The Reactions of Potassium Ethyl Xanthate in Aqueous
THE REACTIONS OF POTASSIUM ETHYL XANTHATE IN
AQUEOUS SOLUTION
BY
NORMAN ROBERT TIPMAN
B.Sc. (Hons.) University of Alberta, 1962
A THESIS SUBMITTED IN PARTIAL';';FULFILMENT OF
THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY
in the Department
of
MINERAL ENGINEERING
We accept this thesis as conforming to the
required standard
THE UNIVERSITY OF BRITISH COLUMBIA
December, 1970 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of British Columbia, I agree that the Library shall make it freely available for reference and study.
I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives. It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.
Depa rtment
The University of British Columbia Vancouver 8, Canada
Date ABSTRACT
The reactions of potassium ethyl xanthate (KEtX) and its oxidation product diethyl dixanthogen (EtX)^ were studied in neutral and mildly alkaline aqueous solution. Dissolved oxygen-was shown to be ineffective in oxidizing KEtX in homogeneous systems (solution), but
(EtX)^ was produced electrochemically at a platinum electrode by application of a small anodic overpotential.
A slow hydrolysis reaction was proposed for the alkaline decomposi• tion of aqueous ethyl xanthate in which carbon disulfide and ethyl alcohol were the decomposition products. The first order rate constant -4 -1 k^ = 7.6 Z 1.0 x 10 hr (22°C) was independent of pH in the region pH 7 to 11. -5 -7 -1 Low concentrations of (EtX)^ (10 to 10 moles liter ) were determined by a new method based on extraction of the aqueous (EtX^ by hexane followed by ultraviolet spectrophotometric determination of the (EtX)^ in the hexane extract. A sensitive technique for measuring saturation of (EtX^ in water using light scattering photometry; was also developed. The solubility of (EtX^ in water was found to be
1.27 x 10~5 moles/liter at 22°C from pH 2 to 8.5.
At pH > 8.5, aqueous (EtX)^ was shown to react with hydroxyl ion by a bimolecular displacement (S^2) mechanism which resulted in the formation of one mole of ethyl xanthate and one mole of an inter• mediate ethyl xanthate sulfenic acid. Decomposition of the intermediate
compound resulted in carbon disulfide and other reaction products. The
second order rate constant (k^ = 13.5 t 0.5 liter mole "Snin ^ at 22°C) was determined. Rest potential measurements using a platinum electrode in potassium ethyl xanthate-diethyl dixanthogen solutions showed that the
Nernst equation for a one electron reaction was obeyed over a wide range of both KEtX and (EtX^ concentrations. Dissolved oxygen was found to generate mixed potentials to which no overall reaction could be assigned.
It is expected that the rapid oxidation of ethyl xanthate to diethyl dixanthogen observed in mineral systems proceeds by a catalytic oxidation or electrochemical reaction at or near the mineral surface. TABLE OF CONTENTS
CHAPTER 1. INTRODUCTION 1
1.1 Chemical Reactions of Xanthates 1
1.2 Reaction of Xanthates with Mineral Surfaces 4
1.3 Xanthate Adsorption Theories 4
1.4 Problems in Explaining the Role of Oxygen in Flotation 6
1.4(a) Adsorbent System: The Reaction of Oxygen with the Mineral Surface 6
1.4(b). Adsorbate System: The Reaction of Oxygen
with the Dissolved Xanthate 7
1.5 The Unresolved Points 8
1.6 Plan of the Present Work 11
CHAPTER 2. EXPERIMENTAL METHODS •.. 12
2.1 Materials 12
2.2 Ethyl Xanthate Decomposition Experiments 13
2.3 Diethyl Dixanthogen Decomposition Experiments .... 13
2.4 Ultraviolet Spectroscopic Analysis for Aqueous Ethyl Xanthate and Carbon Disulfide 14
2.5 Ultraviolet Spectroscopic Analysis for Diethyl Dixianthogen 18
2.6 Light Scattering Method for Determining the
Solubility of Diethyl Dixanthogen 19
2.7 Infrared Spectroscopic Methods 25
..; 2.7Ca) The KBr Pellet Technique 25
2.7(b) Infrared Cell for Aqueous Solutions 25 2.7(c) Attenuated Total Reflectance (ATR) Spectra
of Platinum Surfaces 26
2.8 Electrochemical Methods 28
2.8(a) Cell for Rest Potential Measurements 28
2.8(b) Cell for Electrode Polarization Measurements 28
CHAPTER 3. RESULTS AND DISCUSSION 32
3.1 Decomposition of KEtX in Aqueous Alkaline Solutions 32
3.1.1 Determination of Rate Law - 33
3.1.2 Effect of Oxygen on the Decomposition Rate. 3&
3.1.3 Effect of Hydroxyl Ion on the Decomposition
Rate 41
3.1.4 Evaluation of Reaction Products ^3
3.1.5 Quantitative Measurement of Carbon Disulfide 43
3.1.6 Mechanism of Xanthate Decomposition in Alkaline Solution 46 3.1.7 Decomposition of KEtX in the Presence of Possible Catalytic Agents 50
3.1.7(a) Methylene Blue 51
3.1.7(b) Platinized Platinum 52
3.1.7(c) Ferrous Sulfate 53
3.2 Decomposition of Diethyl Dixanthogen in Aqueous Alka• line Solution 54
3.2.1 Stoichiometry of the Reaction of Diethyl Di• xanthogen with Hydroxyl Ion;. 55
3.2.2 Elimination of Possible Interference by Iodide . 55
3.2.3 Evaluation of Reaction Products ... 57 3.2.4 Contribution of the Reaction Products to the 301 my Xanthate Absorption 58
3.2.5 Mechanism of the Decomposition of Dixanthogen 60
3.2.6 Reaction of Dixanthogen with Other Nucleophiles ...... 61
3.3 Rate Studies on the Reaction Betweeen Dixanthogen and Hydroxyl Ion 59
3.3.1 Reaction Order with Respect to Diethyl Dixanthogen Concentration 65
3.3.2 Reaction Order with Respect to Hydroxyl Ion Concentration 6g
3.4 Solubility of Diethyl Dixanthogen in Water 72
3.4.1 Measurement of Saturation by the Extraction Technique y2
3.4.2 Measurement of Saturation by the Turbidi- metric Technique 73
3.5 Electrochemical Studies on Aqueous Ethyl Xanthate. 75
3.5.1 Derivation of the Nernst Equation for the Oxidation-^of Ethyl Xanthate to Diethyl Dixanthogen 76
3.5.2 Rest Potential Measurements with a Platinum Electrode .... 77
3.5.2Ca) Ethyl Xanthate Solutions Free of Diethyl Dixanthogen 77
3.5.2(b) Ethyl Xanthate Solutions Saturated
with Diethyl Dixanthogen ...... 79
3.5.2(c) Effect of pH on the Rest Potential 86
3.5.2(d) Effect of Diethyl Dixanthogen on the Rest Potential . 88 3.6 Polarization of the.Platinum Electrode in Ethyl Xanthate-Diethyl Dixanthogen Solutions 91
3.7 Infrared Investigations on the Platinum Electrode Surfaces 97 3.8 Interpretation of the Electrochemical Results .... 103
CHAPTER, 4. SUMMARY AND CONCLUSIONS . .... 107
APPENDIX I - Curve Fitting and Error Analysis I12
APPENDIX II - Decomposition of Ethyl Xanthate in Acid Solution 115
REFERENCES 120 LIST OF FIGURES
Figure Page
1 Ultraviolet absorption spectra of aqueous ethyl xanthate, hydroxyl ion and carbon disulfide 16
2 Calibration curve for potassium ethyl xanthate and carbon disulfide in water 17
3 Ultraviolet absorption spectra of diethyl dixanthogen in hexane 20
4 Calibration curve for diethyl dixanthogen in hexane vs. hexane reference .. 21
5 Material balance between the concentrations of ethyl xanthate titrated with iodine and diethyl dixanthogen determined by hexane extraction 22
6 Effect of iodide on the solubility of diethyl dixantho• gen as determined by light scattering, pH 7.5-7.9 ... 24
7 Multiple internal reflection effect illustrating the penetration of the IR beam through the KRS-5 crystal surface . . 27
8 Apparatus for measuring pH and rest potential 30
9 Apparatus for polarization studies 31
10 Decomposition of aqueous ethyl xanthate in (a) solutions saturated with oxygen (b) deoxygenated solutions under argon atmosphere 35
11 Log-log plot of the decomposition rate of aqueous ethyl xanthate vs. initial ethyl xanthate concentration for deoxygenated solutions under an argon atmosphere 38
12 Log-log plot of the decomposition rate of aqueous ethyl xanthate vs. initial ethyl xanthate concentra• tion for solutions saturated with oxygen 40
13 pH. shift of ethyl xanthate decomposition runs in deoxygenated water under argon atmosphere 42
14 Decomposition of ethyl xanthate in deoxygenated water in a sealed cell, pH 7.4 . • . •_ - _j :._ (a) rate of xanthate decomposition, (b) rate of CS2 formation..... ;. . 47 Figure
15 Ultraviolet spectra of sulfur containing compounds in the system CS2~NaOH-alcohol. Cl) dithiocarbonate
in 5 N NaOH; (2) CS2; C3) Na2S in 0.1 N NaOH;
C4) ethyl monothiocarbonate (C2H50C0SK); C5) tri- thiocarbonate CCS3=); (6) potassium ethyl xanthate CC2H5OCSSK); (7) tetrathiopercarbonate (CS4=); (8) trithiopercarbonate (CS30=); (after Hovenkamp (57))
16 First order rate plot for the reaction of aqueous diethyl dixanthogen with hydroxyl ion, Series 4
17 Eyaluation of order for hydroxyl ion in the reaction of diethyl dixanthogen with hydroxyl ion
18 Effect of pH on the solubility of diethyl dixanthogen in water
19 Effect of oxygen on the rest potential of ethyl xanthate solutions free from diethyl dixanthogen, pH 7.0-8.0
20 Effect of oxygen on the rest potential of ethyl xanthate solutions saturated with diethyl dixanthogen, pH 7.0-8.0
21 Eh as a function of dissolved oxygen concentration (moles liter--*-) in distilled water, pH 6.8 (after Natarajan and Iwasaki (66))
22 Rest Potentials of a platinum electrode in solutions saturated with diethyl dixanthogen and under an argon atmosphere
23 Effect of pH on the rest potential of aqueous ethyl xanthate saturated with diethyl dixanthogen
24 Effect of diethyl dixanthogen on the rest potential of a platinum electrode under argon atmosphere and constant ethyl xanthate concentration, pH 6.9-7.7 ...
25 Correlation of the rest potentials of ethyl xanthate and diethyl dixanthogen solutions with the Nernst equation -3 26 Polarization curves for a platinum electrode in 10 M ethyl xanthate in 1 N KC1 electrolyte -3 27 Polarization diagram for a platinum electrode in 10 M ethyl xanthate, 1 M KC1 electrolyte Ca) polarization over a one volt range (b) magnification of (a) to show Tafel behavior. Figure Page
28 Infrared spectra of ethyl xanthate compounds Ca) KEtX, solid, KBr pellet (b) diethyl dixanthogen, capillary film (c) ethyl xanthate ions, aqueous solution (d) platinum xanthate, KBr pellet 99
29 Infrared (ATR) spectra of platinum foil after varying treatments 100
30 Rates and reactions of ethyl xanthate and diethyl dixanthogen in aqueous solution at 22°C 109 LIST OF TABLES
Table Page
1 Decomposition of Ethyl Xanthate Under Argon Atmosphere at 22° C, Series 1 #2 34
2 Decomposition Rates for Aqueous Ethyl Xanthate Under Argon Atmosphere at 22°C 37
3 Decomposition Rates for Aqueous Ethyl Xanthate Under Oxygen Atmosphere at 22° C 39
4 Effect of pH on the Decomposition of Ethyl Xanthate.. 44
5 Relationship between Ethyl Xanthate Decomposition and ptt Increase 45
6 Material Balance Between Ethyl Xanthate Decomposed and CS2 Evolved from Sealed Cell Measurements 48
7 Decomposition of Aqueous Ethyl Xanthate in the Presence of Possible Catalytic Agents 51
8 Reaction of Diethyl Dixanthogen with Hydroxyl Ion at 22°C .. 56
9 Stoichiometry of the Reaction Between Various Nucleophiles and Diethyl Dixanthogen 63
10 Hydrolysis of Dixanthogen with Hydroxyl Ion (Series 4, #3) 67
11 Evaluation of Second Order Rate Constants for Hydrolysis of Diethyl Dixanthogen 71
12 Standard Reduction Potentials of the Xanthate- Dixanthogen Couples 85
13 Vibrational Band Assignments for the Infrared Spectrum of Platinum Ethyl Xanthate 102
14 Correlation Coefficient and Error in Rate Constants for the Oxidation of KEtX Solutions 114
15 Dissociation Constants and Decomposition Rate Constants for Ethyl Xanthic Acid 119 LIST OF SYMBOLS
S KEtX = potassium ethyl xanthate (CILj-CK^O-C-S K )
S S II II
(EtX)2 E diethyl dixanthogen (CH3-CR2-0-C-S-S-C-0-CH2-CH3
S II _ EtX = ethyl xanthate ion (CH -CH -0-C-S )
X = xanthate E alkali alkyl xanthate CR-O-C-S Na (K ))
S S ll II X = dixanthogen (R-0-C-S-S-C-O-R)
R = short hydrocarbon chain
M = molar concentration - moles liter ^ ACKNOWLEDGEMENTS
The helpful guidance of my research supervisor Dr. J. Leja throughout the course of this study is gratefully acknowledged.
I wish to thank the various members of the Department of Mineral
Engineering, Dr. L.G. Harrison of the Department of Chemistry, and -my colleagues for -many helpful discussions.
The following scholarships are gratefully acknowledged,
Sherritt-Gordon Graduate Fellowship (1965)
Giant Yellowknife Mines Graduate Scholarship (1966)
U.B.C. Graduate Fellowship (1967)
U.B.C. Graduate Fellowship (1968). CHAPTER I
INTRODUCTION
Xanthates, or alkyl dithiocarbonates, are a special class of organic thiocarbonate salts that have important industrial uses. Since the turn of the century, xanthates have been used extensively in the manufacture of viscose rayon, and to a small extent, as pesticides. In
1924, CH. Keller discovered that .xanthates could be used to separate selectively sulfide minerals from their ores. Currently, the widest use of xanthates is in the mineral industry, where they constitute the most useful chemical for the selective separation of sulfides from complex ores. Considerable research has shown that the chemical behavior and surface chemical properties of aqueous xanthates are very complex.
1.1 Chemical Reaction of Xanthates S ii - + +
Alkali metal xanthates of formula R-O-C-S K (Na ) where R refers to a short hydrocarbon chain, are easily solubilized in water (i.e.
120 gms/1. at 20°C for potassium ethyl xanthate) and readily ionize to form completely dissociated xanthate ions. Under acidic conditions the xanthate ion is protonated, forming xanthic acid which is unstable and decomposes rapidly into carbon disulfide and alcohol. S II R-O-C-SH y ROH + CS
Xanthic acid has been isolated by extraction into non-aqueous solvents such as benzene or carbon tetrachloride and upon purification was found to have a melting point of ca. -53°C and disproportionated at 25°C (1).
Under neutral and basic conditions, xanthates decompose slowly, but the mechanism and products of these reactions have not yet been clearly established.
Xanthates undergo esterification when reacted with alkyl halides or alkyl sulfates (2). Both reactions proceed rapidly at room tempera• ture in ethanol solvent. R^ and represent different alkyl hydro• carbon chains.
" _ + II
Rj-O-C-S K + R2C1 >• R1-0-C-SR2 + KC1
S S II _ + II
R..-0-C-S K + KR.SO. y R-0-C-SRo + K.SO. 1 2 4 2 2 4
The esters can be hydrolysed by reflux in alcoholic KOH but desulfurization takes place during the reaction (2), resulting in the formation of alkyl monothiocarbonates.
'I II _ +
R-0-G-SR2 + KOH y R-O-C-0 K + R SH
An important chemical property of xanthates is their ability to under• go reaction with a number of oxidizing agents to form dixanthogens. Examples of the reaction of aqueous xanthates with iodine, copper sulfate and potassium tetrathionate have been given by Cambron and
Whitby (3).
S S H II 2R-0-C-S -K + + KI > R-0-C-S-S-C-O-R + 3KI.
S S S S II _ + II » H 4R-0-C-S K + 2CuS0. > R-O-C-S-S-C-O-R + 2R-0-C-S-Cu 4
+ 2KoS0. 2 4
S S II II
2R-0-C-S K + KoS.0, y R-O-C-S-S-C-O-R + 2KoSo0„ 2 4 6 I I 3
Diethyl dixanthogen is a light yellow solid, melting point
31.5-32°C. It is sparingly soluble in water (ca. 1.3 x 10 ^ moles/ liter), but completely soluble in organic solvents and can be purified by vacuum distillation (b.p. 107-9°C at 0.05 mm. Hg) (1).
Diethyl dixanthogens can be desulfurized by treatment with potassium cyanide or sodium arsenate in refluxing ethanol (3), resulting in the formation of xanthogen monosulfides.
S S S II II II
(R-0-C-S)2 + KCN y R-O-C-S-C-O-R + KSCN
S S S II II II
(R-O-C-S) + Na As03 • R-O-C-S-C-O-R + Na AsSO Xanthates also react with metallic ions to form relatively insoluble heavy metal xanthates. The bulk properties of heavy metal xanthates have been given in detail by Sheka and Kriss (4). Poling;-, (5) has shown that the aqueous solubilities are related to the degree of covalency of the metal sulfur bond. For example, covalently bonded -30 -20
gold and copper xanthates are least soluble (Ksp = 10 to 10 ) while ionic metal xanthates have solubilities in the order of 2-5 M.
Extensive studies on the infrared spectra of metal xanthates have been carried out by Leja, et.al. (6) and by Watt and McCormick (7). X-ray crystallographic studies on the structure of Pb, Zn, Sb, As, and
Ni xanthates have been reported (8,9,10,11,12).
1.2 Reaction of Xanthates with Mineral Surfaces
The flotation of minerals has been shown to occur because of a chemical reaction between the xanthate and.the mineral surface. The reaction results in a change of the surface condition from hydrophilic
to hydrophobic once a sufficient coating (usually less than a monolayer) of heavy metal xanthate has been formed. The chemisorption of xanthates onto mineral surfaces has been widely studied and a number of theories have been proposed to explain this phenomenon.
1.3 Xanthate Adsorption Theories
Taggart, et.al. (13) originally proposed that metal xanthates are
formed by a chemical exchange mechanism whereby the less soluble metal xanthate is precipitated on the more soluble mineral surface by well known chemical reactions. However, when solubility data were obtained for lead xanthate and were applied to this theory, it became apparent that PbS could not react with any appreciable amount of xanthate.
The discovery that oxygen played an essential role in flotation complicated many theories. Sutherland and Wark (14) and Gaudin (15) then proposed that xanthate reacted by displacement of ions previously formed on the mineral surface. This included adsorbed ions such as hydroxide or carbonate, and ions formed by oxidation of the sulfide mineral, such as sulfate or thiosulfate. This theory enjoys the largest popularity among workers in the field.
In considering that oxidation of the mineral surface resulted in a large negative electrical double layer potential, Cook, et.al. (16,17) reasoned that adsorption of the negatively charged ions of xanthate through this potential barrier would be impossible. A new theory was proposed whereby an uncharged molecule such as xanthic acid was considered to be the effective adsorbing species. Noting that the concentration of free xanthic acid in neutral to basic solutions would be almost negligible, Leja (18) suggested that dixanthogen, formed by oxidation of xanthate in solution, would be a suitable uncharged molecule that could penetrate the electrical double layer and be effective in at least the first layers of adsorption. This theory, called the "hydrolytic" adsorption mechanism is the most recent and controversial of the three theories. 1.4 Problems in Explaining the Role of Oxygen in Flotation
It is well known that dissolved oxygen is a.necessary reagent for the flotation of most sulfide mineralsStudies on.the effect of dissolved oxygen have considered its interaction with the mineral surface (adsorbent) and its interaction with the xanthate (adsorbate).
1.4.(a) Adsorbent system: The reaction of oxygen with the mineral
surface
Considerable research has been carried out to identify the reaction products formed by oxidation of the mineral surface. In the case of galena (PbS) , Hagihara (19) identified PbSO^ as an oxidation product.
Poling and Leja (21) and Greenler (20), using infrared spectroscopic methods, showed that PbS20^ was formed during the initial stages of oxidation. Recently, Eadington and Prosser (22), using chemical methods, ' found that SO^ , S^O^ an^ S^O^ were present after exposure of the mineral surface to aqueous oxygen. Plaksin (23) and Tolun and Kitchener
(39) showed that the electrochemical potential of galena is signifi• cantly altered by the oxidation products formed on the surface.
Studies on the rate of absorption of oxygen by galena and chalcopyrite have been carried out by Cusack (25).
The result of this research has shown that a number of oxidationJ products are formed on mineral surfaces at different rates. Further studies on the chemisorption of xanthate have been complicated by the fact that not all the oxidation products react with the dissolved xanthate at the same rate. For example, xanthate has been shown to reacted more rapidly with PbS„0„ than PbSO. formed on galena surfaces (21). 1.4.(b) Adsorbate System: The reaction of oxygen with the dissolved
xanthate.
Gaudin, et.al. (26) first recognized that oxidation of xanthate to dixanthogen in aqueous solution may play an important role in the reaction between the mineral surface and xanthate. Dibutyl dixanthogen has been shown to be a more effective collector for the flotation of cement copper than K-n-butyl xanthate (27). Poling and Leja (21) showed that multilayers of lead xanthate resulted from the reaction between oxidized galena surfaces and deoxygenated aqueous solutions containing diethyl dixanthogen. Considering that oxygen alone may be effective in converting xanthate to dixanthogen, PomianowskL and Leja
(18) postulated the reaction in neutral and mildly alkaline soltuions
S S S II _ + II l|
4R-0-C-S K + C>2 + 2H20 +- 2R-0-C-S-S-C-0-R + 40H
xanthate dixanthogen
The reaction model of Pomianowski and Leja proposed that xanthate, dixanthogen, xanthic acid, carbon disulfide and alcohol would reach complex equilibrium depending on the relative concentrations of all of
the species present. Rao and Patel published a series of papers (28,
29,30,31) in which they studied the decomposition of aqueous xanthates
in neutral and alkaline conditions with and without the presence of metallic salts and under oxygen and carbon dioxide atmospheres. In
the presence of oxygen, xanthate solutions (28) were postulated to undergo
two types of reactions: oxidation to dixanthogen and hydrolytic decomp• osition to K„S, K CO., and alcohol. In the presence of ferric (29) or lead (31) salts, the oxidation of aqueous butyl and amyl xanthates were increased but ethyl xanthate was decreased. Copper salts (30) were shown to enhance the oxidation of all xanthate homologues at neutral pH.
The decomposition of xanthate solutions in strong alkaline medium
(0.1 to 10 M) was studied by Philipp and Fichte (32) while investi• gating the ripening of viscose rayon (cellulose xanthate). Subsequent work by Hovenkamp (33), Dautzenburg and Philipp (34) and others (35,36) have postulated that the decomposition of xanthate would proceed by three possible mechanisms.
S II
(1) ROCS + H20 —ROH + CS2 + OH
S II (2) ROCS + OH * ROH + CS20
S S II II _ - OH (3) ROCS + ROH + C03 + 2SH + OH > ROCO + SH
The conditions under which each reaction would be effective was postulated to be dependent upon [OH ], but considerable overlapping of rates was expected, particularily at [OH ] > 1.0 M.
The problem of xanthate oxidation in aqueous solution is clearly shown to be complex and no systematic conclusions appear to be available.
1.5 The Unresolved Points
Pomianowski and Leja (18) developed a reaction model for the decomposition of aqueous ethyl xanthates. They noted that the concentrations of ethyl xanthate, diethyl dixanthogen and carbon disulfide (as determined spectroscopically) reached a stationary state after 15 to 20 days of reaction in a closed vessel at pH 6.5-7.1. These workers concluded that an equilibrium had been reached between the various species and the dissolved oxygen participated as the oxidizing agent in the reactions.
This reaction model was quesioned by Finkelstein (37) who demonstrated that oxygen had no effect on the decomposition rate of aqueous ethyl xanthate. He postulated that the aqueous ethyl xanthate decomposed by a complex hydrolytic mechanism that was based on the work of Philipp and Fichte (32) and has been summarized in the previous review section (1.4.(b)). Since Finkelstein had not undertaken any measurements for the concentration of dissolved diethyl dixanthogen
during the course of his experiment, he had not confirmed whether diethyl
dixanthogen was or was not a reaction product. Consequently, there was no basis for rejecting either the approach of Pomianowski and Leja,
or that of Finkelstein.
Electrochemical investigations on the oxidation of xanthate have
shown that dixanthogen was a product of the reaction, which has been written
S S S n H II 2R0CS ROCSSCOR + 2e
Rest potential measurements have been carried out using a platinum
electrode (38,39,40), silver electrode (41,42), and mercury electrode
(43). The majority of results are in general agreement that the standard electrode potential for the ethyl xanthate-ethyl dixanthogen couple is approximately -0.07 volts. Finkelstein (37), using this value and combining it with the oxygen-hydroxyl ion half cell reaction (viz.
02 + 2H20 + 2e~ -—r 40H~ E° = +0.401 V) calculated that the equilibrium constant obtained from spectroscopic data by Pomianowski and Leja (18) was in error by -0.17 V. Finkelstein (37) attempted to resolve this discrepancy by employing redox indicators in ethyl xanthate solutions and measuring the of the ethyl xanthate-diethyl dixanthogen couple without the introduction of an electrode surface. The method could be questioned since the possibility of a chemical reaction between the ethyl xanthate and the indicator would lead to erroneous results.
A number of questions regarding the chemistry of aqueous xanthate solutions have become apparent from the studies conducted by the various workers.
(1) What is the mechanism of the decomposition reaction of ethyl xanthate under neutral and slightly basic conditions? Does it follow an oxidation mechanism to dixanthogen as suggested by Pomianowski and
Leja (18), or a hydrolytic decomposition mechanism as postulated by
Phillpp and Fichte (32)?
(2) What are the reactions of dixanthogen under basic conditions and are they different from the reaction of xanthate?
(3) Are the electrochemical measurements indicative of a reversible electrode potential between xanthate and dixanthogen as has been suggested or are the measurements subject to interference by the formation of other compounds such as metallic xanthates? The studies on the base decomposition of (EtX)^ have shown that
the concentration of (EtX)2 is rapidly reduced when the aqueous solution contained small amounts of nucleophilic reagents. This observation
appears to be inconsistent with other observations that (EtX)2 is present in mineral systems in detectable quantities. Such evidences have been obtained by Fuerstenau, et al. (71), Gaudin, et al. (72) and Majima and Takeda (40) who have shown that the formation of dixanthogen is necessary for the flotation of pyrite. In order to explain the formation of 1:1 metal:xanthate surface complexes for adsorption of ethyl xanthate on galena, Leja (73) considered that dixanthogen was necessary for attachment to the mineral surface in at least the first adsorbed layer. Studies on adsorption of xanthate and dixanthogen on galena (5), copper (27) and mercury (90) have shown that dixanthogen is an effective reagent in the absence of oxygen. Obviously, the oxidation of xanthate to dixanthogen at the mineral surface is an important process in the flotation of minerals. 1.6 Plan of the Present Work
The plan of work covered in this thesis was a study of the problems
of homogeneous reactions of aqueous ethyl xanthate faced by the previous workers in the field using modern analytical techniques. It was hoped
that the lack of agreement between these workers could be resolved by
evaluating each chemical system for the mechanism of its reaction. The
objectives can be given as follows:
(a) to evaluate the decomposition of aqueous ethyl xanthate in
argon and oxygen atmospheres and determine the mechanism from the rate
and products of the reaction.
(b) to evaluate the reaction between diethyl dixanthogen and
hydroxyl ion and if possible, to establish a mechanism for the reduction
reaction.
(c) to review the results of the electrochemical investigators
and determine if the electrochemical oxidation of ethyl xanthate and
the aqueous oxidation of ethyl xanthate are two compatible systems. CHAPTER 2
EXPERIMENTAL METHODS
2.1 Materials
The potassium ethyl xanthate (KEtX) was prepared by a method described by Little and Leja (44) and Bulmer and Mann (45). Sodium ethoxide was;prepared by dissolving one mole (40 gms) NaOH in 500 mis
ethanol. The product was cooled in an ice bath and one mole (76 gms)
CS^ was added dropwise with stirring while maintaining the temperature
at less than 10°C. The solid ethyl xanthate product was filtered,
recrystallized three times from acetone, then washed with diethyl
ether and dried in vacuum. The final product analysed 98+% purity
by iodine titration.
Diethyl dixanthogen ((EtX^) was prepared by the addition of aqueous
iodine to neutral aqueous KEtX solutions, and the milky suspension
was extracted with either hexane or diethyl ether. The organic
layer was washed several times with distilled water and the final
product was prepared by vacuum evaporation of the solvent. The melting
point of the solid yellow diethyl dixanthogen product was 32°C which
agreed with the value reported in the literature (1).
Double distilled water was deoxygenated by boiling in a narrow
necked flask and bubbling argon through the solution. The flask was
sealed with a rubber septum and allowed to cool under a positive argon pressure. The argon, "Linde" high purity grade, with specified oxygen content (< 5 ppm) was further purified by passing through a gas train consisting of silica gel (H^O removal), ascarite (CO^ removal), copper gauze heated to 300°C (0^ removal) and drierite (I^O removal). Oxygen saturated water was prepared from deoxygenated double distilled water by bubbling1"Linde" high purity oxygen through the solution for 2 to 3 hours. The "Linde" high purity oxygen was further purified by an ascarite (CC^ removal) and drierite gas train.
2.2 Ethyl Xanthate Decomposition Experiments
Decomposition runs were carried out in 1000 ml volumetric flasks that had been filled with deoxygenated double distilled water. The flask was sealed with a rubber septum and samples were removed with a
50 ml hypodermic syringe. Argon or oxygen was admitted into the flask, also by hypodermic needle, to maintain a positive gas pressure inside the flask during sampling. The flasks were stored in a covered constant temperature bath to exclude light from the samples.
2.3 Diethyl Dixanthogen Decomposition Experiments
200 ml aliquots of stock xanthate solution in deoxygenated double distilled water at 22°C and neutral pH were poured into a 250 ml flask in an argon glove bag. Known concentrations of (EtX)^ were then prepared by injecting small volumes (< 1 ml) of 0.1 M KI^ into the KEtX solution. pH adjustment of the KEtX solution to basic conditions was also made by injection of small volumes of strong KOH with a 1 ml syringe. The solution was stirred vigorously with a magnetic stirrer and samples were removed as required for ultraviolet spectroscopic analysis or for the
determination of (EtX)2 concentration.
The blank for the reference beam of the ultraviolet spectrophoto• meter contained Kl in a concentration equal to the concentration of KI^ used for preparing the (EtX^. By this method, the iodide absorbances at 225 my and 192 my were blanked out.
2.4 Ultraviolet Spectroscopic Analysis for Aqueous Ethyl Xanthate
and Carbon Disulfide
Stock solutions of KEtX were prepared by addition of weighed quantities of the fine crystalline powder to slightly basic (pH 7.5-
8.5) deoxygenated double distilled water. The absorption spectrum obtained on a Perkin-Elmer Model 450 Ultraviolet Spectrophotometer showed two strong peaks at 301 my and 226 my (Figure 1). Calibration of the two peaks showed that Beer's law was obeyed for these solutions
(Figure 2). The molar extinction coefficients were determined for the 301 my absorption (e^g-^ = 17,50D0O liteliter i mole ^"cm "*") and the 226 my -1 "I absorption (^26 = ^»^-^ liter mole cm ). These values compared with £301 = 17,750 liter mole "*"cm and ^26 = ^>^^® liter mole "*"cm ^ determined by Pomianowski and Leja (18) and e^oi = 17,460 liter mole "*"cm ^ and ^226 = ^>7^ liter mole "'"cm ^ as determined by Ma j ima (46).
The xanthate absorption in the ultraviolet region has been attri• buted to electronic transitions in the functional group of the xanthate molecule. Shankaranarayana and Patel (89) have tentatively assigned the 301 my absorption to a TT—TT transition and the 226 my absorption to a n-a transition based on an analysis of the behavior of each band to changes in the polarity of the solvent. In this representation
77 denotes a bonding and TT and antibonding TT orbital, n denotes a non- bonding electron localized on the sulfur of the dithiocarbonate group * 1 and a denotes an antibonding a orbital.
Analysis for carbon disulfide was also carried out spectrophoto- metrically by measurement of the C=S absorption at 206.5 my (Figure 1).
Calibration of the peak was obtained by injection of small volumes (.5,
.7, 1.0 yl) of liquid CS^ from a calibrated 5 yl syringe into liter flasks filled with distilled water. The flasks were completely filled and sealed. After the small droplets had dissolved (3-5 hours) the flasks were stirred and 1 cm cuvettes were filled for U.V. determination
The assignment of electronic transitions for the U.V. spectrum of xanthates using the method of Shankaranarayana and Patel is not definitive in giving the correct transitions. Dr. Bree, of the
Department of Chemistry, U.B.C. recommended that more exact methods must be used to ascertain the nature of the band absorptions. The additional information might be obtained by the following methods:
(1) experiments in which polarized ultraviolet light is directed on either single crystals of xanthate or frozen aqueous xanthate solu• tions. From the polarization angle of the emitted radiation, assign• ments of either a or TT bond absorption for each of the 301 my and 226 my frequencies can be made.
(2) Hiickel Molecular Orbital calculations would be required to define the probable electronic transitions and the relative energies of the TT—TT and n-a transitions. Comparison of the results of these calcu• lations with experimental data would establish the correct band assignments. Wavelength (my) Figure 1. Ultraviolet absorption spectra of aqueous (^ethyl xanthate, hydroxyl ion and carbon disulfide. of the CS^ absorbance.
The C=S molar extinction coefficient was found to be 72,000 liter mole ^cm ^. Previous evaluation by Pomianowski and Leja (18) had placed e between 60,000 and 70,000 liter mole ^cm CS 2
2.5 Ultraviolet Spectroscopic Analysis for Dixanthogen
Rapid and accurate methods for the determination of small quantities of dissolved dixanthogen did not appear to exist in the literature.
A method reported by Shankaranarayana and Patel (47) involved the reduction of dissolved dixanthogen by heating the test solution with
NH^NO^ and KCN, then oxidizing the resultant SCN with and back titration of the iodide produced with Na2S20.^ Recently, Mamiya and
Majima (48) have reported an analytical method wherein the dixanthogen is extracted from the aqueous solution with pyridine, and the optical density of the pyridine-dixanthogen extract is determined at 420 my 4 -1-1 (e = 1.59 x 10 liter mole cm ). Unfortunately, both analytical methods are subject to interference by the presence of metallic ions and were consequently not suitable for the bulk of the present testwork.
A new analytical method was developed based on the extraction of dixanthogen by hexane combined with ultraviolet spectroscopic analysis of the organic layer. Numerous organic solvents, such as chloroform, benzene, toluene, carbon tetrachloride, hexane and diethyl ether were tested, but only spectral grade hexane (Baker "Analysed", boiling point range 0.2°C) was found to have a suitable spectral window, a high extraction coefficient and did not affect the dissolved xanthates or metal xanthates that were present. A typical U.V. spectrum for (EtX)„ in hexane is given in Figure 3. The "Baker" hexane was found to have a useful window region to 215 mp. -4 -1
Beer's Law was obeyed up to concentrations of 10 mole liter
(EtX)2 in hexane (Figure 4). The molar extinction coefficient of the
two major peaks were also determined, eOQO = 8,600 liter mole ''"cm ^ and Zoo
£238 = ^-7,800 liter mole ''"cm "*".
The accuracy and reproducibility of the method was tested by
extraction of known quantities of (EtX)^ from KEtX solutions. (EtX)2 was produced stoichiometrically by the reaction between iodine and stock
KEtX solutions as given by the equation
2EtX + I3 > (EtX)2 + 31 (1)
A material balance between the quantity of iodine added, the moles
of EtX consumed and the moles of (EtX)2 extracted is given in Figure 5. —6
The limit of detectability of the method was estimated to be 10
mole liter ^ (EtX)2 in hexane, but the sensitivity could be increased by using a hexane:aqueous sample ratio less than one, with the
result that aqueous (EtX>2 concentrations as low as 1-x 10 ^ moles liter "*"
could be measured.
2.6 Light Scattering Method for Determining the Solubility of
Diethyl Dixanthogen
Turbidity measurements were used by Kakovskii, et.al. (49) to
determine the solubility of bis-dithiophosphate esters in aqueous
solution. Unfortunately, the experimental details of their method have not been reported.
Concentration (moles liter x 10 ). Figure 4. Calibration curve for diethyl dixanthogen in hexane vs. hexane reference. 10.0
8.0 .
T (fl 6.0 " d)s rH O e
CO/v 4.0 - O (EtX>2 determined from addition C o o
(EtX)^ determined from hexane 2.0 - extraction
—i— 1.0 2.0 3.0 4.0 5.0 6.0 7.0
(EtX)2 (moles x 10 )
Figure 5. Material balance between the concentrations of ethyl xanthate titrated with iodine
and diethyl dixanthogen determined by hexane extraction. The principle advanced by these workers was developed for measur• ing the solubility of diethyl dixanthogen in water. The method involved measurement of the turbidity of (EtX) solutions as the concentration of (EtX)2 reached saturation and then became supersaturated or turbid.
The various concentrations of (EtX)^ were produced by addition of known quantities of iodine to standard KEtX solution as given by equation 1. The maximum solubility was determined by extrapolating the turbidity (T) readings to zero turbidity (see Figure 6). The
technique was found to be very sensitive to small changes in concentra• tion at the solubility limit, and the experimental error was estimated as ± 0.5 x 10 ^ moles liter
The measurements were made with a Brice-Phoenix Model 2000 Light-
Scattering Photometer using a 3 cm cell. After addition of iodine
to the stock KEtX solution, the flask was agitated and a portion of
the solution was immediately transferred to the 3 cm cell. Readings for G(s) and G(w) were obtained from measurements of the galvanometer deflection at 0° and 90° to the incident light (436 mp). The absolute
turbidity (T) was calculated from the equation
T = k(G(s)/G(w))
The constant (k) included terms related to the instrument and operating
conditions, such as type of cell, solvent, refraction effects and
filters used.
Figure 6 illustrates three runs in which various concentrations of
Kl were added to determine the effect of the residual iodine on the 100 J
90
80
70 •H T3 •H X> 60 S-i
4J ' 50 -|
s 40 No- I added 30
4 20 - O 10 M KI
A 5 x 10 M KI 10 -
0.5 1.0 1.5 2.0 2.5 3.0 3.5 Concentration of (EtX) (moles liter 1 x 105)
Figure 6. Effect of iodide on the solubility of diethyl dixanthogen as determined by
light scattering, pH 7.5-7.9. solubility of (EtX^. Accurate measure of the concentrations of
(EtX)2 were obtained using the hexane extraction method (Section 2.5) and these results were compared with the concentrations of iodine added.
The [(EtX)2] in the solution measured from the two methods were in agreement to T 0.1 x 10 ^ mole liter.:"^.
2.7 Infrared Spectroscopic Methods
2.7(a) The KBr Pellet Technique
The KBr pellet method, developed by Stimson and O'Donnell (50) is now a standard method for preparing solid samples for infrared spectroscopic study. The KBr pellet die (Perkin-Elmer Model 186 - 0025) was loaded with a powder mix of 0.5 gm KBr containing 2.0 to 3.0 mg of
sample. The die was then evacuated for five minutes and pressed at
30,000 lb. load for one minute. The pellets were removed from the die
and suspended in the infrared beam with a special holder.
2.7(b) Infrared Cell for Aqueous Solutions
Ordinary infrared cells (NaCl) were not suitable for obtaining
spectra of aqueous xanthate solutions because of their solubility in water. KRS-5 (thallium bromide-iodide) was not suitable because of
the reaction between KRS-5 and xanthate to form thallium xanthate. A
special infrared cell was constructed from a bag of "Handi-Wrap" poly•
ethylene containing a 0.2 mm IR window spacer. The test solution was
injected into the bag which was then sealed to NaCl windows with "Nuj;ol"
hydrocarbon oil. This procedure eliminated the problems produced by
interference fringes when the bag alone was placed into the infrared
beam. The cell had a window region from 1400 cm ^ to 700 cm which was suitable for the detection of xanthate species.
2.7(c) Attenuated Total Reflectance (ATR) Spectra of Platinum
Surfaces
The development of the ATR technique as a tool for use in infrared spectroscopic studies was reported independently by Fahrenfort (51) and Harrick (52). The latter has recently published a book (53) which contains an exhaustive treatment of the theory and includes examples illustrating the wide use of ATR in the study of powders, fibers, paints and chemisorbed materials on surfaces. Tipman and Leja
(54) employed this method in a study of xanthate adsorption on evaporated
copper substrates.
The infrared spectra were obtained on a Perkin-Elmer model 521
Grating Infrared Spectrophotometer using a "Wilks Model 9" multiple
ATR attachment. The beam entered the KRS-5 prism (Figure 7) from the
45° bevelled face and was totally internally reflected since the angle
of incidence of the beam on the prism face was greater than the
critical angle for refraction. The size of the crystal was 5 cm x 2 cm
and approximately 25 reflections were required before the beam left the
crystal and was refocused into the IR spectrophotometer. Since the beam
could penetrate through the crystal surface a distance of about 1/10 of
the wavelength of the light used, it was possible to study the deposits
on poorly reflecting metal surfaces such as platinized platinum by
varying the pressure applied to the metal foil, and hence the proximity
of the foil to the metal surface. The reactivity of (EtX)2 toward TIBr (KRS-5 crystal matrix) was
tested in a separate experiment where the TIBr was contacted for 20 minutes with saturated aqueous (EtX)^ solution. No change in the
concentration of (EtX)2 was noted before and after the experiment. An
infrared pellet of 0.5 gm of the reacted TIBr powder showed no infrared
absorption bands. Infrared spectra of TIBr pellets containing 2.5 mg
(EtX)^ did not show any formation of thallium xanthate. These results
eliminated the possibility that the KRS-5 crystal surface may have
reacted with the (EtX)^ deposited on the surface of the platinum electrodes.
The platinum foil specimens were cut to suit the available crystal
area and were clamped into the cell immediately after removal from the
reaction vessel. No significant differences were noted between the
spectra obtained by this technique and those obtained from KBr pellets
or "nujol" mulls.
IR beam
crystal
Figure 7. Multiple internal reflection effect illustrating the penetra•
tion of the IR beam through the KRS-5 crystal surface. 2. 8 Electrochemical Methods
2.8(a) Cell for Rest Potential Measurements
The electrochemical cell used for the electrode potential measure• ments shown in Figure 8, followed a design given by Gardiner (58).
The electrical connection between the removable platinum electrode and the Ag/AgCl/4M KC1 reference electrode was made through a Luggin capillary containing 1 M KC1. The capillary minimized the effect of the possible IR drop through the solution. Potential measurements were made using a Keithley Model 610 B electrometer.
The cell also included pH electrodes which were connected to a
Corning Model 10 pH meter. A similar capillary connection between the solution and the Ag/AgCl/4 M KC1 reference electrode was used to eliminate the possibility of any contamination of the reference electrode by the xanthate solution. All potential measurements were corrected
to the normal hydrogeJ n electrode (E. ,. „^ .. „„. = + 0.1872V vs SHE). Ag/AgCl/4 M KC1
Test solutions were added to the cell with vigorous stirring by a magnetic stirrer. Potential and pH readings were recorded after they had remained constant for 5 minutes. Frequent measurement of the solution temperature showed that it remained between 24 i 1.0°C.
2.8(b) Cell for Electrode Polarization Measurements
The three electrode cell (Figure 9) used for electrode polarization measurements was constructed after the design of Tolun and Kitchener (39).
The platinum electrode (5 cm x 2 cm) could easily be removed from the cell for infrared analysis by the ATR method. The potential measurements using a Model 610 B Keithley electrometer were made through a Luggin capillary connected to an Ag/AgCl/4 M KC1 reference electrode as described previously. Current measurements were made with a Simpson ammeter connected in series with the polarization circuit.
Xanthate solutions were purged with purified argon for one hour in the auxiliary vessel and then introduced into the cell under argon atmosphere.
Electrode potential readings were taken after 1 minute of polar• ization. Concentrations of 1 M, 0.1 M and 0.01 M KC1 electrolyte were used in the xanthate solutions along with a corresponding concen• tration of KC1 only in the salt bridge and reference cell. No appreciable differences in shape of the polarization curves was noticed between the various electrolyte concentrations. Algon purge
K&dbthley Corning Model 10 Electrometer pH meter 1 <2>
1 Reference electrode \ Reference Electrode
Luggin f 1 M KC1 capillary solution
platinum foil
Figure 8. Apparatus for measuring pH and rest potential. Argon purge
Keithley Simpson Electro- Ammeter mete.
•reference electrode QaTt>bridge
Argon ~ r4 purga 11
-Platinum foil
Reference Cell Luggin Working Cell capillary
gure 9. Apparatus for polarization studies. CHAPTER 3
RESULTS AND DISCUSSION
3.1 Decomposition of KEtX in Aqueous Alkaline Solutions
Experiments on the alkaline decomposition of aqueous xanthates have been discussed by a number of workers and most consider that the decomposition is first order with respect to xanthate ion. However, the long reaction times have led to scattered results as the reaction approached completion and some doubt could be cast on the assumption that the first order rate was correct. Finkelstein (37) evaluated a number of first order rate constants for KEtX, but no information was given regarding the length of time the reaction was carried out. Rao and Patel (28) also studied the decomposition of aqueous ethyl xanthate solutions but were unable to fit their data to any rate law.
The results of Pomianowski and Leja (18) suggested that oxygen actively participated in the xanthate decomposition mechanism. However,
Finkelstein (37) was unable to demonstrate any significant difference between the decomposition rates for KEtX solutions under nitrogen or oxygen atmospheres, and he concluded that oxygen was not involved in the decomposition process. His study involved the use of NaH^PO^ and
Na2B^0^ pH buffering agents, and it could be argued that these additives affected the rate and the mechanism of decomposition.
This series of experiments were undertaken to determine the true order of the reaction and to evaluate the effect of oxygen on the decomposition rate without the presence of pH buffers. Two groups of experiments were carried out under identical conditions in a constant temperature bath (22 + 1°C) with the samples shielded from daylight.
The decomposition rates of 5 samples each of aqueous KEtX were observed in (1) purified argon atmosphere over deoxygenated double distilled water and, (2) purified oxygen atmosphere over oxygen saturated double distilled water.
The following measurements were made during the course of a run:
1. the absorbance of the 301 my band of the ultraviolet spectrum was used to determine the ethyl xanthate concentration.
2. the pH was measured
3. the diethyl dixanthogen concentration was measured by the hexane extraction method. (Section 2.5, not obtained for every sample).
Since volumes of aqueous sample were removed from the reaction vessel and the volume replaced with either purified argon or oxygen, the concentration of the gaseous products, notably CS2, could not be measured. A second group of experiments in which the reaction vessel was completely filled are reported in Section 3.1.5.
3.1.1 Determination of Rate Law
The data recorded during a typical run are given in Table 1.
Examples of concentration - time plots for two decomposition runs under each of argon and oxygen atmospheres are given in Figure 10.
The method of determining the true order of a reaction by the measurement of initial rates is described by Laidler (56). The slopes Decomposition of Ethyl Xanthate under Argon Atmosphere at 22°C, Series 1 #2.
Sample Time Absorbance Ethyl Xanthate pH Diethyl Number (hours) 301 my Concentration dixanthogen (moles liter~lx 10^) concentration
1 0 0.453 0.259 7.2 lO"7
2 42 0.418 0.239 7.2 — 3 110.5 0.348 0.199 7.3 — 4 160 0.338 0.193 7.5 lO"7 5 207 0.318 0.181 7.5 — 6 259 0.308 0.176 7.4 lO"7 7 313 0.281 0.161 7.6 — 8 408 0.253 0.145 7.6 lO"7
From the.least squares analyses for slope and intercept (see Appendix 1)
Sloped^ = -2.68 + .65 x 10_8 moles liter'^r"1
• -4 _i Intercept = .244 ± 0.004 x 10 moles liter 9.0
Ca) 8.0
7.0
CU •H 6. 0 iH CO cu O e 5. 0 o X
•oH J-l 4. 0 cfl
4—' c a C o o 3. 0 X 4-1 w w
2.0 Ca) • Solutions saturated with oxygen
Solutions under argon 1.0 J 00
r 1 1 1 r— —i— 50 100 150 200 250 300 350 400 Time in hours
Figure 10. Decomposition of aqueous' ethyl xanthate in Ca) solutions saturated with oxygen, (b). deoxygenated solutions under argon^atmosphere. of concentration-time plots are related to the concentration of a reactant by the equation
n rate = -d[EtX~L-j^—]- = k[Etn7X-v-i J
where k = rate constant
n = order of the reaction
By taking-the common logarithms of each side
log rate = log k + n log .[EtX-]
the order of the reaction was found from the slope of the double logar• ithmic plot of rate against initial KEtX concentration. The slope obtained for KEtX decomposition in argon was found to be 0.965
(Figure 11). Similarily, the result for KEtX decomposition in oxygen was 0.975 (Figure 12). Both values are very nearly 1.0 and confirm the decomposition reaction is first order in KEtX. The observed rate law can be written
_ dlgxj. = klEtx-]
3.1.2 Effect of Oxygen on the Decomposition Rate
The decomposition rate constants observed for KEtX solutions under argon or oxygen atmospheres were found to be equivalent (Tables 2 and 3).
The results are in agreement with the observations made by Finkelstein
(37) and summarized in Table 4. Decomposition Rates for Aqueous v Ethyl Xanthate Under Argon Atmosphere at 22°C.
Rate Initial log(-rate) log Initial First Order (moles liter~~-'-hr-"'' x 10^) Concentration Concentration rate constant _1 - 4 -1 4 (moles liter x 1(T) (moles liter ! x io ) k^hr x 10 )
-2.68 ± 0.65 .244 ± 0.009 -3.57 ± 0.11 -4.612 ± 0.015 10.9 ± 0.7 -7.05 ± 0.87 .984 ± 0.012 -3.15 ± 0.05 -4.007 ± 0.005 7.1 ± 0.9 -24.8 ± 3.3 3.44 ± 0.04 -2.61 ± 0.06 -3.463 ± 0.005 7.2 ± 0.4 -61.5 ± 9.3 7.23 ± 0.09 -2.21 ± 0.07 -3.141 ± 0.005 8.5 ± 1.0 -649 ± 44 94.6 ± 0.5 -1.15 ± 0.05 -2.024 ± 0.002 6.9 ± 0.5 log Co-
Figure 11. Log-log plot of the decomposition rate of aqueous ethyl xanthate vs. initial ethyl xanthate concentration for deoxygenated solutions under
an argon atmosphere. Decomposition Rates for Aqueous Ethyl Xanthate under Oxygen Atmosphere at 22°C
Rate Initial log (-rate) log Initial First Order (moles liter_lhr_l x 10 ) Concentration Concentration rate constant (moles liter"1 x.104) k (hr-1 x 104) 1
2.15 ± 0.30 .207 ± .004 -3.67 ± 0.06 -4.684 ± 0.008 10.4 ± 0.3 6.98 ± 0.62 .944 ± 0.009 -3.16 ± 0.04- -4.025 ± 0.004 7.4 ± 0.7 26.6 ± 2.0 3.66 ± 0.02 -2.58 ± 0.03 -3.437 ± 0.003 7.3 ± 0.2 i 72.0 ± 5.0 9.71 ±,.0.06 -2.14 ± 0.03 -3.012 ± 0.003 7.2 ± 0.6 ^ 754 ± 40 95.1 ± 0.5 -1.12 + 0.02 -2.022 ± 0.003 7.9 ± 0.5 ' -5.0 -4.6 -4.2 -3.8 . S3..4 ' -3.0 -2.6 -2.2 -1.8
log Co
Figure 12. Log-log plot of the decomposition rate of aqueous ethyl xanthate vs. initial ethyl xanthate concentration for solutions saturated with oxygen. Complete exclusion of oxygen from the reaction flask was not expected in the experiments using argon. However, the procedure employed in deoxygenating the water was designed to reduce the oxygen content of the water (normally 43 mg/1 at 22°C) by at least two orders of magnitude. Since no variation in rate constant between argon and oxygen solutions was observed, it was concluded that oxygen was not involved in the rate controlling steps of the decomposition process.
3.1.3 Effect of Hydroxyl Ion on the Decomposition Rate
The pH of the samples taken during the course of the runs for xanthate decomposition under argon atmosphere are given in Figure
13. The results for decomposition under argon and oxygen atmospheres are presented in Table 4 where comparison with the values reported by
Finkelstein can be made. The observed rates obtained in unbuffered solutions are in general agreement with the results obtained by
Finkelstein, and the constancy of the rates indicate that the reaction rate is unaffected by [OH J in the region pH 7 to 11.
Since the samples were not buffered, it was necessary to take the pH measurement immediately after the sample was withdrawn from the reaction flask. After a short time (less than 5 minutes) the absorption
of C02 from the atmosphere caused the pH to fall rapidly and such readings would be unreliable.
During the course of a run, the pH increased- indicating that OH was a product of the decomposition reaction. In Table 5, the change in IOH J has been related to the decrease in lEtX ]. A comparison of [EtX ] = 94.6 x 10 M o
o- [EtX ] = 7.23 x 10 4M
[EtX ] = 3.44 x 10 4M o
EtX ] = 2.58 x 10 M o
[EtX ] = 0.98 x 10
i —i— —I— —1— 100 150 200 250 300 350 time in hours Figure 13. pH shift of ethyl xanthate decomposition runs in deoxygenated water under argon atmosphere. the figures shows a general correspondence between the overall pH increase and xanthate decrease, but because of the difficulties in pH measurement it was not possible to obtain a quantitative relationship between [EtX ] decomposed and [OH ] produced.
3.1.4 Evaluation of Reaction Products
The measurement of diethyl dixanthogen concentration during the KEtX decomposition experiments revealed that this compound was not a major decomposition product (Table 1) for KEtX in either deoxygenated water or oxygen saturated water. has been recognized as a major decomp• osition product along with possibly other gases such as COS or H^S, and quantitative data were obtained for these products.
The gaseous products from the decomposition were bubbled
through CCl^, and an infrared spectrum of the resultant CCl^ solutions was taken; Only one major peak at 1515 cm \ attributable to CS^
absorption, was observed.
50 ml of the gas was injected into a Pye Gas Chromatograph which was equipped with a 10 foot column containing a neutral substrate.
The single peak was confirmed as CS^ from the retention time in the
column. No other peaks were observed.
3.1.5 Quantitative Measurement of Carbon Disulfide
In slightly alkaline solutions (pH 7-9) the absorbance of OH in
the ultraviolet region (400 my to 195 my) is negligible. Under these
conditions, the aqueous CS^ concentrations were determined directly
from the absorbance at 206.5 my (e = 72,000 liter mole "''cm "*", Section
2.4). Effect of pH on the Decomposition of Ethyl Xanthate
Literature Values Present Study
1 1 Atmosphere Temp. pH k (hr ) Atmosphere Temp. pH k1(hr'- )
Air 22±3°C 6 8.8 x .104 (1)
Nitrogen 25° 6 185 X 10_4(2)
4 7 26.5 X lO"4 Argon 22±2°C 7.0 7.1 ± 0.9 x 10" -4 -4 8 13.3 X 10 8.0 ' 7.2 ± 0.4 x 10 -4 -4 9 12.2 X 10 9.0 8.5 ± 1.0 x 10 -4 -4 10 12.1 X 10 10.5 6.9 ± 0.5 x 10 -4 11 12.4 X 10
Oxygen 25° 6 185 X lO"4 Oxygen 22+2°C 6.9 7.4 ± 0.7 x 10~4 -4 7 28 X 10 -4 8 13.3 X 10 -4 -4 9 12.3 X 10 9.5 7.2 ± 0.6 x 10 -4 -4 10 11.1 X 10 10.5 7.9 ± 0.5 x 10 -4 11 12.3 X 10
(1) Pomianowski and Leja (18). (2) Finkelstein (37). Relationship Between Ethyl Xanthate Decomposition and pH Increase
Decomposition Under Decomposition Under Argon Atmosphere Oxygen Atmosphere
KEtX Decomposed OH Increase Xanthate Decomposed OH Increase (moles liter ^) (moles liter "*") (moles liter (moles liter ^)
-5 -6 -8 1.14 x 10 5 x 10 8.5 x 10 6 x 10
-5 -7 -5 5 2.88 x 10 4 x 10 2.8 x 10 10" -5 -5 -5 -5 8.5 x 10 5.0 x 10 9.0 x 10 4 _ x 10 -4 -4 -4 -4 2.20 x 10 1.2 x 10 2.5 x 10 1.3 x 10 -3 -3 -3 -4 2.24 x 10 1.3 x 10 2.5 x 10 5 • x 10 CS^ is sparingly soluble in water (ca. 2 gms/1 at 22°C (1)) and a new series of decomposition runs were undertaken in quartz U.V. cells fitted with teflon plugs. The cells were filled with aqueous
KEtX solution so that no gas volume was present, then stoppered and sealed with wax. The rate of decomposition of KEtX and the evolution of
CS^ were observed. In Table 6, the material balance calculations showed that CS^ evolution corresponded to the xanthate ion decrease. The rate of CS^ formation was unaffected by the oxygen dissolved in the water.
3.1.6 Mechanism of Xanthate Decomposition in Alkaline Solution
Most workers agree that CS^ results from aqueous decomposition of xanthate, but a large number of other products have also been proposed.
Glembotskii (57) reported K^CS^ and K^CO^ as other major products, but he was unable to confirm their presence. Philipp and Fichte (32) described a decomposition mechanism as due to a slow hydration of the carbon sulfur bond in the xanthate which resulted in an intermediate of the structure
S~ I R-O-C-SH I OH
The hydration product was unstable and decomposed into other products, namely carbonate, hydrosulfide, dithiocarbonate and trithiocarbonate.
Hovenkamp (33) concluded from experiments with ethyl xanthate in dilute alkaline solution that: 0.7
0.6
4-1 0.5 • d[E^ ] = -4.38 + 0.8 x 108 mole liter 1hr 1
in r-l o e, 0.4 CM o
C n)
0.3
o
t-i 0.2 C o d[CS2] 8 ii C 1 O dt =4.Ix 10°+ 0.4 mole liter V u 0.1 •
i I —I— —I— 50 100 150 200 Time in hours
Figure 14. Decomposition of ethyl xanthate in deoxygenated water in a sealed cell, pH 7.4.
(a) rate of xanthate decomposition, (b) rate of [©S? formation. Material Balance Between Ethyl Xanthate Decomposed and CS^ Evolved
from Sealed Cell Measurements
Condition Initial First Order [EtX ] [CS ] Evolved Ratio
pH Rate Constant Decomposed [CS2]/[EtX ] (moles liter ) _1 (moles liter ) (hr )
4 6 -6 Deoxygenated 7.2 7.1 ± 0.8 x 10 , 8.0 x 10 7.3 x 10 0.91 -4 -5 -5 7.4 7.9 ± 1.6 x 10 1.77x 10 1.78 x 10 1.00
-4 -6 -5 Oxygen Saturated 7.5 7.0 ± 0.5 x 10 7.3 x 10 6.4 x 10 0.87 -4 -5 -5 7.3 8.8 ± 0.8 x-10 1.19x 10 1.08 x 10 0.91 "As to the decomposition of ethyl xanthate, in our opinion a reaction in which carbon disulfide is liberated has to be considered to be the main reaction."
+ 'C2H5OCS2 + H * C2H5OH + CS2"
Recently Dautzenburg and Philipp (34) postulated that the alkaline decomposition of ethyl xanthate can be written
C2H5OCS2 + H20 » C2H5OH + • CS + OH
in the region pH 8 to 13. Above 1 M NaOH, the CS2 was considered to recombine with the excess OH to produce a dithiocarbonate ion which decomposed rapidly to carbonate and sulfide according to the scheme
+ CS2 + OH • CS2OH —l CS20 + H
CS20 + 20H » C03 + HS
The results presented in this section are in complete agreement with Dautzenburg and. Philipp (34) for the region pH 8 to 13.
Experimental evidence in support of this reaction scheme can be summarized into the following points:
1. the decomposition rate is first order with respect to ethyl xanthate ion, indicating that only one molecule of ethyl xanthate is involved in the rate determining step.
2. the reaction was found to be unaffected by oxygen or base. 3. the increase in pH indicated that OH was product of the reaction.
4. CS^ was the only sulfur containing product observed during the reaction, and a quantitative measure showed that one mole of CS^ was produced for each mole of ethyl xanthate decomposed.
3.1.7 Decomposition of KEtX in the Presence of Possible Catalytic
Agents
Xanthate decomposition experiments under argon and oxygen atmos• pheres at neutral to slightly basic conditions (pH 7 to 9) were carried out in the presence of three possible catalytic agents
1. methylene blue
2. platinized platinum _3 3. ferrous sulfate (10 M).
3.1.7(a) Methylene blue
Finkelstein had reported that an equilibrium electrochemical potential between ethyl xanthate and diethyl dixanthogen could be measured with a redox indicator such as methylene blue. The possibility that (EtX)^ would be catalytically formed by methylene blue as an alternate product to the hydrolysis reaction was investigated.
Under argon atmosphere, the decomposition rate was slower than similar observations made for ethyl xanthate decomposition without methylene blue (Table 7). However, under oxygen atmosphere, EtX decomposed rapidly and only about 10% remained after 30 minutes. A true linear first order rate was not observed, but because of the rapid Decomposition of Aqueous Ethyl Xanthate in the Presence of Possible Catalytic Agents
Description of Run Temp. Initial Initial Xanthate First Order pH Concentration Rate Constant [EtX~l (hr"1)
-5 Methylene Blue (2 x 10 M)
9 -5 -4 1. Oxygen free 23 C 8.1 8.7 x 10 2.3 ± 0.5 x 10 -5 2. Oxygen saturated 23°C 8.5-6.3 6.8 x 10 400 ± 100
Platinized Platinum -5 -4 1. Oxygen free 23°C 8.0 5.6 x 10 9.7 ± 0.7 x 10 -5 -4 2. Oxygen saturated 23°C 7.6 7.3 x 10 15.7 ± 1.0 x 10
-3 Ferrous Sulfate (10 M) -4 23°C 8.3 1.0 x 10 3.6 ± 0.5 x 10" 4 1. Oxygen free -5 -4 23°C 8.6 8.5 x 10 42. ± 0.5 x 10 2. Oxygen saturated reaction the concentrations of EtX were difficult to measure accurately.
For comparison, the approximate first order rate constant was 400 hr or six orders of magnitude faster than the decomposition rate of xanthate in oxygen free water. No (EtX)^ could be found by hexane extraction either during or at the end of these runs.
Methylene blue, or 3,9-bis-dimethylaminophenazothonium chloride contains a positive sulfur and'four amine groups. The ethyl xanthate ion was presumbly capable of combining with the positively charged sulfur atom to form a weak disulfide which undergoes decomposition reactions in the presence of oxygen. Another possibility was the complexing of the xanthate ion with the amino groups, resulting in the formation of adducts which may also have undergone rapid decomposition.
The result pointed out that oxidation-reduction indicators cannot be used with xanthates unless care is taken to insure that the indicator does not decompose the xanthate.
3.7.1(b) Platinized Platinum
The decomposition rate of KEtX in water containing platinized 2 platinum (20 cm geometrical surface area in 900 ml R^O with strong agitation) was observed under both argon andAoxygen atmospheres (Table
7). Each value is a composite of two runs, where a different surface pretreatment for the platinum was used for each run. In the first run, the platinum was platinized, cleaned by cathodic polarization where H. was evolved from a dilute H„S0. solution, then washed with I 2 4 copious quantities of dilute KOH and distilled water and then placed in the reaction flask. The second platinum pretreatment involved cleaning in aqua regia, then dilute KOH. and distilled water. No differences in the decomposition rates were noted and the two results for each run in both argon and oxygen saturated solutions are plotted together.
The rate of decomposition is approximately doubled in oxygen saturated
KEtX solutions over similar rates found for solutions not containing platinum. This rate, however, is still slow and it appears that platinized platinum is not an effective oxidation catalyst.
The diethyl dixanthogen content of the solution, determined after 30,80 and 120 hours showed no significant concentration present. After completion of the run, the platinum foil was removed from the reaction vessel and washed with hexane. The U.V. spectrum of the extract showed no evidence of diethyl dixanthogen.
3.1.7(c) Ferrous Sulfate
The possibility that ferrous and ferric ions, 2 and 30 hours in the present experiments revealed no measurable quantities (< 10 mole liter ^) at these time intervals. The possible catalytic agents for converting xanthate to dixanthogen in the presence of oxygen did not yield fruitful results. However, I j [ | | | other metallic ions such as Pb , Cu , Zn could significantly accelerate the decomposition rate, but this was not investigated. The experiments on the behavior of aqueous xanthate and oxygen were carried out in a homogeneous environment, and thus cannot be applied directly to flotation systems. The presence of sulfide minerals creates a heterogeneous environment in which the catalytic oxidation of xanthate to dixanthogen is likely. The subject of catalytic oxidation will be discussed further in the Summary and Conclusions chapter of this thesis. 3.2 Decomposition of Dixanthogen In Aqueous Alkaline Solution In the previous tests, the oxidation of ethyl xanthate was carried out under argon and oxygen atmospheres and the products were shown not to be diethyl dixanthogen, but CS^ an^ OH formed by a hydrolytic reaction. To further define the chemical nature of the aqueous xanthate-dixanthogen system, the reaction of hydroxyl ion and other reagents with diethyl dixanthogen was studied. It was tentatively believed that the reduction of diethyl dixanthogen with base would result in non-stoichiometric production of xanthate, i.e. one mole of dixanthogen reacted with OH would not produce two moles of xanthate. This formed the initial part of the study; the second part was the measurement of the rate and order of the reduction reaction with respect to each of the reactants so that a mechanism could be postulated. 3.2.1 Stoichlometry of Reaction of Diethyl Dixanthogen with Hydroxyl Ion Initially, stock solutions of known quantities of ethyl xanthate and diethyl dixanthogen were prepared by the addition of KI^ to neutral KEtX solutions. A small volume of concentrated KOH was added to pHi>ll and the flask was stirred vigorously until the reaction was estimated to have gone to completion. At this time, the concentration of remaining CEtX)^ was measured by the hexane extraction technique (Section 2.5) and the EtX concentration was measured by ultraviolet absorption (Section 2.4). The molar ratio of the amount of EtX generated from the amount of (EtX)2 decomposed was then calculated. The result of 18 reactions between OH and (EtX)^ are given in Table 8. The value for the molar ratio clearly shows that one mole of xanthate was generated for each mole of dixanthogen that was reduced. 3.2.2 Elimination of Possible Interference by Iodide Since the controlled amounts of (EtX)2 were prepared by titration of the KEtX solution with KI^, there was some question that the residual I could have interfered with the results. Three hydrolysis runs (Series 2) were carried out with stock solutions of (EtX^ pre• pared by agitating liquid (EtX)^ in deoxygenated double distilled water, which eliminated iodide ion. These results were equivalent to those obtained by iodine titration. Further evidence was provided in tests -4 -3 in which the reduction of (EtX)^ was attempted in 10 M and 10 M Kl. Under these conditions, the rate was shown to be sufficiently slow to be negligible compared with the rate of base reduction (Table 9). Reaction of Diethyl Dixanthogen with Hydroxyl Ion at 22° C Run # Atmosphere Run Time pH [(EtX>2] [EtX~] Molar Ratio hydrolysed recovered [EtX~]/[(EtX) ] -15 -15 (moles liter x 10 )(moles liter x 10 ) Series 1 #1 Air 15 12.15 2.60 2.59 1.00 #2 Air 15 11.92 0.45 0.49 1.09 Series 2 #1 Argon 45 11.05 1.35 1.54 1.14 #2 Argon 15 11.85 0.45 0.44 0.98 #3 Argon 15 11.58 1.00 1.05 1.05 Series 3 #1 Oxygen 85 10.35 1.10 1.12 1.02 Series 4 #1 Air 45 11.70 2.20 2.15 0.98 #2 Air 45 11.25 2.20 2.31 1.05 #3 Air 45 11.53 2.20 2.25 1.02 #4 Air 30 11.97 1.50 1.50 1.00 #5 Air 30 12.12 1.50 1.50 1.00 Series 6 #1 Air 40 11.40 1.12 1.10 0.98 #2 Air - 60 11.08 1.08 1.12 1.04 #3 Air 45 12.05 0.95 1.02 1.07 Series 7 #1 Air 60 10.84 1.10 1.15 1.05 #2 Air 50 11.15 1.05 1.01 0.91 #3 Air 50 11.17 1.00 0.95 0.95 #4 Air 50 11.18 0.95 0.98 1.03 Mean value for molar ratio [EtX ]/[(EtX) ] = 1.02 ± 0.05. 3.2.3 Evaluation of Reaction Products The reaction of (EtX)^ with OH resulted in the formation of CS^ which could be measured qualitatively by following the ultraviolet absorbance at 206.5 my. Unfortunately, at pH >10, the OH absorption became significant and the CS^ band was masked. Consequently, quanti• tative evaluation of CS^ concentrations by this method was not possible. Extraction of the reaction products into organic solvents was investigated, but this resulted in simultaneous extraction of (EtX)^ and did not yield fruitful results. Because of the low concentration of products, an infrared spectrum of the decomposed aqueous solution was not possible using the solutions at which the U.V. experiments were made. A separate experiment was devised whereby 5 N KOH was agitated with liquid (EtX)^ for 24 hours in a sealed flask. The reacted solution was centrifuged to remove any excess (EtX)^ and an infrared spectrum was obtained from the remaining solution using the cell described in Section 2.7(b). The infrared spectrum over the region 1400 cm to 700 cm yielded -1 -1 -1 three strong peaks at 1060 cm , 1110cm and 1040 cm which correspond with the absorption peaks at 1058 cm \ 1110 cm and 1040 cm ^ for EtX ions previously reported by Poling, (5) (see Figure 28). A U.V. absorption spectrum of the solution showed that approxi• mately 1.6 moles of EtX was produced for each mole of (EtX)^ hydrolysed. The discrepancy between.this result andithe previous experiments reported in Table 8 can be explained by considering that the hydrolysis was carried out in a strongly basic medium and some of the CS„ produced during the hydrolysis was likely to recombine with the alcoholate in solution and regenerate ethyl xanthate ions. The only reaction products that were detected were CS^ and EtX . 3.2.4 Contribution of the React ion Products to the 301 my Xanthate Absorption In the previous experiments on ethyl xanthate decomposition, the pH was less than 10 and the possibility of reaction between the product CS^ and OH was avoided. In the diethyl dixanthogen decomposition runs, the pH was greater than 11, and the possibility of other products contributing to the 301 my xanthate absorption band was considered. Rovenkamp (57) lists the reactions between CS^ and OH as follows: + CS„ + 0R~ y CSo0H y CSo0 + H 2 2 -< 2 CS20 + 20H y C03 + 2SH H2CS20 y H2S + COS rapld = COS + 30H~ > SET + C03 + H20 CS„ + SH y HCS_ y CS„ + H+ 2 3 -« 3 CS2 + CS20 y CS3 + COS The U.V. spectra of CS20 , SH , CS2, CS3 , CS^ , CS30 and ethyl xanthate have been reported (33) and are given in figure 15. The contribution of any of these products to the 301 my absorbance of ethyl xanthate does not appear to be significant, and for the present experiments,1 301 my band was considered to be due solely to absorption by the ethyl xanthate ion. toooo a '\cs2o"' ImoC'crri1 ® CD-/ \ SHV a (cs,) / 5J300 \ III 40 w 20- a .1 ImorW (D 15JDO0- \ / r® 10000 X V, Xanthate 5*00 n. ImoC lyroo 10JQO0 5TJ00 400 380 360 340 320 300 260 260 240 mi figure 15. Ultraviolet spectra of sulfur containing compounds in the system CS2~NaOH-alcohol. (1) dithiocarbonate in 5 N NaOH; (2) CS, (3) Na„S in 0.1 N NaOH; (4) ethyl monothiocarbonate (C^OCOSK) ; (5) trithiocarbonate (CS^) ; (6) potassium ethyl xanthate (C^OCSSK) ; (7) tetrathiopercarbonate - (CS^ ) ; (8) trithiopercarbonates(CS30~); (after Hovenkamp (57)) 3.2.5 Mechanism of the Decomposition of Dixanthogen The hydrolysis of diethyl dixanthogen with OH can be considered to be a specific example of the reactions between disulfides and negatively charged bases or nucleophilic reagents. In a review paper by Davis (58), the general form of these types of reaction was written: XS-SX + Y~ y XSY + XS~ The reaction was considered to proceed by an S^2 or nucleophilic displacement mechanism. Such "a mechanism was first proposed by Foss (59), who considered that all disulfides, including dixanthogen undergo polar fission according to the scheme XS-SX «—y (XS+) (SXT) Foss postulated that attack by nucleophilic agents stronger than the xanthate ion would result in displacement of the xanthate and the forma• tion of a new disulfide adduct. The reaction between diethyl dixanthogen and OH can be illustrated following the mechanism proposed by Foss (59). The initial step of the displacement reaction can be written S S S S M II _ II H _ ROCSSCOR + OH y ROCSOH + ROCS One of the products, a sulfenic acid, is known to be highly unstable. Danehy (60) has observed that proteins are disproportionated by base hydrolysis and that alkyl disulfides undergo rearrangement and decomp• osition until alkyl sulfonic acids are formed. The mechanism of xanthate sulfenic acid decomposition has not been explained, although Danehy (60) has explained that only dialkyl disulfides undergo oxidation to sulfonic acids, and all other disulfides undergo disproportionation to a variety of products. Carbon disulfide was detected as a product of (EtX)^ hydrolysis from the U.V. absorbance at 206.5 my. Attempts at quantitative measurement were unsuccessful because of interference by the OH absorbance which becomes significant above pH 10. The disproporation- ation of xanthate sulfenic acid was considered to proceed by the following reaction S ll ROCSOH ROH . + CS2 + OH 3.2.6 Reaction of Dixanthogen with Other Nucleophiles The reactivity of bases toward the sulfur-sulfur bond has been discussed by a large number of workers (61,62). Foss considered that the displacement of a sulfide anion with another thiol is a measure of the base strength of the thiol. The displacement of thiocyanate, o-o-dialkyl dithiophosphate, xanthate and mercaptide ions from bis (o- nitrophenyl)disulfide illustrated the increasing order of base strength of each nucleophile and was discussed in terms of S-nucleophilicity. The order of reactivity of some nucleophiles toward the disulfide bond has beencreported by Davis (58) and is given below C^S > CgHgS > CN > 0H~ »02N-^^-S » N3 >> SCN~ In general, the reaction between thiols and disulfides has been considered to be reversible and has been written RSSR + R.S y RS-SR.. + RS 1 •* 1 Such a reaction between diethyl dixanthogen and thiosulfate ion was postulated by Foss (59) and can be written S S S S II II = II _ II _ ROCSSCOR + S203 ROCSS203 + ROCS S S II = II _ ROCSS„0„ + S.0- y ROCS + S.O, 2 3 2 3 -< 4 6 The kinetics of this reaction have been studied by..-Wells (24). The reaction of disulfides with non-sulfur nucleophiles such as CN and OH has been shown to result in unstable adducts which de• compose to other products and equilibrium reactions have not been observed. Cambron (2) observed that the reaction of CN with diethyl dixanthogen resulted in the formation of thiocyanate (SCN ) and xanthogen monosulfide. The reaction between disulfides and OH have been reported by Danehy (60) and are summarized in Section 3.2.5. Experiments on the reaction between;- diethyl dixanthogen and thio• sulfate ion, sulfite ion, and iodide ion are reported in Table 9. The reduction of (EtX) with K2S203 would have shown 2.0 moles of xanthate produced for each mole of (EtX)2 decomposed had the reaction gone to Stoichiometry of the Reaction Between Various Nucleophiles and Diethyl Dixanthogen Nucleophile Concentration pH [ (EtX) ] Reacted [EtX ] Recovered Molar Ratio - [EtX ]/[(EtX)2] 3 -5 -5 s2o3= 5 x 10 9.10 1.39 x 10 2.51 x 10 1.81 4 = -5 -5 so3 lO" 7.60 1.75 x 10 1.62 x 10 0.92 3 I- lO" 6.85 0.74 x 10-5 0.7 x 10-6 0.10 4 I- lO" 6.90 0.74 x 10-5 not detected 0.00 completion as postulated by Foss (59). In the case of sulfite reduc• tion, the inability of the sulfite anion to dimerize probably resulted in disproportionation of the adduct, and the formation of CS^ and other reaction products. Since no reaction was observed between iodide and diethyl dixanthogen, the iodide ion has been considered to be a weaker base than the ethyl xanthate ion. From these results and an order of base strengths for nucleophiles given by Davis (58), an order for the base strengths of these nucleophiles can be postulated. S = = CN~> 0H~ > S03 > S203 > ROCS" > SCN~ > i" 3.3 Rate Studies on the Reaction Between Dixanthogen and Hydroxyl Ion The reaction between diethyl dixanthogen and hydroxyl ion was postulated to proceed by a bimolecular displacement mechanism (or S^2 mechanism). Studies on S^2 displacement reactions in carbon systems have been reviewed by Ingold, who demonstrated that the incoming group enters on the "backside" of the substrate and forms a weak bond with the substrate ion. The timing of the bond breaking and bond forming process requires that both the attacking group and substrate must be in close proximity (or collision) before the reaction can occur. The concentration of the intermediate is directly proportional to the number of bimolecular collisions between them which in turn is proportional to the rate constant. The general rate expression for a bimolecular reaction can be written -d[(EtX)2] m n = k2[(EtX)2] [OH ] dt In the case of an S^2 displacement mechanism, the reaction becomes first order in [(EtX)2] and [OH ], or m = n = 1. To provide evidence for the proposed mechanism, the order of the reaction with respect to each species was determined experimentally. 3.3.1 Reaction Order with Respect to Diethyl Dixanthogen Concentration The method of initial rates (56), which was used to determine the order of the ethyl xanthate decomposition reaction was not suitable for evaluating the order of decomposition of diethyl dixanthogen. Since the reaction could be followed to 80% completion, and the [OH ] remained constant, the order of the reaction could be determined from the integrated form of the assumed first order rate expression. Consider m = 1 d.[ (EtX) 2] n = k2I(EtX)2] I0H ] dt with JOH ln constant -di (EtX) 2] n - k [(EtX)2] kx = k2 JOH J dt -d ln[(EtX)2] dt ln[(EtX)2] = C - kx t In Table 10, the results of a typical diethyl dixanthogen hydrolysis run are given. The [(EtX)^^ was calculated indirectly from measurements of [EtX ] . Since the stoichiometric runs had shown that one mole of xanthate was formed for each mole of (EtX) reacted _ [(EtX) ] = [(EtX).] - : [EtX ] + [EtX ] l t z o t o [(EtX)2^ was determined by hexane extraction at the beginning of each run. The contribution of (EtX)2 to the 301 my absorbance was estimated from the height of the EtX baseline and subtracted from the absorbance. Semilogarithmic plots of [(EtX)2^ vs. time are given for all Series #4 runs in Figure 16. A linear plot was obtained over 90% of the reaction and established that rate was first order, consistent with the postulated mechanism. 3.3.2 Reaction Order with Respect to Hydroxyl Ion Concentration The order of the reaction with respect to hydroxyl ion was evaluated by considering the rate of decomposition of diethyl dixanthogen at a number of pH values -d[(EtX)2] n = k2[(EtX)2]IOH J dt -d In I (EtX)2] n = kx = k2 [OH ] dt log k± = log k2 + n log [OH ] = log k_ - n pOH Hydrolysis of Dixanthogen with Hydroxyl Ion (Series 4 #3) Sample No. Time pH Absorbance [EtX ]fc [(EtX)2]fc=[(EtX)2]q-[EtX ]t+[EtX ] log[(EtX) ] (minutes) 301 my -.4 . 1 _ -1 ..4 , ,.^ -1 x 10 mole lxter x 10 mole liter 1 0 11.54 1.010 0.577 0.210 -4.678 2 2.00 - 1.119 0.639 0.168 -4.775 3 4.00 - 1.168 0.677 0.140 -4.854 4 6.84 11.53 1.244 0.711 0.097 -5.013 5 9.17 - 1.263 0.722 0.086 -5.066 6 12.00 - 1.304 0.745 0.062 -5.208 7 14.30 - 1.326 0.758 0.050 -5.301 ON 8 17.50 11.53 1.353 0.773 0.034 -5.469 9 21.00 - 1.372 0.785 0.023 -5.638 10 26.50 11.53 1.391 0.795 0.013 -5.886 2 -1 From least squares analysis , k = 4.5 ± 1.4 x 10 min OpH 11.25 QpH 11.53 ApH 11.70 •pH 11.97 • pH 12.-12 3x10 2xH?0 10"64 5 10i 15 20i 25 10 20 Time in minutes Figure 16. First order rate plot for the reaction of aqueous diethyl dixanthogen with hydroxyl ion, Series 4. = log k. - n ,(pKr - pH) z w = (log k0 - n pK ) + n pH z w Tabulated values of k^ and calculated values for k^, along with the error at 95% confidence limits are given in Table 10. A plot of log k^ vs. pH, given in Figure 17, yielded a slope of 0.995 ± 0.08 which confirmed that n = 1 and the rate expression was written -d[(EtX) ] = k2[(EtX)2][OH ] which was correct for a S„2 displacement reaction. A model for the reaction intermediate between diethyl dixanthogen and hydroxyl ion is drawn below; following examples for such inter• mediates given by Gould (65). S II S-C-OR s II R-O-C OH The model has been shown to be consistent with similar models drawn for S^2 reactions for aliphatic systems (65). Similar intermediates would be expected for the reaction between (EtX)2 and other nucleophilic reagents such as CN , , and SO^ all of which are stronger bases than the ethyl xanthate ion, and would displace ethyl xanthate ion from diethyl dixanthogen. Evaluation of Second Order Rate Constants for Hydrolysis of Diethyl Dixanthogen Run No. pH pOH k, x 10 min log k.. log k9 k„ liter mole min Series 4 #1 11.70 2. 30 7. 4 + 1.0 -1. 13 + 0.06 1.17 + 0.06 14.0 + 2.0 #2 11.25 2. 75 2. 9 + 0.1 -1. 53 + 0.02 1.12 + 0.02 13.2 + 0.6 #3 11.53 2. 47 4. 5 + 1.4 -1. 34 + 0.01 1.13 + 0.01 13.5 + 0.3 #4 11.97 2. 03 11. 9 + 0.4 -0. 92 + 0.01 Lll + 0.01 12.9 + 0.3 m 12.12 1. 88 18. 2 + 1.9 -0. 74 + 0.05 1.14 + 0.05 13.8 + 1.7 Series 5 #1 11.30 2. 70 3. 1 + 0.6 -1. 51 + 0.08 1.19 + 0.08 15.5 + 3.0 Series 6 #1 11.40 2. 60 3. 5 + 0.2 -1. 46 + 0.03 1.14 + 0.03 13.8 + 1.0 #2 11.08 2. 92 1. 44 + 0.08 -1. 84 + 0.02 1.08 + 0.02 12.0 + 0.6 Series 7 #1 10.84 3. 16 0. 92 + 0.04 -2. 09 + 0.02 1.12 + 0.02 13.2 + 0.6 #2 11.15 2. 85 2. 1 + 0.2 -1. 69 + 0.04 1.14 + 0.04 13.8 + 1.4 #3 11.17 2. 83 2. 1 + 0.1 -1. 68 + 0.02 1.15 + 0.02 14.2 + 0.6 #4 11.18 2. 82 2. 2 + 0.2 -1. 66 + 0.03 1.16 + 0.03 14.5 + 1.0 3.4 Solubility of Diethyl Dixanthogen in Water Following the development of an analytical method for determining the concentration of diethyl dixanthogen in water, a series of experiments were undertaken to determine the concentration of saturated diethyl dianthogen solutions. Two methods were used: 1. Solutions of deoxygenated-double distilled water were agitated with liquid (EtX)^ for suitable periods of time. Samples were removed, centrifuged and analysed for (EtX^ content by the hexane extraction method (Section 2.5). 2. Saturation was estimated by iodine titration and measurement of turbidity (Section 2.6). 3.4.1 Measurement of Saturation by Extraction Technique The length of time of agitation of the diethyl dixanthogen^water solution was tested by determining the amount of diethyl dixanthogen that had dissolved after 4, 8, 16 and 30 hours at 22 + 1°C.After 4 hours the concentration of dissolved diethyl dixanthogen had reached 1.0x10 ^ moles liter ^,after ;8 hours 1.26x10 ~* moles liter ^ and after. 16" hours 1.25 moles liter 1 ;(see Figure 18, pH = 6.1).. After 16 hours.agitation' the solutions .were considered to be saturated with (EtX^- Some variation in solubility was observed, as is noted in Figure 18. The difficulty with this procedure was that any minute traces of liquid (EtX)^ that may have remained suspended in the solution would create large anomalous readings. Hence, all samples were analysed in duplicate and when a large variation was found, the analysis was repeated until the precision was i 0.1 x 10 ^;M. 3.4.2 Measurements of Saturation by Turbidimetric Technique This method was found to be more sensitive for evaluating solubility than the dissolution method. The effect of KI, and excess KEtX has been discussed in Section 2.5. Three measurements of (EtX)^ concentration, determined at pH 5.8,6.8 and 7.5 are given in Figure 18. These values compare favorably with the measurements by the agitation technique. The average value for the limit of solubility, using both methods, was observed to be 1.25 x 10 M over a pH range of 2.8 to 8.4. This value compares with 1.3 x 10 ^ M .- as determined by Pomianowski and Leja (18) at pH 6. In all cases during these tests, varying amounts of EtX were present. The effect of changing EtX concentrations did not appear to affect the solubility of (EtX)2- The solubility measurements were limited to pH 9 because above this value the rate of the hydrolysis reaction became significant and interfered with the solubility determination. The low pH region was not completely explored, although a qualitative test showed that (EtX)£ was stable at pH 2 for at least 20 hours. These results expanded on the data of Pomianowski and Leja (18) who obtained a value of 1.3 x 10 M for the solubility of (EtX) 2 at pH 6. A value of 10 was reported by Goldstick (42) who -measured the Tyndall effect of aqueous-organic mixtures containing (EtX^- The present measurements established the saturation concentra• tion of (EtX)^ which was necessary for evaluating the activity of CEtX)„ solutions used in the electrochemical experiments. • Determined by dissolution method rj Determined by turbidimetric method -ri-r**-*"0™* T -6 -1 I ~L 1 x 10 mole liter 1 1 i i it i 1 2 3 4 5 6 PH Figure 18. Effect of pH on the solubility of diethyl dixanthogen. 3.5 Electrochemical Studies on Aqueous Ethyl Xanthate The electrochemical oxidation of xanthate to dixanthogen was written "2(EtX) + 2e~ 2EtX~ Values for the standard reduction potential (E°) of this reaction have been determined by rest potential methods using platinum electrodes (38,39,40), potentiometric titration methods (41,42) and redox methods (37,63), and are summarized in Table 12. Although the majority of E° values obtained using the various methods are in general agreement with each other, some questions regarding the reversibility of the reactions at the electrode surfaces remained unanswered. Specifically, two reasons are apparent for this problem. 1. The reactiyity of the platinum electrode toward aqueous xanthate was not known. This included the possibility that platinum xanthate may be formed on the surfaces of platinum electrodes, causing a significant overpotential. 2. The standard state for dixanthogen has been set at the saturation of dixanthogen in water. The effect oh the rest potential of concentrations of dixanthogen less than saturation has not been investigated. The results reported in this section evaluated the electrochemical rest potentials resulting from varying concentrations of ethyl xanthate, diethyl dixanthogen, oxygen and hydroxyl ion on a platinum electrode, along with qualitative analysis of the surface products by infrared spectro• scopy. A series of polarization experiments were also carried out, along with analysis of the surface products. 3.5.1 Derivation of the Nernst Equation of the Electrochemical Oxidation of Ethyl Xanthate to Diethyl Dixanthogen For the ethyl xanthate-diethyl dixanthogen couple, the following reactions apply: CEtX)2 + 2e --j—*- 2EtX then yCEtX)„ + 2ye" = 2yCEtX~) 2y 2y P e" (EtX ) (EtX)2 -2FE .= 2ye_ = 2y°EtX_) + RT In a^ - - RT In a^^ 2FE° = - [2W°(Etx_) - V(EtX)2] E = E°- || In ^1 2F a (EtX)2 The measured potential is correct if the assumption yi (in Pt) = ug (solution) is valid. The standard state for the activity of ethyl xanthate ion a. = 1.0, has been established, when [EtX ] = 1.0 M. The standard (.EtX ; state for diethyl dixanthogen (a = 1.0) has been chosen at its maximum solubility in water which has been determined experimentally to be 1.27 x 10 5 moles liter"1. 3.5.2 Rest Potential Measurements with Platinum Electrode 3.5.2(a) Ethyl Xanthate Solutions Free of Diethyl Dixanthogen The series of experiments on the electrode rest potentials were carried out in the apparatus described in Section 2.8(a). The stock solutions were prepared from purified KEtX, which was substantially free from (EtX)2, and were added to the cell where potential measurements were recorded as argon purging proceeeded. After the potential had become constant, oxygen was purged through the solutions to generate mixed potentials. The change from argon to oxygen purge was accompanied by a rapid change to anodic potential as can be seen in Figure 19. The concentration of (EtX)2 was measured after argon purging and after oxygen purging of the neutral (pH 7-8) solution. Determinations for (EtX)£ were carried out during argon purging of the solutions, and after oxygen purging. In each case, (EtX-)2'was not detected. Even after runs lasting 24 hours, no appreciable (EtX)2 was found. This result concurs with the results of KEtX decomposition tests reported in the section on alkaline decomposition of KEtX in the presence of platinized platinum where (EtX)2 was also.not detected, The effect of electrode pretreatment was important in determining the value of the maximum cathodic potential that could be reached under +0.190 • ( ^ Argon purge -5 9.8 x 10 M KEtX -4 4.96 x 10 M KEtX +0.120 oo +0.110 If 9.67 x 10 4 M KEtX > +0.100 . w •u g +0.090 A Oxygen purge w +0.080 +0.070 - +0.060 . +0.050 —r~ -r- i I —i— 0 140 1®§ 200 20 40 60 80 100 120 180 Time in minutes Figure 19. Effect of oxygen on the rest potential of ethyl xanthate solutions free from diethyl dixanthogen, pH 7.0-8.0. argon purge. Electrodes that were prepared by anodic treatment (i.e., anodic polarization in dilute KOH followed by washing with distilled water) were initially at more positive potentials than electrodes pre• pared under cathodic conditions (i.e., cathodic polarization in dilute H^SO^ followed by washing with dilute KOH and distilled water). As the solution was purged with argon a larger potential drop was observed for the anodically treated platinum than for the cathodically treated material. After oxygen addition,.the potential changes were larger with the anodically treated platinum. 3.5.2(b) Ethyl Xanthate Solutions Saturated with Diethyl Dixanthogen Solutions of KEtX containing (EtX)^ at pH 7-8 were prepared either by addition of small quantities of KI^ to the stock KEtX solution or by agitating liquid (EtX)^ with distilled water for 16-24 hours and subsequent addition of KEtX to the solution. The effect of the residual I produced by titration method was tested by measuring -4 -4 -3 the rest potential of 10 M KEtX solution containing 10 and 10 M Kl. No significant change in potential was observed, and the results using the two techniques are reported together. For KEtX solutions saturated with (EtX^, the rest potentials rapidly reached a constant value with argon purging. Upon introduction of oxygen, the potential was observed to become more anodic, which agreed with the results of Tolun and Kitchener (39). The magnitude of the potential difference between argon and oxygen purges was found to be a -3 maximum of 18 mV for 10 M KEtX and decreased as the KEtX concentration -3 was decreased below 10 M. This was attributed to the establishment +0.210 +0.200 4 Argon purge Oxygen purge +0.190 5.4 x 10 5 M KEtX, +0.180 J f 1.2 x 10-5 M (EtX). +0.170 j oo w +0.160 o CO +0.150 4 x 10 4 M KEtX, CO 5 > +0.140 1.2 x 10~ M (EtX), cn o 9.32 x 10 M KEtX, > +0.130 -5 A.A—A- 1.3 x 10 M (EtX). rC +0.120 w +0.110 M-A- +0.100 —i r r To" 30- 50 70 90 110 130 150 Time in minutes© Figure 20. Effect of oxygen on the rest potential of ethyl xanthate solutions saturated with diethyl dixanthogen, pH 7.0-8.0. of mixed potentials resulting from the coupling of the cathodic reduction of oxygen with the anodic oxidation of ethyl xanthate. The measurements with a platinum electrode in distilled water at known concentrations of oxygen have been reported by Natarajan and Iwasaki (66) (Figure 21). Using -nitrogen purging for removal of the dissolved oxygen they reported a potential decrease of 60 mV for a ten-fold change in the dissolved oxygen content. The maximum change over the experimental region was reported as 160 mV which corresponded to a change in dissolved oxygen content of 2.8 orders of magnitude. These authors explained that the slope of 60 mV/loglO^] resulted from adsorption of oxygen in the electrode followed by a one electron transfer step such as the one given by the equation (02)ads + e~ + (02~)ads This mechanism is in agreement with Hoare (67) who considered that (02 )ads underwent further reduction to produce (H202)ads. The small potential difference observed for argon and oxygen purged ethyl xanthate solutions which were saturated with diethyl -dixanthogen Indicated that the dissolved oxygen content had a non-stoichiometric Influence, on the rest potential. Unfortunately, because of the complexity of the oxygen half cell, no overall reaction could be assigned to this mixed potential value. Solutions of KEtX saturated with (EtX)2 at pE 8 were left in the cell overnight under oxygen purge. The concentration of (EtX)2 determined the following morning showed that the (EtX)9 content had decreased from -0.56 H -0.54 J -0.52 1 -0.50 -0.48 J -0.46 -0.44 J -0.42 -0.40 +1.0 +2.0 , +3.0 +4.0 +5.0 +6.0 +7.0 +8.0 lo -1 §H02] moles? liter Figure 21. Eh as a function of dissolved oxygen conentration (moles liter xn distilled water, pH 6.8 (after Natarajan and Iwasaki (66)). 1.3 x 10 ^ if to 2.0 to 2.8 x 10 ^ M and was accompanied by a potential rise of 5 mV. Similar tests carried out under argon purge also showed a decrease in (EtX^ content to the same value, and indicated that the reduction of dixanthogen proceeded independently from the amount of oxygen present in the solution. The reduction in diethyl dixanthogen concentration was attributed to the base decomposition reaction which was proposed in the previous section of this thesis (Section 3.2.5). The values obtained for the rest potentials of platinum electrodes in KEtX solutions saturated with (EtX)^ (Figure 22) and purged with argon (deoxygenated) are compared with the results obtained by previous workers (Table 12). Solutions saturated with (EtX)^ were at the standard state a(EtX) = 1' t'ie ecluation became E = E° - JJ In [EtX_] Upon extrapolating log [EtX-] to '[EtX-] = 1.0 M, E = E° = -0.060 V. The linear behavior of the slope (60 mV/logI:EtX_]) is consistent with the one electron reaction for ethyl xanthate oxidation on the assumption that the diethyl dixanthogen concentration is always constant and at unit activity. This assumption was shown to be valid by the experi• ments where the solubility of (EtX)^ in water was determined. The (EtX)2 concentrations were shown to be constant at neutral pH for at least the duration of the present experiments, and over the pH ranges used in this investigation. +8.3 +0.2 • results of present experiments +0.1 • O Majima and Takeda (40) 9'9' A A Tolun and Kitchener (39) -0.0 s''A 0 Stepanov, et al. (38) tM Du Rietz (41) -0.1 Finkelstein (37) } i i —i— i -1 -3 -4 -5 -6 -7 log [KEtXJl Figure 22. Rest potentials of a platinum electrode in solutions saturated with diethyl dixanthogen and under an argon atmosphere. Standard Reduction Potentials of the Xanthate-Dixanthogen Couples (RX) + 2e 2RX E potential (S.H.E.) Reference Method Temp. Methyl Ethyl n-Propyl Iso- n-Butyl Iso- n-Amyl Hexyl propyl Butyl Stepanov Rest Potential Ambient -0.013 -0.037 -0.068 -0.100 -0.132 -0.157 et al (38) Du Rietz (41) Potentiometric Ambient -0.069 -0.095 -0.120 -0.145 -0.140 -0.155 Titration Maj ima & Takeda (40) Rest Potential 25° -0.003 -0,049 -0.092 -0.096 -0.127 -0.127 -0.160 Goldstock (42) Potentiometric 17.5° -0.053 Titration 33° -0.053 Tolun & Rest Potential Ambient -0.081 Kitchener (39) Finkelstein (37) Redox Indicator 25° C-0.08- -0.131 Foss (63) Disulfide Ambient ca.-0.30 Exchange 3.5.2(c) Effect of pH on Rest Potential Unbuffered solutions of KEtX saturated with (EtX)2 at pH 6.0- 7.0 were introduced into the cell (Figure 8) and were purged with argon until a constant potential was attained (approximately 60 minutes). The pH was changed by dropwise addition of 1 M KOH solution to tbe cell with a 1 ml syringe. The pH and rest potential were recorded after reaching a value that remained unchanged for 10 minutes. The results of four such runs given in Figure 23 show a constant potential in the pH region between pH 6 to 8. A pH > 8, a potential decrease of approxi• mately 30 mV/pH was observed which was attributed to two factors: 1. The reaction between hydroxyl ion and (EtX)2 had reached a sufficient rate to decrease the concentration of (EtX)2 from saturation conditions. According to the Nernst equation this change in (EtX)^ concentration would drive the potential to more negative values. 5 2. At pH 9, the I0H~] = 10~ M which is equal to the I(EtX)2]. At higher pH the {OH ] is larger than (EtX) and the negative change in potential could be attributed to the formation of mixed potentials between the xanthate, dixanthogen, H^O and OH . Natarajan and Iwasaki (66) have shown that the rest potential of a platinum electrode in distilled water purged with nitrogen obeyed the relationship Eh = 0.80 - 0.059 pH The potential change of 30 mV/pH observed for the present experiments could not be attributed solely to the effect of the oxygen potential, but to mixed potentials resulting from the combined effects of oxygen +300 +250 H Deoxygenated distilled water +200 w S3 CO +150 to > to •u rH O +100 ^ 5 x 10 M KEtX A ;0-4 M ,KEtX +50 • 1 -3 ^ 10 M KEtX 5 x ,10-3 M KEtX 0.0 T —i— 6.0 7.0 8.0 9.0 10.0 11.0 12.0 13.0 pH Figure 23. Effect of .pH on the rest potential of aqueous ethyl xanthate saturated with diethyl dixanthogen. and diethyl dixanthogen on the rest potentials. 3.5.2(d) Effect of Diethyl Dixanthogen on the Rest Potential The rest potentials of KEtX solutions containing concentrations of (EtX), less than saturation were also studied. The results plotted in Figure 24 were obtained by addition of small quantities of KI^ solution to a KEtX stock solution at pH 7.0. The (EtX)^ concentrations were measured by comparing the results of the hexane extraction method (Section 2.5) with the quantities of normal iodine solution added. The results obtained from these experiments were combined with those obtained from all the previous tests on the rest potentials of argon purged ethyl xanthate-diethyl dixanthogen solutions are plotted in Figure 25. The calculations for the activity of (EtX)^ were made on the basis that unit activity (Standard State = 1) was reached at maximum solubility of 1.27 x 10 ^ M. The Nernst equation obtained for the electrochemical oxidation of .xanthate was written previously as E « E° -M m aRtx F 1/2 U(a ); (EtX)2 1/2 a a The values of log EtX~/^ (Et.x) ^ were plotted against the observed rest potential and this resulted in two separate plots. Those experi• mental values obtained at pE 6 to 8 (plot (a)) have a slope of 58 mV which is in agreement with the results of previous workers reported 1/2 in Figure 22. The intercept of this plot (log a„ -/(a,Ei.v> ) ' =0) t,ZX ^JitA.^2 was found to yield E°= -0.063 + 0.01 V which is also in agreement with +0.140 Saturated Solution (1.27 x 10 5 M) +0.130 • -O- O — -o- +0.120 V w 6 w 9.8 x 10 M (EtX) +0.110 • CO / • 7.9 x 10 M (EtX), C>O 4J 5 x 10 M KEtX Q 6.8 x 10-6 M (EtX), O +0.100 . > [(EtX) ]= 0.5 x — A -6 10 M • 4.4 x 10 M (EtX), +0.090 -i— - i —r- —r 20 30 40 60 0 10 50 Time in minutes Figure 24. Effect of diethyl dixanthogen on the rest potential of a platinum electrode unde argon atmosphere and constant ethyl xanthate concentration, pH 6.9-7.7. 4-0.1230 +0.1^0 3 5 x 10 5 M KEtX, pH 8.2-11.5 +o.iao 5 x 10 4 M KEtX, pH 6.0-7.7 +0.1Q0 +0.1^0 10 4 M - 5 x 10 5 M KEtX, pH 9.0-9.3 +0.1(50 w 10 M KEtX, pH 7:0-8.0 +0.1^0 CO +o.i(3;o CO rH +0.120 o O > +0.110 +0.100 V +0.090 - +0.080 +0.070 +0.060 +0.050 -4.4 -4.0 -3.6 -3.2 -2.8 -2.4 -2.0 EtX log a(EtX)^/2 Figure 25. Correlation of the rest potentials of ethyl xanthate and diethyl dixanthogen solutions w,ith the Nernst equation. the results obtained by previous workers (see Table 12). "The values which deviated from the straight, line- plot (Figure 25) were obtained at. pH 9 or at low EtX and (EtX) concentrations. Under these conditions, "dissolved oxygen, OH or minor contaminants could have, interfered with the attainment of equilibrium. Thus, no interpretation can be given for those results. 3.6 Polarization of Platinum Electrodes in Ethyl Xanthate- Diethyl Dixanthogen Solutions The- three electrode apparatus used in the polarization studies was constructed according to the design given by Tolun and Kitchener (39) and is described in Section 2.8(b). The solutions were prepared in the side vessel and were purged with argon before being added to the cell, which was also under inert atmosphere. The oxygen content could not be tested since any analytical methods which depended on the quantitative reduction of oxygen for its determination would be poisoned by the presence of ethyl xanthate. However from data supplied by Natarajan and Iwasaki (66),the^concentration of oxygen in —6 —1 distilled water was reduced to less than 2 x 10 moles liter (from -4 -1 10 mole liter ) after purging with nitrogen for one hour. Since the conditions used in the present experiments were similar to those used by Natarajan and Iwasaki (66), the supply of oxygen to the electrode was considered to be reduced to negligible quantities in comparison with the concentrations of the other reagents. The three polarization curves given in Figure 26 were typical of similar curves obtained under a wide range of experimental conditions. -3 Descriptions of these curves, obtained for 10 M KEtX in 1 M KC1 electrolyte were detailed as follows: 1. The base line (Curve (1)) was established with electrolyte only and with 0.1 M ammonium acetate and 0.1 M sodium borate buffers and electrolyte. The current was noted to undergo through a hysteresis at the start of each experiment, i.e., in proceeding from anodic to cathodic sweeps, the current differed by 10 micro amperes at equivalent potentials. This effect was insignificant compared to the magnitude of the current when KEtX was present and had no effect on the subsequent polarization curves. 2. The application of anodic overpotential onto the platinum -3 electrode in argon purged 10 M KEtX resulted in the electrochemical oxidation of KEtX to (EtX)„. At an applied potential E„ „ _ = +0.42 V, the electrode was removed and an infrared spectrum of the region 1400 cm ^ to 700 cm ^ was obtained using the ATR technique. The spectrum confirmed the presence of (EtX)2 as can be seen in Figure 29. The circuit was opened at +0.55 V and the potential value decayed to the rest potential value +0.123 V (S.H.E.) in twenty minutes. The current observed during cathodic polarization resulted from reduction of the adsorbed "(EtX) to EtX and proceeded until...r hydrogen was evolved. The anodic and cathodic curves were identical for solutions buffered with ammonium acetate, sodium borate and without buffer. In the latter case, the pH of the solution, initially at 8.5 dropped to 5.5 during anodic polarization when oxygen began to evolve ' and then rose to 9.5 during the cathodic polarization sweep when hydrogen was evolved. 3. Mixed polarization curves were obtained from KEtX. solutions -3 purged with oxygen. The rest potential in 10 M KEtX was +0.136 V Figure 26. Polarization curves for a platinum electrode in 10 M ethyl xanthate in 1 M KC1 electrolyte. (S.H.E.) or 13 mV more anodic than the same solution under argon purge. Similar results are given in Figure 20. The anodic polarization curve followed the oxidation of air free KEtX closely until the potential exceeded +0.3 V. The cathodic polarization sweep illustrated the reduction of oxygen along with (EtX^. The rate of oxygen reduction was sufficiently rapid that the reduction of (EtX)2 could not be observed. In addition to the observations presented in Figure 26, several other experimental parameters were investigated which did not result in any marked effect on the overall shape of the curves. These results can be summarized as follows: -4 -3 1. Polarization curves for 10 M and 5 x 10 M KEtX were obtained with no resulting change in their shape. The rest potential for 10-4 M KEtX was +0.187 V (S.H.E.) and for 5 x 10~3 M KEtX was +0.103 V (S.H.E.). 2. Electrolyte solutions of 0.1 M and 1.0 M KC1 had no effect on the shape of the polarization curves, but slightly lower currents were observed for 0.1 M KC1 solutions. 3. The results obtained from shiny and platinized platinum electrodes were likewise similar. Current densities obtained for platinized platinum were larger than those observed for shiny platinum, due to the increased surface area of the platinized electrode. The results presented in Figure 26 were replotted in Figure 27 to show the Tafel behavior of the anodic and cathodic polarization sweeps. In the place of the ammeter -used to obtain the results presented in Figure 26, a special potentiostat was employed which improved the current measuring accuracy of the apparatus. Anodic and cathodic polarization potentials were applied in steps of 10 mV and the steady state current -3 was measured. The plot for 10 M KEtX solution containing 1 M KC1 electrolyte purged with nitrogen is given in Figure 27. From the initial anodic polarization, a linear Tafel region was observed, xvith a slope of approximately 60 mV/log i. From the Tafel equation r\ = a + b log i b = 2.303 ^| r =0,059 „ 0.060 L'U Suchaa value for n was expected for a one electron reaction. During cathodic polarization, a linear Tafel region was not observed. The slope of 80 mV/log i was obtained from a line drawn from the rest potential to a tangent of the cathodic curve (see Figure 27(b)). The non-linearly may be explained by considering that the supply of uncharged dixanthogen molecules was limited to their rate of diffusion to the electrode, and thus a region controlled by activation over- potential would be difficult to ,obtain. The exchange current density (n = 0) was 4 x 10 1 amp/geometric cm . Potter (68), in a discussion of the significance of exchange current values pointed out that i for various electrode processes cover o r —18 —2 2 a range from about 10 to 10 amps/geometrical cm of the electrode. -7 2 If the value of iQ is low (e.g., 10 amps/cm ) it would be unlikely except in systems of exceptional purity that the reversible potential +240 1 +220 / +200 - +180 " J1, +160 +140 'O^ (b) +120 +100 ' +80 . +60 • +40 o I —I— -4 -3 ~^2 -4 -3 -2 log 1 amperes/geometrical cm -3 gure 27. Polarization diagram for a platinum electrode in 10 M ethyl xanthate, 1 M KC1 electrolyte (a) polarization over a one volt range, (b) magnification of (a) to show Tafel behavior. of the electrode process could be attained practially before some alternative process capable of sustaining a current density greater than i assumed the control of the potential. For the reversible o _3 hydrogen electrode at platinum cathodes, i is about 10 amp/geometrical 2 cm and impurities that would interfere with the reversible hydrogen reaction would be present in easily detected quantities. The exchange current betweeen ethyl xanthate and diethyl dixanthogen was found to be 4 x 10 amps/geometrical cm . From the previous examples, this current should be sufficiently large to exclude the effect of any minor impurities in the solution, with the result that the rest potential measurements would accurately reflect the electro• chemical equilibrium between ethyl xanthate and diethyl dixanthogen. 3.7 Infrared Investigations of Platinum Electrode Surface The reference infrared spectra for KEtX (Figure 28(a)) was obtained using the KBr pellet technique, the spectrum of diethyl dixanthogen was obtained from a capillary film (Figure 28(b)) and the spectrum of aqueous ethyl xanthate ion (Figure 28(c)) was obtained using a special cell constructed from polyethylene film. From observations of the spectra, the difference between xanthate dixanthogen ' and platinum xanthate can easily be distinguished on the basis of their characteristic peaks in the region 1300 cm 1 to 700 cm \ The assignment of molecular vibrations in the xanthate molecule to the infrared absorption bands has been carried out by Little, Poling and Leja (69). The strong band for KEtX at 1050cm"1 and_at 1040cm"1 for (EtX)2 and (EtX ) has been attributed to the C=S stretching mode. The multipeaked band in the region 1080 to 1180 cm 1 for KEtX was designated as resulting from stretching vibrations of the C-O-C linkage. Similar C-O-C bands are observed at 1260 and 1240 cm 1 for (EtX)2 and at 1160 cm 1 for aqueous ethyl xanthate ions. Spectra of the material deposited on the surface of platinum electrodes were obtained using the multiple ATR technique described in Section 2.7(c). A summary of thednfrared results presented in Figure 29 follows: 1. The platinum foil, whether shiny or platinized platinum yielded a flat base line as can be seen in Figure 29(a). The three small peaks observed at 1040, 925 and 770 cm 1 resulted from the light transmission characteristics of the KRS-5 crystal and were not due to absorption bands from the platinum. -3 2. After the platinum foil had been anodically polarized in 10 M KEtX at +0.43 Y (S.H.E.) for one minute.(see Figure 26) the electrode was withdrawn from the solution and the spectrum, characteristic of (EtX)2 was obtained (Spectrum (b)). Ether washing resulted in removal of the adsorbed layer from the electrode as was observed in Spectrum (c). 3. Diethyl dixanthogen, which had been deposited on the electrode by anodic polarization was then subjected to cathodic polarization for one minute at -0.6 V (S.H.E.). The resultant spectrum (d) illustrates the reduction of all of the (EtX)^ deposited by anodic polarization. _3 4. Platinized platinum foil was also immersed in 10 M xanthate solution without polarization. Under basic conditions, pH 8 to 11, no (EtX)„ was observed, but pH 6.0, with 30 minutes exposure time in the 50 100 o •H CO CO •H e (EtX) (liquid) CO c H 50 100 C (c) CU O S-i CU CM KEtX (aqueous) 50 100 Pt(EfX) (solid) 1400 1300 1200 1100 1000 900 800 700 wavenumber cm 1 Figure 28. Infrared spectra of ethyl xanthate compounds, .(a) KEtX, solid, KBr pellet; (b) diethyl dixanthogen, capillary, film; (c) ethyl xanthate ions, aqueous solution; (d) platinum xamttehate, KBr pellet. (a) Pt foil after cleaning in aqua regis and cathodic polarization for 15 mins. IX 50 100 (b) Pt foil after anodic polarization at +0.2 V for 1 min. in 10"3 M KEtX 50 100 OS) (b) after ether wash (d) (b) after cathodic polariza• tion 50 100 Pt foil after 30 mins. in 10-3 M KEtX (not^eolarized) pH 6.0 i 1400 1300 1200 1100 1000 900 800 -1 Wavenumber cm Figure 29. Infrared (ATR) spectra of platinum foil after varying treatments. solution, weak bands attributed to the formation of (EtX)^ were observed (Spectrum (e)). The formation of platinum oxides and hydroxides at the platinum surface has been suggested as contributing to the non-reversibility of the Pt-02-H20 system (67). By analogy, platinum xanthates in the Pt- KEtX-IL^O system could also be found on the electrode, and the potential measurements would reflect a non-equilibrium process. Platinum xanthate was prepared by a method described by Watt and McCormick (70) and its infrared spectrum was obtained using the KBr pellet technique (Figure 29(d)). The absorption bands corresponded with those obtained by Watt and McCormick (70) (Table 12). The results presented in Figure 29 show that platinum xanthate is not observed under the conditions investigated. The amount of platinum xanthate required to block the electrode processes between KEtX and (EtX)2 can be assumed to sufficient to cover the active sites, which is less than a monolayer. The coverage of platinum xanthate on the electrode for the geometric surface area can be calculated assuming °2 that the area of a xanthate molecule is 28 A . The electrode would contain 2 16 ^2 electrode area (4 x 10 cm x 10 A x 321 gm" 2 cm mole 28A°2 6.02 x 1023 molecules - • mole 6 = 9 x 10T gm Pt(EtXl2 Because of surface roughness, the actual area of the electrode surface is at least twice the geometric area for a shiny platinum electrode (68). TABLE 13 Vibrational Band Assignments for the Infrared Spectrum of Platinum Ethyl Xanthate (after Watt and McCormick (70)) Pt(EtX)2 Assignment Description 1389 w 6 CCH2) deformation 1371 m 6CCII3I deformation 1320 sh,w 1285 -vs . vCC-Oi stretching 1144 w OJCCF^) wagging 1114 s v(C=S) stretching 1058 vw v(CH2) twisting 1017 vs P (CH3) rocking 1004 m 853 m v(C-C) stretching vw - very weak w - weak m - medium sh - sharp vs - very sharp Platinizing increased this area by a factor of 10 to 100 times. Under these conditions, monolayer coverages of platinum xanthate could be detected by infrared spectroscopy. The absence of any infrared bands that may be related to platinum xanthate indicates that if it were present, the coverage would be less than a monolayer. 3.8 Interpretation of the Electrochemical Results The electrochemistry of xanthate solutions saturated with dixanthogen has been widely studied. Some of the results of these various studies on ethyl xanthate are presented in Figure 2 2 - where the standard potential was determined by extrapolating the measured E values at various xanthate concentrations to [EtX ] = 1.0 M. Similar investigations with other xanthate homologues have also been reported and are given in Table 12. The present results obtained for ethyl xanthate and diethyl dixanthogen were undertaken to add information about the effect of oxygen on the electrode potential measured in ethyl xanthate solutions, and of the effect of diethyl dixanthogen concentrations less than saturation on the electrode potential. The results of this work can be summarized as follows: 1. The electrode potentials observed for xanthate solutions saturated with oxygen are considered to arise from two sources: Ca) either they represent a non-equilibrium reaction at the electrode'surface, which results from immeasurably slow reactions between ethyl xanthate and oxygen or; (b) the oxygen and ethyl xanthate do not react at all, In which case the observed potential is a mixed potential dependent upon the relative exchange currents of various species contribut• ing to the electrode process. The evidence obtained in support of this conclusion can be briefly stated as follows: (i) the relative concentrations of ethyl xanthate and diethyl dixanthogen were not influenced by either the pH or oxygen content of the solution in the pH region 6 to 9. (ii) the difference in potential between argon and oxygen purged solutions of ethyl xanthate saturated with diethyl dixanthogen was a maximum of 18 mV., although argon purging over a period of 60 minutes would have reduced the oxygen content to 2 x 10 ^ moles liter \ or 2% of saturation (66). This indicates that oxygen has a non-stoichiometric effect on the electrode reaction. (iii) The non-linear relationship between the potential and the oxygen content indicates mixed potentials were observed which could not be related to any overall electrode reaction. (iv) The value obtained for the exchange current of the ethyl xanthate-diethyl dixanthogen couple on a platinum electrode (4 x -4 2 10 amps/geometrical cm ) was sufficiently larger than the value -9 obtained for the oxygen-hydroxyl couple (viz 1.3 x 10 amps/ 2 geometrical cm (67)) which indicates that the ethyl xanthate- diethyl dixanthogen couple would control the measured electrode potentials. 2. Within the limitation of these experiments, the oxidation of ethyl xanthate to diethyl dixanthogen has been shown to be a reversible reaction in an electrochemical sense. The evidence obtained in support of this conclusion can be briefly stated as follows: ' Ci) The rest potentials of varying concentrations of ethyl xanthate in solutions saturated with diethyl dixanthogen correlate with the Nernst equation in that a linear plot with a slope of 60 mV/log [EtX ] was observed. (ii) The rest potentials of varying concentrations of diethyl dixanthogen at constant ethyl xanthate concentration was tested with the Nernst equation and a linear plot with a slope 1/2 of 59 mV/log a^ -/(a/r,tV. ) was observed. The extrapolated EtA (btX;2 value for E° = -0.063 V correlated well with the results of other workers. (iii) The lack of evidence of platinum xanthate added support for the absence of chemical reaction between the platinum electrode and the xanthate solution. The formation of platinum xanthate on the electrode surface as an intermediate step during the oxidation of xanthate or the reduction of dixanthogen is expected to be quite possible. The present results have shown that this reaction does not interfere with the overall transfer of electrons between xanthate and dixanthogen. Similar observations have been made by Goldstick (42) on the reactions between xanthate and the silver electrode. Although silver xanthate was formed on the electrode, the surface film was found to be sufficiently conducting so that the oxidation potential of xanthate to dixanthogen could be measured. From these measurements, Goldstick (42) concluded that the reversible electrode potential of the xanthate-dixanthogen couple could be measured with silver electrodes. The adsorption of xanthate on copper surfaces resulted in the rapid growth of a cuprous xanthate film which insulated the metal surface from the xanthate solution. The cuprous xanthate film could not be removed by application of cathodic potential as:was found for platinum and silyer. Copper electrodes were concluded to behave irreversibly and were unsuitable for the measurement of xanthate-dixanthogen electrode potentials. CHAPTER 4 SUMMARY AND CONCLUSIONS The evaluation of ethyl xanthate decomposition, diethyl dixanthogen decomposition and the electrochemical behavior of the ethyl xanthate- diethyl dixanthogen couple has demonstrated that each of these systems undergo a different set of reactions in neutral to mildly alkaline conditions. The main results relating to each system can be briefly described as follows: (1) The decomposition of KEtX in neutral to mildly alkaline aqueous solution was shown to proceed by a hydrolytic mechanism given by the equation S II _ C H H + CS + H C2H5-0-C-S + H20 y 2 5° 2 ° The kinetic studies supported the postulated mechanism by providing the following information (i) The quantity of CS2 evolved was in a 1:1 molar ratio for the quantity of KEtX decomposed (Table 6). (ii) The rate of CS2 evolution was equal to the rate of KEtX decomposition,(Figure 14). (iii) The reaction was first order in KEtX, indicating that only one molecule of KEtX was involved in the rate controlling steps of the decomposition process. The first order rate constant v was also measured (k^ = 7.6 + 1.0 x 10 4 hr "*") (Section 3.1.1). (iv) The decomposition rate was independent of pH over the range tested (pH 7-12). The increase in pH observed during the course of a decomposition run was also related to the postulated reaction (Figure 13). (v) (EtX)^ was not detected as one of the reaction products. Auxilliary experiments were also carried out which explored the possibility that (EtX^ could be produced by oxidation of aqueous KEtX with oxygen in the presence of catalytic agents. Of three possible agents tested, ferrous sulfate had a slight retarding effect 6n the decomposition rate of KEtX, while platinized platinum and methylene blue accelerated the rate by 2 and 10 times respectively. (EtX)^ was not detected during the course of the runs (Table 7). (2) The reaction between diethyl dixanthogen and hydroxyl ion was found to proceed in two steps in which the first step was a nucleophilic attack of the sulfur-sulfur bond of (EtX)2 by OH which resulted in the formation of ethyl xanthate and an unstable sulfenic acid intermediate. In the second step of the reaction; the disproportion- ation of ethyl xanthate sulfenic acid resulted in the formation of CS2 and other products. The reactions are given by the following equations S S OH II II C2H50-C-S-0H + CH 0-C-S S W C2H50H + CS + OH CoHr0-C-S-0H 2 The observations which supported the proposed mechanism can be briefly stated as follows: (i) Material balance studied showed that one mole of EtX was generated for each mole of (EtX)2 decomposed (Table 8). (ii) Kinetic studies confirmed the rate to be first order with respect to each of the concentrations of (EtX)2 and. OH (Section 3.3). _d[(EtX) ] dt = k2 [(EtX)2][0H ] This rate expression is consistent with that required for the postulated bimolecular displacement mechanism. The experimental second order rate constant was measured (^2 = 14.4 + 0.3 liter mole "'"min ^) . (3) The electrochemical experiments indicated that aqueous solutions of ethyl xanthate and diethyl dixanthogen came to a reversible electrochemical equilibrium at a platinum surface. It was also shown that ethyl xanthate was readily oxidized to dixanthogen by the applica• tion of a small anodic overpotential, but not by dissolved oxygen. The experimental evidence supporting the conclusion that ethyl xanthate and diethyl dixanthogen behave reversibly can be briefly given as follows: (i) Rest potential measurements were obtained for varying KEtX concentration at constant saturated aqueous (EtX)2 concentra• tions (Figure 23). The results fitted the Nernst equation for a one electron oxidation reaction. (ii) Rest potential measurements were also obtained for varying (EtX)„ concentration at constant KEtX concentration, and these also indicated a one electron transfer reaction (Figure 25)• (iii) The Tafel slope resulting from the initial stages of anodic polarization of the platinum electrode indicated a one electron oxidation reaction. Infrared spectroscopic examination of the platinum surface confirmed (EtX^ as the only detectable reaction product • (Figures 27 and 29). (iv) No changes in the (EtX^ content of KEtX solutions were noted after purging with oxygen. This evidence supported the conclusion that dissolved oxygen was not effective in the oxidation of KEtX to (EtX) 2'.(Section 3.5.2(b)). Numerous studies on the decomposition of KEtX under acidic conditions have shown that the ethyl xanthate anion is protonated, forming xanthic acid which decomposes into ethyl alcohol and carbon disulfide according to the following equation S S k H - K I' + 2 + CoHr0-C-S + Ho0 -^-»- C„Hc0-C-SH + H C„Hc0H + CS0 + H 2 5 2 •< 2 5 2 5 2 A review of the published work on this reaction is presented in Appendix 2. Determinations of the dissociation constant for xanthic acid -2 -1 formation generally agree on a value of 2.8 ± 0.3 x 10 moles liter The second order rate constant for xanthic acid decomposition has been reported as k^ = 230 Z 20 liter mole "'"min "*". The decomposition of aqueous ethyl xanthate and diethyl dixanthogen has been summarized in Figure 30 where the half-life time of the decomposi• tion is plotted against pH. The figure illustrates the instability of xanthate under acidic conditions and the instability of dixanthogen under basic conditions. solution at 22°C. The present results also provide some new information to the mech• anisms of depression of some minerals, notably pyrite. KCN, ^2820^ and CaO are depressants for pyrite and a partial explanation for this result is the reactivity of these reagents toward the adsorbed dixanthogen. The effect of this reaction is only a partial answer to the problem of explaining depression since depressants also affect the zeta potential and adsorption of oxygen on the mineral surface. APPENDIX I Curve Fitting and Error Analysis Curve fitting was obtained by. first transposing the data to fit a set of linear coordinates and then applying linear regression analysis to those results. The calculations were carried out on a Hewlett Packard Model 9100 A calculator which contained a program that yielded values for the intercept a, the slope b, and the correlation coefficient r. The calculated values of |r| are given in Table 1. These are compared with tabulated values obtained from Neville and Kennedy (74) for a significance level of 95%. A genuine correlation at this level is shown to exist when the calculated value of |r| exceeds the tabulated value. The equation for a straight line is given as y =. a + bx where b is the slope of the line The error in slope is given by b +1 t R where t R is the confidence limit. The t values for a 95% confidence limit were obtained from a table in Neville and Kennedy (74) which related the number of experimental points to the probability value. R is defined by £ (y^y)2-mE (y^-y) Cx,.-x) where x and y are the mean values, x_^ and y_^ are the experimental values and n is the number of trials. Values for t and R are also \ given in Table;1. The results obtained from studies on the kinetics of xanthate and dixanthogen decomposition reactions and those obtained from the electrochemical experiments were treated with the same methods, and the levels of significance for the results are given with those results. Correlation Coefficient and Error in Rate Constants for the Oxidation of KEtX Solutions Description of Run |r| |r| theoretical t R x 10 First Order Rate Constant ± t R Methylene Blue 1. oxygen free 0.95 0.81 2.132 .23 2.3 ± 0.05 x 10 4hr 1 -1 2. oxygen saturated 0.92 0.93 2.015 .50 400 ± 100 hr Platinized Platinum 1. oxygen free 0.99 0.95 2.92 .25 9:7 ± 0.7 x 10 4hr 1 2. oxygen saturated 0.997 0.75 2.015 .48 15.7 ± 1.0 x 10-4hr-1 Ferrous Sulfate 1. oxygen free 0.98 0.86 2.132 0.23 3.6 ± 0.5 x 10 4hr 1 2. oxygen saturated 0.975 0.93 2.015 0.19 4.2 ± 0.5 x 10~4hr-1 APPENDIX II Decomposition of Ethyl Xanthate In Acid Solution Alkyl xanthates of Na+ or K+ ionize readily and react with water under acidic conditions to form xanthic acid. S ' S II II R-O-C-S + HO y- R-O-C-SH + OH Aqueous xanthic acid is highly unstable and decomposes rapidly into carbon disulfide and alcohol S II R-O-C-SH y ROH + CS The reaction rate constant has been shown by Iwasaki and Cooke (75,76) to be second order,and the kinetics of the decomposition of xanthic acid have been described by the following equation ^=g- = k (X")(H+) Guadin (72), in his studies on the flotation of pyrite with ethyl xanthate reported that the decomposition of the xanthate at low pH had to be taken into account to explain the high xanthate consumption during his tests. A number of studies have been carried out to determine the ionization constant and the decomposition rate of ethyl xanthate. This section is a summary of the works published on the decomposition of ethyl xanthate in acid solution. The first investigators of the acid xanthate reaction concluded that the decomposition resulted from combination of the xanthate ion with hydrogen ion to form xanthic acid, which decomposed into carbon disulfide and alcohol, von Halban and Hecht (77) mixed ethyl xanthate solutions with hydrochloric acid, stopped the reaction at suitable times by the addition of sodium carbonate and then determined the remaining ethyl xanthate by iodine titration. The ionization constant was calculated for the dissociation of ethyl xanthic acid (0.030) and the second order rate constant (16.0 liter mole "*"min ^) for the decomposition reaction was determined (see Table 14). King, et.al. (78,79,80) measured the decomposition rate by following the increase in vapor pressure of the reaction due to the formation of carbon disulfide and other products. The rates for a number of xanthate homologues were studied by Komylev (81) who correlated the half-life of the reactions with. pE, temperature and chain length of the particular xanthate. Cook and Nixon (82), using a pE technique reported a value for the equilibrium constant considerably lower than those reported previously. In a later paper they admitted that their method was subject to a considerable error because of the non-stoichiometric: consumption of hydrogen ions in the reaction. All of the more recent studies (since 1958) have been made employing the ultraviolet spectrophotometric method for measuring the xanthate ion concentration. Iwasaki and Cooke (75,76) using this method evaluated the rate constant for ethyl xanthate to be 4.3 min 1 and the ionization constant as 0.020. Klein, et.al. (83), also using spectro• scopic methods postulated a decomposition mechanism based on an activated complex where the hydrogen ion is in.the vicinity of the oxygen atom on the R-O-C chain. Majima (84), in a detailed investigation into this system, recognized that the contribution of ethyl xanthic acid to the spectral absorbance of the ethyl xanthate ion was very small, a consideration overlooked by Iwasaki and Cooke. Majima also evaluated the ionization constants and the decomposition rate constants for the homologous series and found that both constants decreased as the chain length of the hydrocarbon was increased. His investigation also showed that ionic strength had a small effect on the decomposition rate. A similar study by Zahradnick (85) at 0.5°C showed no definite trend in both constants in passing from methyl to amyl xanthate, but then showed a sharp decrease in decomposition rate with the longer chain derivatives. Pomianowski and Leja (86) also evaluated the decomposition rate of ethyl xanthate as a part of a study on the formation of complexes between ethyl xanthate and alkyl trimethyl ammonium bromide (C^TAB). They pointed out that the decomposition rate a neutral pH values was affected by small quantities of metallic ions which could act as oxygen carriers or catalysts. In a recent study, Hopstock (87) proposes that a rapid equilibrium is reached between the xanthate ion and xanthic acid. A small concentra• tion of a zwitterionic form of xanthic acid is present as a result of bonding of the hydrogen atom to one of the electrons on the oxygen atom. The zwitterionic form is unstable and decomposes rapidly into alcohol and carbon disulfide. The results of these workers have been given in Table 14 and show a reasonable consistency between the ethyl xanthate decomposition rate constants obtained by all workers except those of the very early investigations. The average value of the decomposition rate used in the calculation in Figure 30 was 230 + 20 liter mole 1 min 1 obtained from an average of the results of the more recent workers, notably Majima and Eopstock. TABLE 15 Dissociation Constants and Decomposition Rate Constants for Ethyl Xanthic Acid Method Reference Temp.°C Dissociation Decomp. Rate _v Constant Constant -1 iodometric von Halban 0.030 16.0 liter mole titration & Hecht (77) min-1 vapor King et.al. 25 0.031 270 liter mole -1 pressure (78,79,80) min~l -1 Spectrophoto• Pomianowski & 23 165 liter mole metry Leja (86) min-1 pH determina• Cook & Nixon 0.007 tion (82) -1 Spectrophoto• Iwasaki & 23.5 0.020 4.3 min metry Cooke (75) Spectrophoto• Cooke & 21 0.0239 metry Iwasaki (76) -1 Spectrophoto- Majima (84) 25 0.029 226 liter mole metric min~l Spectrophoto- Toernell (88) 25 0.046 metric -1 Spectrophoto- Homylev (81) 25 185 liter mole metric min~l -1 Spectrophoto- Klein (83) 10 0.0255 29.0 liter mole metric min"! -1 0.0230 10.5 liter mole -.cm"-'- -1 Spectrophoto- Hopstock (87) 23.5 0.0263 240 liter mole metric min~l REFERENCES 1. Handbook of Chemistry and Physics, 50th Edition, The Chemical Rubber Company, 1970. 2. A. Cambron, Can. J. 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