf. [1992GRE/FUG, Appendix B]), we obtain log K°(2, 298.15 K) = 32.86 and £°(2,298.15 K) = 0.972 V. This chapter was partly written by Hans Wanner, HSK, Villigen, Switzerland 290 JNC TN8400 99-011 The uncertainty in the resulting potential is quite large due to the uncertain ion interaction parameter. If we assume a maximum uncertainty of ±0.1 in Ae, cf. [1992GRE/FUG, B.I.3], and an initial uncertainty of ±0.007 V in the reported redox potential {i.e., ±0.24 in log K), an uncertainty of ±0.54 is calculated in log K°{2), which corresponds to ±0.016 V in E°{2). Table 9.1: Experimental equilibrium data compiled for the redox pair Pd2+/Pd(s), according to the equilibria Pd2+ + 2e~ <=> Pd(s). These data were chosen for the evaluation of recommended values in the present report. Additional information for the reference see Section 9.7. Method: pot = potentiometry. log K Reference Comments I (M) Medium Method log K: Pd2+ + 2er <=> Pd(s) 33.37 [1943TEM/WAT] T= 298.15 K, 1=4 HC1O4 pot Table 9.2: Additional, experimentally determined equilibrium data compiled for the redox pair Pd27Pd(s), according to the equilibria Pd2+ + 2e~ <=> Pd(s). These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text. Method: pot = potentiometry. log K Reference Comments I (M) Medium Method log K: Pd2+ + 2e <=> Pd(s) 30.93 n967IZA/EAT1 T= 298.15 K, 1=0. Pd(ClO4)? pot Table 9.3: Thermodynamic data for the redox pair Pd27Pd(s) taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. logK Reference Comments log K: Pd2* -h 2e <=> Pd(s) 33.35 [1952LAT] T= 298.15 K, 1=0 31.07 [1968GOL/HEP] T= 298.15 K, 1=0 30.93 [1980BE1OTEA] T= 298.15 K, 1=0 30.92 [1982WAG/EVA] T= 298.15 K, 1=0 30.93 [1985BAB/MAT] T= 298.15 K, 1=0 30.92 [1985COL] T= 298.15 K, 1=0 30.92 fl988PHI/HALl T=298.15 K, 1=0 291 JNC TN8400 99-011 9.2 Hydrolysis ofpalladium(II) 9.2.1 Hydrolysis ofpalladium(II) The Pd2+ aqua ion is a fairly strong acid: Non-complexing solutions of Pd2+ are stable at pH = 0 but start hydrolyzing at pH = 0.7 [1984WAN]. However, the reproducibility of the visible absorption spectra around pH = 0.7 and higher is poor, indicating that the reaction may not be a + simple, mononuclear hydrolysis reaction leading to PdOH , Pd(OH)2°, etc. In contrast, the hydrolysis of Pd2+ most probably involves polynuclearization and the subsequent formation of colloidal species. Papers reporting hydrolysis constants of Pd2+ are thus to be regarded with great care (see discussions of the papers mentioned hereafter in the Section 9.7 ,,Comments on selected references"). The published hydrolysis data of Izatt, Eatough and Christensen [1967IZA/EAT] and Nabivanets and Kalabina [1970NAB/KAL] are unreliable. Wood [1991 WOO] observed an increase of the solubility of Pd metal at pH values above 11, but the solution composition varied strongly. On this basis it is not possible to select any hydrolysis constants for palladium(II). The suggestion of Byrne and Kump [1993BYR/KUM] to relate the stepwise formation constants of the Pd(II) hydroxide complexes to those of the Pd(II) chloride complexes, i.e., Kj+l/K- is similar for OFT and Cl" as a ligand, is unlikely to be correct due to the strong tendency of the Pd(II) hydroxide complexes toward polymerization, contrary to the Pd(II) chloride complexes. Predicted hydrolysis constants are unlikely to be any better, except perhaps at extremely low Pd(II) concentration where polynuclearization can be expected to be negligible. However, it is at present unknown at what Pd(II) concentration this will be the case. Several hydrolysis studies of Pd(II) have been carried out in seawater or otherwise using chloride as a background electrolyte [1967KAZ/PTI, 1984MIL/BUG, 1989KUM/BYR, 1991TAI/JAN]. Of course, due to the high stability of the chloro complexes of Pd(II), 2+ 2 hydrolysis of Pd (more correctly: PdCl4 ") starts at considerably higher pH values under these conditions than in the absence of any complexants. In seawater (0.558 M Cl") hydrolysis of Pd(II) starts at pH values between 7 and 8 [1989KUM/BYR, Figure 2]. From the spectroscopic work of Kump and Byrne [1989KUM/BYR] we can derive a constant for the formation of 2 2 2 PdCl3OH ~, and a limiting constant for the formation of PdCl2(OH)2 ", both from PdCl4 ~, see comments on Ref. [1989KUM/BYR] below. These constants are valid at / = 0.7 M, but the correction factor to infinite dilution (/ = 0) should be small because of the well-balanced (isocoulombic) charge pattern of the reaction. 2 = PdCl3OH -+ Cl" \ogK°= 4.8 2 = PdCl2(OH)2 -+ 2 Cr \ogK°< 9.3 Using the hydrolysis notation, we obtain (with log A^w° = -14.0): 2 2 + PdCl4 " + H2O(1) = PdCl3OH - + H + Cl" log *K° = -9.2 2 2 + PdCl4 " + 2 H2O(1) = PdCl2(OH)2 -+ 2 H + 2 Cr log *K° < -18.7 292 JNC TN8400 99 - Oil The hydrolysis data base of Pd(II) selected in this review is quite incomplete. It should be noted that these data allow to predict Pd(II) hydrolysis with confidence only at chloride concentrations of 0.5 M and higher, and at pH values lower than about 9 to 9.5. At lower chloride concentrations other, less chlorinated hydrolysis species might become dominant. 9.2.2 Additional equilibrium data compiled for the hydrolysis ofpalladium(II) Table 9.4: Experimentally determined equilibrium data compiled for the palladium(II) hydroxide system, 2+ 2 m + according to the equilibria Pd + mH20 <=> Pd(OH)m ~ + mH . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text and in Section 9.7. Method: sol = solubility measurements, sp = spectrophotometry, pot = potentiometry, and tit = titration (pH) log P,.m Reference Comments KM) Medium Method + + + log fii.f P^ + H2O <=> PdOH + H -1.38 ' [1967IZA/EAT] T= 298.15 K, 1=0.1 0.1 Pd(C104)2 tit 2 -1.00 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(ClO4)2 tit 2 -1.60 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(ClO4)2 sp -2.28 3 [1970NAB/KAL] T= 290 K, 1=0.1 0.1 HC1O, sol -9.23 4 [1984MII7BUG] T= 298.15 K, 1=0.5 0.5 NaCl pot -9.3 4 [1984MEL/BUG] T= 298.15 K, 1=1 1 NaCl pot -9.35 4 [1984MIL/BUG] T= 298.15 K, 1=1.5 1.5 NaCl pot -9.39 4 [1984MIL/BUG] T= 298.15 K, 1=2 2 NaCl pot -9.45 4 [1984MIL/BUG] T= 298.15 K, 1=2.5 2.5 NaCl pot -9.61 4 [1984MIL/BUG1 T= 298.15 K, 1=3 3 NaCl pot + + log j3!2: P& + 2H2O <=> Pd(OH)2°+ 2H -2.36 ' [1967IZA/EAT] T= 298.15 K, 1=0.1 0.1 Pd(ClO4)2 tit 2 -2.20 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(ClO4)2 tit 2 -1.50 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(C104)2 sp 3 -4.42 [1970NAB/KAL] T= 290 K, 1=0.1 0.1 HC1O4 sol -7.14 5 [1991 WOO] T= 298.15 K, 1=0-0.35 n/a pot + + PcP + 3H2O <=> Pd(OH)f+ 3H 3 -16.57 [1970NAB/KAL] T= 290 K, 1=0.1 0.1 HC1O4 sol -19.14 5 [1991WOO1 T= 298.15 K, 1=0-0.35 n/a pot pd2+ 4H 2 + log Pi.4- + 2® <=> Pd(OH)4 ~+ 4H 3 -29.57 n970NAB/KALl T= 290 K, 1=0.1 0.1 HC1O4 sol 293 JNC TN8400 99-011 Table 9.4: continued 2+ + log p4t4: 4PJ + 4H2O <=> Pd4(OH)4"-+ 4H -28.81 4 [1984MIL/BUG] T= 298.15 K, 1=0.5 0.5 NaCl pot -29.1 4 [1984MIL/BUG] T= 298.15 K, 1=1 1 NaCl pot -29.63 4 [1984MIL/BUG] T= 298.15 K, 1=1.5 1.5 NaCl pot -29.88 4 [1984MIL/BUG] T= 298.15 K, 1=2 2 NaCl pot -30.35 4 [1984MIL/BUG] T= 298.15 K, 1=2.5 2.5 NaCl pot -30.51 4 [1984MIL/BUG1 T= 298.15 K, 1=3 3 NaCl pot 1 log Kw used -13.78 2 extrapolated to I = 0 by [1969IZA/EAT] with extended Debye-Huckel equation 3 log Kw used by [1970NAB/KAL]: -14.0 4 Pd concentration = 2.5 - 40 raM. Formation of chloro complexes. 5 Pd concentration = 0.0001-0.1 raM.I not constant. Calculated in this report with a log K Pd(cr)/Pd2 of 32.86 (cf Section 9.1). Table 9.5: Thermodynamic data for the palladium(II) hydroxide system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log P,,m Reference Comments I (M) Medium 2 + + log p,j: Pd + H2O <=> PdOH + H* -1.30 [1968GOL/HEP] T= 298.15 K, 1=0 0 -2.00 ' [1976BAE/MES] T= 298.15 K, 1=1 1 HC1O4 -1.38 2 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1 -1.00 [1976SMI/MAR] T= 298.15 K, 1=0 0 -2.10 [1981BAE/MES] T= 298.15 K, 1=0 0 -1.3 [1985BAB/MAT] T= 298.15 K, 1=0 0 -1.2 [1987BRO/WAN] T= 298.15 K, 1=0 0 -1.0 3 [1988MOU/WOO] T= 298.15 K, 1=0 0 -2.1 [1988PHI/HAL] T= 298.15 K, 1=0 0 -2.3 [1989WOO/MOU] T= 298.15 K, 1=0 0 -2.3 [1992WOO/MOU] T= 298.15 K, 1=0 0 -5.7 4 [1993BYR/KUM] T= 298.15 K, 1=0.5-1 0.7 seawater -1.87 5 |"1992PEA/BER1 T= 298.15 K, 1=0 0 2 + log pu: Pd + 2H2O <=> Pd(0H)2°+ 2H+ -1.90 [1968GOL/HEP] T= 298.15 K, 1=0 0 -4.00 ' [1976BAE/MES] T= 298.15 K, 1=1 1 HC1O4 -2.36 2 [1976SMI/MAR] T= 298.15 K, 1=0.1 0.1 -2.20 [1976SM1/MAR] T= 298.15 K, 1=0 0 -1.9 [1985BAB/MAT] T= 298.15 K, 1=0 0 -3.42 [1987BROAVAN] T= 298.15 K, 1=0 0 -2.2 3 [1988MOUAVOO] T= 298.15 K, 1=0 0 -4.6 [1988PHI/HAL] T= 298.15 K, 1=0 0 -4.4 [1989WOO/MOU] T= 298.15 K, 1=0 0 -4.4 [1992WOO/MOU] T= 298.15 K, 1=0 0 -12.2 4 [1993BYR/KUM] T= 298.15 K, 1=0.5-1 0.7 seawater -3.80 5 [1992PEA/BER1 T= 298.15 K, 1=0 0 294 JNC TN8400 99-011 Table 9.5: continued 2+ + log Pu: Pd + 3H2O <=> Pd(OH)f+ 3H -12' [1976BAE/MES] T= 298.15 K, 1=1 1 HC1O, -3.6 3 [1988MOUAVOO] T= 298.15 K, 1=0 0 -16.6 [1989WOO/MOU] T= 298.15 K, 1=0 0 -16.6 [1992WOO/MOU] T= 298.15 K, 1=0 0 -19.9 4 [1993BYR/KUM] T= 298.15 K, 1=0.5-1 0.7 seawater -15.94 5 [1992PEA/BER1 T= 298.15 K, 1=0 0 2+ log P,,,: Pd + 4H2O <=> Pd(OH)/-+ 4H+ -13 ' [1976BAE/MES] T= 298.15 K, 1=1 1 HC1O4 -5.2 3 [1988MOU/WOO] T= 298.15 K, 1=0 0 -29.6 [1989WOO/MOU] T= 298.15 K, 1=0 0 -29.6 [1992WOO/MOU] T= 298.15 K, 1=0 0 -28.5 4 [1993BYR/KUM] T= 298.15 K, 1=0.5-1 0.7 seawater -29.36 5 fl992PEA/BER] T=298.15K, 1=0 0 1 estimated by [1976BAE/MES] based on [1970NAB/KAL] 2 log Kw used -13.78 3 P,",and P12 values selected from [1967IZA/EAT], p13and pu estimated by [1988MOUAVOO] 4 [1993BYR/KUM] recalculated log B values based on [1970NAB/KAL] and [1991 WOO]; (log Kw used -13.76) 5 values from [1970NAB/KAL] corrected to 1=0 with Davies equation by [1992PEA/BER] 295 JNC TN8400 99-011 9.3 Solid palladium(H)-oxide/hydroxide In presence of oxygen Pd(cr) is oxidized to PdO(cr) at 600 CC. In solutions which contain Pd(II) an amorphous, yellow-brown palladium-oxide-hydrate precipitates, PdO-H2O(precip) (or Pd(OH)2(precip)) [1955GLE/PEU, 1995WIB]. In absence of water, Pd(OH)2(precip) dehydrates to PdO(cr) above 90 °C, while in presence of water above 100 °C, a PdO(s) solid with disturbed lattice is produced, indicating the incorporation of H2O in the lattice [1955GLE/PEU]. 9.3.1 Pd(OH)2(precip) Only two experimental determinations of the solubility of Pd(OH)2(precip) in aqueous solutions have been found. [1970NAB/KAL, 1971NAB/KAL] measured a constant Pd(II) concentration of 4xlO"6 M between pH 3 and 11 in 0.1 M perchlorate media, corresponding to a log K*s2 of 5.4 for the reaction Pd(OH)2° <=> Pd(OH)2(precip). They did not indicate any detection limit and we suspect that the measured minimum Pd(II) concentration reflects the detection limit of the analytical method used. [1991WOO] determined a constant Pd(II) concentration of approx. 9xl0~8 M between pH 8 and 11 in diluted solutions. Unfortunately, [1991 WOO] was not able to show whether Pd(s), which he used as a starting material, or Pd(OH)2(precip), which he expected based on the Eh measurements, was the solubility limiting phase present. However, it is interesting to note that above a pH of 12 the curves determined by both [1970NAB/KAL, 1971NAB/KAL] and [1991 WOO] agree well. No solubility product for Pd(OH)2(precip) is proposed in this report. 9.3.2 PdO(cr) The standard Gibbs molar energy of formation of PdO(s) has been determined at 700 - 1000 K by [1982LEV/NAR] and [1983MAL/SRE]. Both authors obtained after extrapolating their measurements to 298 K a AfG° of approx. -83 kJ/mol from which a log K°*so of ~ 6 can be 2+ calculated for the reaction Pd + H2O & PdO(cr) + 2H+ (Table 9.6). [1983MAL/SRE] defined their product with X-ray analysis as PdO(cr). This agrees well with the observations of [1955GLE/PEU] who observed that, in absence of water, PdO(cr) crystallizes above 90 °C. The precipiation product from aqueous solutions, however, is PdO-H2O(precip) or Pd(OH)2(precip) and not PdO(cr) [1955GLE/PEU]. 296 JNC TN8400 99- Oil Table 9.6: Experimentally determined equilibrium data compiled for the formation of the solid Pd(II)- oxide/hydroxide. These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text. Method: emf = emf measurements, sol = solubility. logK Reference Comments KM) Medium Method log K*S2: Pd(OH)2° <=>Pd(OH)2(precip) 2 5.40 [ 1970NAB/KAL] T= 290 K, 1=0.1 0.1 HCIO4 sol 7.00 3 [1991WOO1 T= 298.15 K, 1=0-0.002 0 NaOH sol 2+ + log K'SQ: Pd + H2O <=>PdO(crt) + 2H 5.98 4 [1982LEV/NAR] T= 298.15 K, I=n/a emf 5.86 4 [1983MAL/SRE] T= 298.15 K, I=n/a emf 1 Extrapolated to 1=0 with Debye-Hiickel 2 Pd=l.lmM 3 Pd = 0.0001-0.1 mM; solubility limiting phase may Pd(s) and not Pd(OH)2(precip) 4 extrapolated by the respective authors from measurements at 700-1000 K; Calculated in this report with a log K of 32.86 for the reaction Pd2+ + 2er = Pd(s) (Section 9.1) Table 9.7: Thermodynamic data for the for the formation of the solid Pd(II)-oxide/hydroxide taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log K Reference Comments I (M) Medium 2+ + log K*so: Pd + 2H2O <=> Pd(OH)2(s) + 2H 2.98 [1952LAT] T=298.15 K, I=n/a 2.65 [1967IZA/EAT] T= 298.15 K, 1=0 0 Pd(C104)2 0.80 [1968GOL/HEP] T= 298.15 K, 1=0 0 0.60 [1976BAE/MES] T= 298.15 K, 1=0.1 0.1 HC1O4 0.50 [1976SMI/MAR] T= 298.15 K, 1=0 0 0.80 [1981BAE/MES] T= 298.15 K, 1=0 0 0.79 [1988PHI/HAL] T= 298.15 K, 1=0 0 1.6 [1992PEA/BER] T= 298.15 K, I=n/a 2+ + log K*so: Pd + H2O <=>PdO(s) + 2H 2.34 [1952LAT] T=298.15 K, I=n/a 3.63 ' [1954COU] T= 298.15 K, I=n/a 14.14 ' [1967WAR] T= 298.15 K, I=n/a 4.41 [1968GOL/HEP] T= 298.15 K, 1=0 6.68 ' [1971NAU/RYZ] T= 298.15 K, I=n/a 2.82 ' [1973BAR/KNA] T= 298.15 K, I=n/a 4.40 [1976BAE/MES] T= 298.15 K, 1=0 5.74 ' [1977BAR/KNA] T= 298.15 K, I=n/a 5.74 ' [1979KUB/ALC] T= 298.15 K, I=n/a 4.70 2 [1987BRO/WAN] T= 298.15 K, 1=0 0 4.24 [1988PHI/HAL] T= 298.15 K, 1=0 4.40 [1992PEA/BER1 T= 298.15 K, I=n/a 1 calculated in this report with a log K of 32.86 for the reaction Pd2+ + 2e" = Pd(s) (Section 9.1) 2 reported in [1987BROAVAN] for 'Pd(OH)2(cr)' 297 JNC TN8400 99-011 9.4 Chloride complexes of palladium(H) 9.4.1 Chloride complexes Among the other ligands considered in the PNC-TDB project, reliable formation constants are + only available for the chloride complexes of palladium(H): PdCl , PdCl2(aq), PdCl3" and 2 PdCl4 ~. We concur with the review of Wood, Mountain and Pan [1992WOO/MOU] that the most reliable study of Pd(II) chloride complexes is the spectrophotometric investigation by Elding [1972ELD] at/= 1 M (HC1OJ. The results of other experimental studies do not differ greatly from Elding's values. The extrapolation from / = 1 to / = 0 is not straightforward. We use the specific ion interaction equation (cf. NEA-TDB [1995SIL/BID]) and estimate Ae as 2+ 2+ + follows: We assume e(Pd , C1O4") - e(Co , C1O4") = 0.34 and e(PdCl , C1O4") - e(CdCl\ C1O4") = 0.25. For the anionic complexes, we only have an indication of the interaction coefficients with Na+, not with H+, cf. Table B.4 [1995SIL/BID]. We use average values of the + + 2 + 2 respective charge patterns, e(PdCl3~, Na ) » e(M", Na ) = 0.00 and e(PdCl4 ", Na ) « e(M ", Na+) = -0.10. In order to compensate the use of the Na+ interaction coefficient for the anionic complexes, we use e(Cl~, Na+) instead of £(C1~, H+) for the last two reactions. In this way we obtain the following constants at / = 0 (with errors between ±0.1 and ±0.2): Pd2+-i-cr <=>Pdcr log p,° = 5.1, Ae = -0.21 2+ Pd -h2cr <=> PdCl2° log p2° = 8.3, Ae = -0.58 2+ Pd -1-3 cr <^ PdCl3- log (33° = 10.9, Ae = -0.43 2+ 2 Pd -H4cr <=> PdCl " log P4° = 11.7, Ae = -0.56 The bromide complexes measured by Elding [1972ELD] are equally reliable, but bromide is not a priority ligand in the present project. The formation constants of Pd(II) with many other ligands, such as carboxylic acids and numerous nitrogen donors, have been determined, but complexes with carbonate, phosphate or sulfate are unknown to our knowledge. It should be mentioned that carbonate complexes of Pd(II) may simply not form at all because of the enormous competition by hydroxide. 298 JNC TN8400 99 -Oil Table 9.8: Experimental equilibrium data compiled for the palladium(II) chloride system, 2+ 2 m according to the equilibria Pd + m Cl~ <=> PdClm " . These data were chosen for the evaluation of recommended values in the present report. Additional information for the reference see text. Method: sp = spectrophotometry. Reference Comments I (M) Medium Method log [3l.m 2+ + log Pu: Pd + Cl- <=> PdCl 4.47 [1972ELD] T= 298.15 K, 1=1 Na,HC104 2+ log p1>2: Pd + 2Ct «• PdCl2° 1.16 [1972ELD] T= 298.15 K, 1=1 Na,HClOd 2+ log pli3: Pd + 3CI- & PdCl3- 10.17 [1972ELD] T= 298.15 K, 1=1 Na,HC104 sp 2+ 2 log PiA: Pd + 4Ct <=> PdCl4 - 11.54 [1972ELD] T= 298.15 K, 1=1 Na,HClO4 9.4.2 Additional equilibrium data compiled for palladium chloride complexes Table 9.9: Additional, experimentally determined equilibrium data compiled for the palladium(II) chloride 2+ 2 m system, according to the equilibria Pd + m Cl" <=> PdClm ~ . These data were not chosen in the present report for the evaluation of recommended stability values. Reasons for not selecting these references are given in the text. Method: sp = spectrophotometry, pot = potentiometry. log !*,.„ Reference Comments I (M) Medium Method 2 log pu.- Pd * + cr <=> Pdcr 4.70 ' [1969GEL/KIS] T= 298 K, I=n/a n/a pot 2 4.47 R991TAI/JAN1 T= 292 K, 1=0.01-0.5 NaCl sp 2+ log Pi.2: Pd + 2CI- <=> PdCl2° 7.70 ' [1969GEL/KIS] T= 298 K, I=n/a n/a pot 2 8.01 [1991TAyjAN] T= 292 K, 1=0.01-0.5 NaCl sp 299 JNC TN8400 99-011 Table 9.9: continued 2+ log Pu: Pd + 3Cl~ <=> PdCl 10.30 • [1969GEL/KIS] T= 298 K, I=n/a n/a pot 10.61 2 T1991TAI/JAN1 T= 292 K, 1=0.01-0.5 NaCl 2+ log PL4: Pd + 4CI- <=> 11.92 ' [1969GEL/KIS] T= 298 K, I=n/a n/a pot 2 11.84 [1991TAI/JAN] T= 292 K, 1=0.01-0.5 NaCl sp 2 log K1A: PdClf + Cl- <=> PdCl4 - 1.433 [1968LEV] T=298K, 1=1 1 LiC104 sp 1.593 [1968LEV] T= 298 K, 1=2 2 LiC104 sp 3 1.77 [1968LEV] T= 298 K, 1=3 3 LiC104 sp 3 2.01 [1968LEV] T= 298 K, 1=4 4 LiC104 sp 1.27 4 [1973GUL/SCH] T= 298 K, 1=0.04-0.3 0 NaCl sp 'I not given, probably not constant 2 I not constant 3 Pd = 0.1 mM; only log K4 measured at different I; 4 Pd = 1.2 mM I not constant, only log K4 measured at different I Table 9.10: Thermodynamic data for the palladium(II) chloride system taken from previous compilations. As pointed out in Section 2 of this report only experimental data were used for the present evaluation. The following table serves only for comparison. Medium: Where data refer to specific electrolyte solutions, this is indicated. log pliir Reference Comments Medium + + log pu: PoF + Cl- <=> PdCl 6.08 [1967AHR] T= 298.15 K, 1=0 0 4.47 [1976SMI/MAR] T= 298.15 K, 1=1 1 6.10 [1976SMI/MAR] T= 298.15 K, 1=0 0 5.08 ' [1980KRA] T= 298.15 K, 1=0 0 CIO4- 3.94 [1982WAG/EVA] T= 298.15 K, 1=0 0 6.00 [1985BAB/MAT] T= 298.15 K, 1=0 0 7.37 2 [1985COL] T= 298.15 K, 1=0 0 5.00 [1987BRO/WAN] T= 298.15 K, 1=0 0 4.98 [1988PHI/HAL] T= 298.15 K, 1=0 0 3.97 [1992PEA/BER] T= 298.15 K,I=n/a 300 JNC TN8400 99-011 Table 9.10: continued 2+ log (Sj,2: Pd + 2CI- <=> PdCl2° 7.74 [1976SMI/MAR] T= 298.15 K, 1=1 1 10.70 [1976SMI/MAR] T= 298.15 K, 1=0 0 8.16 ' [1980KRA] T= 298.15 K, 1=0 0 C1O4- 7.51 [1982WAG/EVA] T= 298.15 K, 1=0 0 10.60 [1985BAB/MAT] T= 298.15 K, 1=0 0 9.44 2 [1985COL] T= 298.15 K,I=0 0 8.77 [1987BRO/WAN] T= 298.15 K, 1=0 0 7.71 [1988PHI/HAL] T= 298.15 K, 1=0 0 7.51 [1992PEA/BER] T= 298.15 K,I=n/a 2+ log pu: Pd + 3CI- <=> PdClf 10.20 [1976SMI/MAR] T= 298.15 K, 1=1 1 13.10 [1976SMI/MAR] T= 298.15 K, 1=0 0 10.58 ' [1980KRA] T= 298.15 K, 1=0 0 cio4- 10.33 [1982WAG/EVA] T= 298.15 K, 1=0 0 12.94 2 [1985COL] T= 298.15 K, 1=0 . 0 11.46 [1987BRO/WAN] T= 298.15 K, 1=0 0 10.99 [1988PHI/HAL] T= 298.15 K, 1=0 0 10.32 [1992PEA/BER] T= 298.15 K,I=n/a 2+ 2 log j3L4: Pd + 4CI- <=> PdCl4 - 12.26 [1952LAT] T= 298.15 K, 1=0 0 14.15 2 [1963GRI/KIS] T= 298.15 K, 1=1 1 10.94 [1968GOL/HEP] T= 298.15 K, 1=0 0 11.50 [1976SMI/MAR] T= 298.15 K, 1=1 1 15.40 [1976SMI/MAR] T= 298.15 K, 1=0 0 11.46 ' [1980KRA] T= 298.15 K, 1=0 0 C1O4- 9.50 [1982SMI/MAR] T= 298.15 K, 1=0.5 0.5 12.04 [1982WAG/EVA] T= 298.15 K, 1=0 0 13.96 2 [1985COL] T= 298.15 K, 1=0 0 13.09 [1987BROAVAN] T= 298.15 K, 1=0 0 11.83 [1988PHI/HAL] T= 298.15 K, 1=0 0 12.04 [1992PEA/BER] T= 298.15 K,I=n/a 1 extrapolated with SIT from data from different sources [1980KRA] 2 Calculated in this report assuming a log K of 34.32 for the reaction Pd2+ + 2e+ = Pd(s) (Section 9.1), corresponding to a AfG° of 187.57 kJ/mol. 301 JNC TN8400 99-011 9.5 Amino complexes ofpalladium(II) Although nitrogen containing ligands are not among the ligands selected by PNC, it may be useful to have recommended constants for the amino complexes, because nitrogen containing ligands form very stable complexes with Pd(II). Rasmussen and J0rgensen [1968RAS/J0R] determined the consecutive formation constants of all four amino complexes of Pd(II) using visible absorption spectroscopy. Like the spectroscopic study of Elding [1972ELD], this study is reliable and can be recommended. Due to the isocoulombic equilibria, the ionic strength dependence will be very small, and we believe that we can neglect it. The constants reported by Rasmussen and J0rgensen [1968RAS/J0R] are thus recommended at any ionic strength (with errors also between ±0.1 and ±0.2): 2+ 2+ Pd + NH3(aq) = PdNH3 log (3,° = 9.6, As = 0 2+ 2+ Pd + 2 NH3(aq) = Pd(NH3)2 log (32° = 18.5, Ae = 0 2+ 2+ Pd + 3 NH3(aq) = Pd(NH3)3 log (33° = 26.0, Ae = 0 2+ 2+ Pd + 4 NH3(aq) = Pd(NH3)4 log P4° = 32.8, Ae = 0 9.6 Conclusions It is unfortunate that the hydrolysis reactions cannot be quantified in a reliable way. The main problem that precludes any reliable quantification of the Pd(II) hydroxide system at this time is the strong tendency toward polymerization. It is likely that the composition of such polymers, changes with pH and Pd(II) concentration. However, the solubility of Pd(II) is very small in the absence of strong complexants such as organic ligands containing amino donors. On the other hand it is known that the affinity of metal ions to oxide-type surfaces is related to their hydrolysis behavior, and it can thus be expected that Pd(II) may form very stable surface complexes on these types of solids. 302 JNC TN8400 99-011 9. 7 Comments on selected references: [1943TEM/WAT]: In this careful work, Templeton, Watt and Garner authors determined the redox potential of the couple Pd27Pd(s) in 0 to 4 M hydrochloric acid and in 4 M perchloric acid at 25°C. Two measurements were performed in HC1O4 (both at / = 4 M). The resulting potentials, corrected by the authors to refer to the molal hydrogen electrode, are E = 0.9849 V and 0.9895 V for Pd2+ concentrations of 9.73 mM and 2.62 mM, respectively, which gives a mean value of 0.9872 with a simple standard deviation of ±0.0033 V. The reported ionic strength was 4.02 M (= 4.89 m) and the hydrogen concentration in the experiments was 4.007 M. [1967IZA/EAT]: Izatt et al. investigated the hydrolysis of Pd2+ at 25°C by pH titration and spectroscopy at pH values between 0.5 and 2.3. The ionic strength was not maintained constant and hence varied between 0.005 and 0.3 M. The concentration of Pd(II) was in the millimolar range. Although 5 hours were required to establish constant pH values, which might indicate polymerization reactions, only monomeric hydrolysis species were considered in the reported data set. From the spectroscopic experiments the authors reported errors as large as 30% and 85% in the extinction + coefficients of ,,PdOH " and ,,Pd(OH)2", respectively. In addition, only the choice of some selected data points resulted in positive values for the hydrolysis constants. The authors used an extended Debye-Hiickel equation to extrapolate the constants to / = 0 and reported the following values from the two methods (pH titrations and spectroscopic measurements): Pd2+ + OH" = PdOH+ log p°, = 13.0 (pH); 12.4 (sp) 2+ Pd + 2 OH" = Pd(OH)2° log p°2 = 25.8 (pH); 26.5 (sp) These constants cannot be accepted in this review due to the reasons mentioned above. In addition, these constants do not follow the usual pattern of formation o constants, in particular those derived from the spectroscopic measurements: log p2 is 1.7 orders of magnitude larger than 2 x log p,°, which casts large doubts on the credibility of the reported values. In addition, the authors determined the redox potential for the reaction Pd2+ + 2 e~ = Pd(s). However, no raw data are given in any form. The only reported parameter is a standard half cell potential of (0.915 ± 0.005) V that had been extrapolated to I = 0 by use of an extended Debye-Htickel equation. It is, however, not clear at what ionic strength the redox measurements had been performed. The authors explained that they used ,,some of the data obtained in the determination of the formation constant of the tetracyanopalladate(II) ion" [1967IZA/WAT]. However, the purpose of that paper was not the determination of a redox potential, and it is not indicated explicitly at what ionic strength the measurements were done. However, the authors corrected for hydrolysis of Pd2+ using the hydrolysis constants from [1967IZA/EAT], which indicates that they did not work at pH = 0. We suppose that they worked at pH = 1 or higher, as the ionic strengths reported in the paper [1967IZA/WAT] was / = 0.1 M 303 JNC TN8400 99-011 or lower. As a result of the insufficient credibility of the hydrolysis constants from [1967IZA/EAT], see assessment above, the reported redox potential cannot be recommended here. [1970NAB/KAL]: Nabivanets and Kalabina used batch solubility measurements of Pd(OH)2(s) in the pH range from 0.7 to 13.5, at 17°C, and in 0.1 M perchlorate solutions to examine the complex formation of Pd2+ with the hydroxide ion. They measured a constant Pd(II) concentration of 4x10~6 M from pH 3 to 11. The reproducibility of the experiments was not verified. The determination of the Pd(II) concentration was done spectrophotometrically with xylenol orange. The minimum concentration of Pd(II) required for this method, as reported by Otomo [1963OTO], was between 0.2 and 0.8 ppm, i.e., between 2X10"6 and 8X10"6 M Pd(II). Hence, the Pd(II) concentration in the solution to be analyzed needs to be even higher due to subsequent dilution caused by the addition of acid and xylenol orange. The measured minimum Pd(II) concentration of 4x10"6 M is thus very likely to reflect the detection limit of the analytical method used. + + The authors reported the following constants, valid at 7 = 0.1 M (H /Na , C1O4~)- For the ion product of water, the authors used K^ = 10"14, which may be somewhat erroneous at the ionic strength of the system. The consecutive constants (log Kn) below refer to the same medium: 2+ + Pd + OH" PdOH log ,(' = 0.1 M)= 11.7 log K\ = 11.7 2+ p Pd + 2 OH" = Pd(OH)2° log p a= 0.1 M)= 23.6 log K2 = 11.9 2+ 2 Pd + 3 OH' = Pd(OH)3- log 0.1 M)= 25.4 log K3 = 1.8 2+ 2 p Pd + 4 OH" = Pd(OH)4 " log pii= 0.1 M)= 26.4 log K4 = 1.0 Pd(OH) 2° Pd(OH)2(s) log K 0.1 M)= 5.4 Looking at these values, it is not plausible why, on one hand, log K2 is larger than log Kx, and on the other hand, log K3 and log K4 so much smaller than log K2, although a steric hindrance is inconceivable. Quite certainly, something is suspect about these data. It is likely that the minimum solubility measured by the authors was equal to the detection limit of the method, and that the actual solubility of the Pd(OH)2(s) used in the experiments was in fact much lower. In addition, extensive formation of colloids would also induce significant changes in the above constants. However, there are other shortcomings in this paper. Comparative dialysis with Co2+ as a reference ion in the pH range between 1.1 and 2.0, where the solubility of Pd(OH)2(s) increases strongly, showed a relative decrease of the Pd(II) concentration after dialysis with decreasing pH. The authors postulated a general polymerization according to p Pd(II) = [Pd(II)], 304 JNC TN8400 99-011 and calculated the ,,polymerization constants" from the dialysis experiments at pH = 1 and Pd(II) concentrations ranging from UxlO"4 M to 3.4x10"* M. Assuming that no polymerization occurs at a Pd(II) concentration of 6.3x10"5 M, the authors calculated the polymerization constants for different polymerization degrees (p = 2 to 6) at each palladium concentration and found more or less constant values only for p = 6 in the concentration range of 1.3x10^ M to 2.4X10"4 M. The fact that only a slight increase of the Pd(II) concentration to 3.4x10"4 M reduces the polymerization constant by a factor of as much as 100, gives rise to serious doubts about the reliability of the reported results. The calculation of log K2 (= log P2 - log P,) from the point at which the solubility of Pd(II) is twice the minimum solubility of 4x10^ M, presumes that + the only monomeric aqueous species are Pd(OH)2(aq) and PdOH . However, according to the resulting constants Pd2+ would also be present at significant concentrations at these pH values. For the derivation of log P2 (at pH = 1.55 to 2), the data were corrected for the formation of the polymer (p = 6), although the polymer data were obtained at a constant pH of 1. However, the pH dependency of the polymer formation, which is expected to be significant in this case, was not examined by the authors. In addition, the suspicion that the minimum solubility of 4x10~6 M Pd(II) merely reflects the detection limit of the analytical method used by the authors casts additional doubt on the credibility of the reported results. The authors also derived log P3 and log P4 from very few points (see [ 1970NAB/KAL, Figure 1]: 4 points in total at pH values of about 12.2, 12.6, 13.0 and 13.6). The basis is too weak for these two constants to be selected. [1984MIL/BUG]: Milic and Bugarcic investigated the hydrolysis of Pd2+ in sodium chloride solutions. They used 0.0025, 0.005, 0.01, 0.02 and 0.04 mM Pd2+ solutions in 3 M NaCl, and 0.01 M Pd2+ solutions in 0.5, 1, 1.5, 2, 2.5 and 3 M NaCl solutions. They found that hydrolysis started near a pH of 8, and they postulated the formation + 4+ of PdOH and Pd4(OH)4 . Hydrolysis increased with increasing Pd(II) concentration, and decreased with increasing NaCl concentration, but the maximum hydrolysis degree in the experiments was only 0.13. The authors qualified the observed strong NaCl concentration dependence of the Pd(II) hydrolysis as ,,medium effect", although the formation of chloride complexes of Pd2+ had been shown spectroscopically by several authors, for example Elding [1972ELD]. The study of Milic and Bugarcic [1984MIL/BUG] is not suitable to derive any hydrolysis constants of Pd2+ and has to be discarded. [1989KUM/BYR]: Kump and Byrne measured the UV absorbance of Pd(II) in seawater at 25°C between pH = 7.1 and pH = 8.7. They used a surface seawater taken from the Gulf of Mexico with a salinity of 36%o. The total chloride concentration was 0.558 M, and the Pd(II) concentrations applied were 10~5 and 5xl0~6 M, respectively. From the change in the UV absorbances at four different wavelengths (290, 300, 310 and 320 nm), the authors derived the following constants: 305 JNC TN8400 99-011 + 0 2 f[PdOH ] + [PdClOH ] + [PdCl2OH"] + [PdCLOH "]) x n _ _V / s-i\ 1 [Pd2+ ] + [PdCl+ ] + [PdCl5 ] + [PdClg ] [PdClJ" ] x H+ ([Pd(OH)5 ] + [PdCl(OH)2 ] + [PdCl2 (OH)!" l) t f D — O) [Pd2+] + [PdCl+] + [PdCl§] + [PdCl3 ] + [PdCl2" ] From their spectroscopic measurements, the authors obtained the values B] = (2.1 ± 9 18 0.6)xl0~ and B2 < 2.4x10~ . At the high chloride concentration of these experiments, we can expect the Pd(II) complexes to be saturated with chloride or hydroxide ligands, i.e., to contain no unhydrolyzed H2O ligands. Hence, the + 0 concentrations of PdOH , PdClOH and PdCl2OH' in Eq. (1), of Pd(OH)2° and 2+ + 0 PdCl(OH)2" in Eq. (2), as well as of Pd , PdCl , PdCl2 and PdCl3" in both equations, are assumed to be negligible. The derived constants Bx and B2 then correspond to the following equations: 2 2 + PdCl4 " + H2O(1) = PdCl3OH -+ H + Cr (3) 2 2 + PdCl4 " + 2 H2O(1) = PdCl2(OH)2 " + 2 H + 2 Cr (4) The equilibrium constants for the above reactions then correspond to K{Z) = 5,x[Cl ] 2 and K(4) = 52x[Cl~] , respectively. The chloride concentration in the seawater used was 0.558 M, and the ionic strength was reported later [1993BYR/KUM] as / = 0.7 M. It is preferable to convert these constants to refer to isocoulombic reactions, i.e., reactions in which reactants and products have the same charge pattern, because their dependence on the ionic strength is minimal: 2 = PdCl3OH - + Cl" (5) 2 4 = PdCl2(OH)2 -+ 2 Cr (6) 2 The constants for these reactions are K(5) = K{3)/Kw and K{6) = K(A)/KW . The constants K(5) and ^(6) are assumed to be independent of the ionic strength. We are further forced to assume that reactions occurring due to the possible presence of other complexants in the natural surface seawater used, have a negligible influence on the values of Bx and B2. The ion product of water at 25°C and / = 0.7 M (NaCl) is about log ^w = -13.72 [1976BAE/MES]. The selected constants for the mixed chloro- hydroxo complexes of Pd(II) are thus: log K°(5, 298.15 K) = log Bl + log [CT] - log Kw = 4.8 log K°(6, 298.15 K) = log B2 + 2 log [Cl"] - 2 log Kw < 9.3 [1991TAI/JAN] Tait et al. used Raman and UV/VIS spectroscopy to determine the speciation of Pd(II) chloro complexes at varying pH and chloride concentration. From the 2 Raman spectra they assigned typical absorption peaks to PdCl4 ~, PdQ3(H2O)~ and 306 JNC TN8400 99 - Oil PdCl2(H2O)2(aq). They concluded that in seawater with a chloride concentration of 0.558 M, they assumed that they had mixed chloro-hydroxo complexes and solid phases of Pd(II). However, they claimed that this would only be possible ,,when OH" serves as a limiting reagent (i.e., only when not enough OH" ligand is present to completely replace the chlorides)". This claim is surprising, because hydroxide ion as a ligand is always available in aqueous solution. In their opinion, in natural systems, mixed chloro-hydroxo complexes would not be stable ,,because OH" readily replaces Cl" over a negligible pH change". Experimental confirmation of this statement is not given. It should be mentioned that Byrne and Kump [1993BUR/KUM] contradicted this statement and claimed that mixed-ligand complex formation would be an important aspect of Pd(II) hydrolysis in natural solutions, cf. comments under [1993BYR7KUM]. [1991WOO] Wood investigated the hydrolysis behavior of Pd(II) and Pt(II) by solubility measurements of the respective metals in 0.0004 M to 10.0 M NaOH solutions. For the solubility measurements Pd metal (,,Pd shot") and Pt wire were both put in contact with the respective NaOH solution and heated to 85°C for 18 days. The temperature was then held at 70°C and 60°C for 14 days each, and then lowered to 25°C. Samples were analyzed after a total of 319 days and 465 days, respectively. The whole procedure was carried out in a glove box under nitrogen atmosphere. Prior to analysis by graphite furnace atomic absorption, the authors separated Pd and Pt by coprecipitation with tellurium. At the end of the experiments, they examined the solid phases microscopically and by X-ray diffraction, and they found no evidence for the presence of Pd(II) hydroxide. However, they detected a white precipitate containing Na, C and O, and they concluded that CO2 had diffused into the reaction solutions, or that some organic matter might have been leached from the polyethylene bottles at higher temperatures. The authors derived the following constants from their measurements: Pd(s) + 2OH" = Pd(OH)2(aq) + 2 e" log p2 = -12.0 Pd(s) + 3 OH" = Pd(OH)3- + 2 e" log p3 = -10.0 (log K, = 2.0) Unfortunately, the ionic strength was not held constant during the experiments. In fact, the ionic strength varied by as much as four orders of magnitude, which per se precludes the recommendation of these data in the selected data set. It should, however, be noted that the authors suspected their own Eh measurements to be in error, because their measured points fall into the stability range of Pd(OH)2(s) (instead of Pd metal) in the Eh-pH diagram. In our view it might be equally justified to question the correctness of the Eh-pH diagram that [1989WOO/MOU] used, which is based on the more than doubtful data of Izatt et al. [1967IZA/EAT] and Nabivanets and Kalabina [1970NAB/KAL]. [1993BYR/KUM] This is a comment paper on the publication of Tait et al. [1991TAI/JAN] including a review of the Pd(II) chloro complexes and hydroxo complexes, as well as extended comments on the formation of mixed ligand complexes. In particular, 307 JNC TN8400 99-011 they contradicted the statement in [1991TAI/JAN] that hydroxide would act as a limiting reagent. They correctly stated that in hydrolysis equilibria ,,hydroxide can never be a limiting reagent". Byrne and Kump [1993BYR/KUM] reported, from their earlier spectroscopic data in seawater (0.558 M Cl") [1989KUM/BYR], the following constant valid for seawater conditions: 2 2 PdCl4 - + OIT = PdCl3OH - + Cr log K(I = 0.7 M) = 4.8 2+ 2 This means that in seawater, hydrolysis of Pd (more correctly: PdCl4 ) starts at a pH between 7 and 8. This finding is consistent with the experimental findings of Kazakova and Ptitsyn [1967KAZ/PTI] and Milic and Bugarcic [1984MIL/BUG]. Byrne and Kump [1993BYR/KUM] presented a model to statistically predict the formation constants of simple mixed ligand complexes. In the case of chloride and hydroxide as ligands, the following equation is used to predict the formation of mixed complexes: (7) Here, a{j accounts for the ,,promotion" of mixed ligand complex formation through statistical effects, and in the present case where (/ + j) = 4, the term a,y = -0.402 [1982BEL/KOL]. The term log 8,y accounts for ,,ligand effects" which are usually considered negligible for the substitution of equally-charged ligands. In this way, by 2 2 using log (34(PdCl4 -) = 11.32 and log p4(Pd(OH)4 ~) = 26.4, they predicted the following constants for the mixed chloro-hydroxo complexes according to Eq. (7): log p31 = 15.69, log P22 = 19.66, and log p13 = 23.23. Finally, Byrne and Kump [1993BYR/KUM] recognized the fact that insufficient experimental information is available to obtain reliable hydrolysis constants of Pd2+, and they estimated them by making the following two assumptions: 1) The value of 2 log P4(Pd(OH)4 ") from [1970NAB/KAL] is correct, and 2) The ratios of the stepwise formation constants of the Pd(II) hydroxide complexes are the same as those of the Pd(II) chloride complexes. In this way they obtained the following + constants: log p,(Pd(OH) ) = 8.1, log P2(Pd(OH)2°) = 15.34, log P3(Pd(OH)3") = 2 21.4 and log P4(Pd(OH)4 ") = 26.5. However, according to this set of constants, Pd2+ would be present as unhydrolyzed aqua ion below a pH of about 5, which is in contradiction to all experimental evidence obtained hitherto in non-complexing aqueous solutions. 308 JNC TN8400 99-011 10 References [1883SCH] Schulze, H. (1883), J. prakt. Chemie, 27, 320 (as cited in [1952LAT]. [1906GOL/ECK] Goldschmidt, H. and Eckhardt, M. 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