REACTIONS OF IRIDIUM AND WITH

ETHTLENEDIAMINETETRAACETIC

DISSERTATION

Presented in Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the Graduate School of The Ohio State University

By HAROLD DALE McBRIDE, B. A., M. S.

The Ohio State University 1958

Approved by

1/Vg-Ar-vq* *" Adviser Department of Chemistry ACKNOWLEDGMENTS

The author expresses his sincere gratitude to Dr. William M. MacNevin for his enthusiastic encouragement and wise counsel during the course of this study.

Acknowledgment is also made to the Department of Chemistry of the Ohio State University for financial aid in the form of assistantships and to the National Science Foundation for a Fellowship from June, 1957» to the completion of this work.

ii DEDICATION

Dedicated in memory of iry son, Gary.

iii TABLE OF CONTENTS

PAGE

STATEMENT OF THE PROBLEM ...... 1

INTRODUCTION...... 2

Theory of Coordination Compounds and Conplex Ions ...... 2

Chemistry of EDTA...... 7

EDTA, General Information and Properties ...... 7

Preparation of EDTA ...... 10

EDTA as a Chelating A g e n t ...... 12

Preparation of EDTA Complexes...... 17

EDTA in Oxidation-Reduction Reactions ...... 21

Analytical Uses of EDTA ...... 23

Chemistry of ...... 26

Oxidation Potentials of Rhodium and I r idium ...... 26

Separations and Analyses of Metals .... 28

BACKGROUND OF THE P R O B L E M ...... 30

EXPERIMENTAL...... 31+

Reactions of Iridium(IV) with E D T A ...... 31+

Job’s Method Applied to Iridium(IV) and EDTA ...... 31+

Other Preliminary Studies . 35

Attempts to Prepare Iridium(IV)-EDTA Complex ...... 38

Titration of Iridium(IV)-EDTA Product . 1+3

Chromatographic Separation of the Product ...... 1+5

Reactions of Iridium(IV) with Na2H2Y ...... 1+6

iv V

PAGE

Reduction of Iridium(lV) with EDTA ...... U 8

Postulated Mechanism...... 56

Additional Evidence for Mechanism ...... $8

D i s c u s s i o n ...... 59

Evidence for Rhodium(III)-EDTA Complex . 62

CONCLUSIONS ...... 66

BIBLIOGRAPHY...... 67

AUTOBIOGRAPHY ...... 70 LIST OF TABLES

TABLE PAGE

I, Directional Characteristics of Covalent Bonds .... 6

II. Solubilities (/lOO ml.) of EDTA and Sodium Salts

in Water and pH of Aqueous ...... 9

III. Logarithms of Formation Constants of EDTA Complexes . . 1 $

IV. Atomic Ratios of Elements Based on Elemental Analysis

of Product ...... 1*2

V. Atomic Ratios of the Elements Based on Elemental

Analysis of Product Prepared from Solutions at pH 3-5 . 1*7

VI. Molar Ratio of Iridium(lV) : EDTA in the Presence

of Excess Iridium(IV)...... 51*

vi LIST OF FIGURES

FIGURE PAGE

1. Job's Method, for EDTA and Iridium(IV)...... 36

2. Absorption Curves for I^IrCl^ and H 2lrCl6 Treated

with E D T A ...... 39

3* Absorption Curve for Product Isolated from Reaction

of H 2lrCl6 and ED T A ...... Uli

1*. Absorption Curves for H2lrCl6 and H3lrCl6 Formed by

Electrolytic Reduction of H^IrClg ...... 1*9

5. Absorption Curves for IrCl£s Formed by Reduction of

^IrClg with EDTA and Hydroxylamine Hydrochloride . . 5l

6. Dependence of Molar Ratio of Reactants on Moles

Iridium(IV) : EDTA Reacting . 55

7. Absorption Curves for Rhodium(lIl)-EDTA Complex . . . 63

vii STATEMENT OF THE PROBLEM

In 19$k Kriege'*' reported spectrophotometrie, complexometric, and

■*■0. H. Kriege, Ph.D. dissertation, Ohio State University (195>U).

hydrolytic evidence for the formation of a complex of tetravalent irid­ ium with ethylenediaminetetraacetic acid in chloride . He found the complex to have the unusual composition of two moles of iridium to one mole of EDTA if EDTA was in large excess in the reaction mixture.

There was also evidence for some 3sl complex when EDTA was not in large excess. Further work on these complexes, such as solution studies on the formulas of the complexes, their preparation, and determination of their stability constants, is indicated.

Kriege found no evidence for a complex of trivalent rhodium with

EDTA in chloride solution. If the chloride complex is too stable for the EDTA complex to form in chloride solution, it can be expected to form in the absence of chloride. It is proposed to prepare the EDTA complex by starting with rhodium in the form of its sulfate, since sul­ fate forms less stable complexes than chloride.

1 INTRODUCTION

Theory of Coordination Compounds and Complex Ions

In general, a coordination compound may be defined as the addition product of the combination of an electron donor and an electron accep­ tor, each capable of independent existence. The electron acceptor or

the donor may be an atom, an ion, or a molecule. Compounds resulting

from the combination are referred to as complex confounds or coordina­

tion confounds, and the resulting ions are complex ions. However, the

above definition, which is based on the formation of a coordinate bond

according to the theory of Sidgwick2 and Lowry,^ includes only a part

2N. V. Sidgwick, J. Chem. Soc., 123, 72$ (1923).

•^T. M. Lowry, J. Soc. Chem. Ind., U2, 316 (1923).

of the types generally included as complexes. Although chemists do not

agree on a simple definition of complexes, the addition compounds includ­

ed as complexes have a wide range of bond types from electronic to cova­

lent, and they possess a wide variety of properties.

Complexes were classified by Biltz^ according to stability. He

^W. Biltz, Z. anorg. Chem., l6U, 3U$ (1927).

called those that undergo reversible dissociation normal complexes, and

those that exhibit little or no reversible dissociation he called pene­

tration complexes. Whereas the term normal complex is associated mainly

2 3 with complexes whose bonding is of a loose nature possessing a high degree of ionic character* penetration complexes are those whose bond­ ing is essentially covalent* The classification is one of convenience rather than strict division*

The modern theory of complexes had its beginning in 1893 when

Alfred Werner^ published his theory explaining how it is possible for

5a . Werner, Z_. anorg* Chem,* 3 , 267 (1893)*

species capable of existence as entities to combine to form molecular

complexes* His theory involved the following ideas* Metals not only

have principal, or primary, ionizable valencies, but also auxiliary, or

secondary, valencies* The secondary valencies, which are non-ionizable,

combine with a maximum fixed number of atoms in a first sphere* This

coordination number is usually k or 6 and sometimes 2 or 8* Since they

are directional, six secondary bonds would form a regular octahedron

and four secondary bonds would form either a plane or a tetrahedron*

Therefore, isomerism is possible* Whereas secondary valencies may be

satisfied by anions or molecules, primary valencies are satisfied only

by anions less firmly attached in a second sphere*

In 1923 Sidgwick^ and Lowry? gave an electronic interpretation to

V. Sidgwick, J. Chem* Soc.* 123 , 725 (1923).

7T. M. Lowry, J. Soc. Chem* Ind., 1*2, 316 (1923). k the Werner theory* They considered Werner's primary valencies to be the formation of electrovalent bonds by the transfer of electrons* His secondary valencies consisted of the formation of electron-pair bonds by the donation of electrons to the atom by the coordinating group.

These are coordinate bonds which are indistinguishable from covalent bonds after they are formed, since the source of electrons in the electron-pair bond has no influence on the strength or direction of the bond* The idea of accumulating electrons on an electropositive metal atom required in the fomation of coordinate bonds is an inherent weak­ ness in their theory when it is applied to coordination compounds and complex ions* The hybridization of bond orbitals as set forth by

Pauling® explains in a very satisfying way how these strong covalent,

®L. Pauling, "The Nature of the Chemical Bond," 2nd. ed., Cornell University Press, Ithaca, 19k0, Ch. III.

directional bonds are fomed* If the three 2p orbitals and the s orbi­ tal of the tetravalent atom were used individually to form bonds, there would be three mutually perpendicular bonds and another weaker one without any particular direction* However according to the quantum- mechanical treatment, if there is a linear combination of the s orbital with the three 2p orbitals, there could be formed four strong bonds of equal strength and directed toward the corners of a regular tetrahedron*

This is consistent with the experimental facts* Also, d orbitals may be used in hybrid bonds to give a variety of arrangements according to Table I«? Whereas some of the possible bond types and corresponding

% . Moeller, "Inorganic Chemistry," John Wiley & Sons, Inc*, Hew York, N. Y., 1952, p. 203.

spatial arrangements are rare among complexes, others are quite common.

It is seen, for example, that metals with a coordination number of 6 may form d^sp3 bonds arranged octahedrally about the metal atom*

It is important that not all the bonds in complexes are covalent as the theory of hybridization of orbitals suggests, since the proper­ ties of many complexes indicate electrovalent bonds and covalent bonds containing considerable electrovalent character* The bond type and the configuration of the complex can be determined by measuring the of the complex, which is roughly a function of the number of electrons and the number of vacant d orbitals* A bond in a complex is usually neither entirely covalent or ionic, but it may be thought of as essentially covalent or essentially ionic*

A donor that is capable of coordinating at two or more positions with the same metal ion forms complexes containing ring structures*

These complexes were named chelates by Morgan and Drew^-O from the Greek

10G. T. Morgan and H* D. K. Drew, J* Chem, Soc,, 117, 1U56 (1920).

word,£yA<; , which means crab's claw* Donors that form them are called chelating agents* The additional stability that the complex has because 6

TABLE I

DIRECTIONAL CHARACTERISTICS OF COVALENT BONDS

Covalence Bond Type Spatial Arrangement

2 sp or dp linear (straight line) p2 , ds, or d2 angular

3 sp^, dp^, d^s, d^ trigonal plane dgp unsymmetrical plane p3 or d2p trigonal pyramid

h sp3 or d3§ tetrahedron dsp2 or d2p2 tetragonal plane d^sp, dp3, or d3p irregular tetrahedron tetragonal pyramid $ dsp3 or d3sp bi-pyramid d2ap2, d^s, d2p3, or dtyp tetragonal pyramid d3pZ pentagonal plane d5 pentagonal pyramid

6 d2sp3 octahedron d^sp or d5p trigonal prism d3p3 trigonal antiprism ? s p 2, d^s, d^p2 mixed

7 d3sp3. sp ZrFy-3(face-centered octahedron) d^sp2, d^p3, d^p2 TaFy" (face-centered trigonal prism)

8 d|*sp3 dodecahedron antiprism d^sp2d|P 2 face-centered prism 7 of chelation was called the "chelate effect" by Schwarzenbach,^- If

■^G. Schwarzeribach, Helv. Chim. Acta, 35, 23UU (1952),

the chelate contains fused rings it is even more stable than one contain­ ing single rings.

The crystal field theory, which is concerned with the splitting of d levels, has been quite helpful in recent years in explaining mapy experimental facts about coordination compounds and complex ions.- ^ “16

19 J. H. Van Vleck, "Theory of Electronic and Magnetic Susceptibilities," Oxford, The Clarendon Press, 1932,

13W. G. Penney and R, Schlapp, Phys. Rev., Ul, 191; (1932).

^ I b i d . , U2, 666 (1932).

^L. E. Orgel, J. Chem. Soc., 1952, U756.

l6L. E. Orgel, J. Chem. Phys., 23, 1819 (1955),

Chemistry of EDTA

EDTA, General Information and Properties

Ethylenediaminetetraacetic acid (EDTA), which is also called

(ethylenedinitrilo)tetraacetic acid in the literature, has the formula,

HOOC COOH nch2 ,ch2 ** N-CH2-CH2-N ^CH2 'CH2 HOOC ^ COOH

It has been of considerable interest in recent years because of its many applications as a completing agent. The sodium salts, ’which are sold under various trade names, such as Versene, Sequestrene, Nullapon,

Calsol, Chelaton, Complexone (III), Imidol D, Nervancid, Idranal (III),

Tetra Ver, and Trilon B, are used more commonly because they are much more soluble than the tetra acid and because complex formation occurs more readily at pH values higher than the tetra acid exists. For con­ venience in writing formulas, Y^~ represents the tetranegative anion of EDTA. Hence, the tetra acid and the sodium salts are represented conveniently as H^Y, NaH^Y, Na2H2Y, etc.

H[jY is a tetrabasic acid which can be titrated with a strong base to give a typical titration curve. Schwarzenbach and Ackermann^

■^G. Schwarzenbach and H, Ackermann, Helv. Chim. Acta, 30, 1798 (19H7).

determined its pK values at 20°C. in 0.1N KC1 to be 1.996, 2.672, 6.161, and 10.262. They state that since two of the protons show very strong acidic character, the tetra acid of EDTA exists as a zwitterion with the structure of a double betaine as follows:

H00C-CH2 CH2C00“ \ + + / nh-ch2-ch2-hn / \ "00C-CH2 ch2cooh

The tetra acid is a white, tasteless and odorless crystalline

which is only slightly soluble in water and all the usual organic solvents, but is soluble in hot formamide. The solubilities of EDTA and its sodium salts and the pH of the resulting solutions are shown in

Table II. 18

1 O ’’Sequestrene," Geigy Industrial Chemicals, p. 2.

TABLE II

SOLUBILITIES (grams/100 ml.) OF EDTA AND SODIUM SALTS IN WATER

AND pH OF AQUEOUS SOLUTIONS

22°C. li0°c. 80°C. pH

Ethylenediaminetetraacetic acid 0 .2 0 .2 0.5 2 .2

Monosodium ethylenediamine tetraacetate l,k 1 .U 2 .1 3.5

Disodium ethylenediamine tetraacetate 1 0 .8 13.7 2 3 .6 iu7

Trisodium ethylenediamine tetraacetate U6.S U6.S U6.5 8.U

Tetrasodium ethylenediamine tetraacetate 60 59 61 1 0 .6

Crabb-^ showed by the use of an automatic recording thermogravi

19^N, Crabb, unpublished work.

metric balance that the tetra acid simultaneously loses four moles of

CO2 in the range 110-230°C. He also showed by the same method that the

disodium dihydrate loses both moles of water at 98°C., two moles

of CO2 in the range 210-280°C., and two more moles of CO2 between 280° 10 and 3f?0°C. This is not in complete agreement with the usual under­ standing that EDTA and its sodium salts are stable on prolonged heating at l£0°C. and that the tetra acid melts by decomposition at 2hO°C.20

20 *lSequestrene,w Geigy Industrial Chemicals, pp. 2-3*

It is generally understood that EDTA and its sodium salts are attacked by the strongest oxidizing agents, such as permanganate and chromic acid, but that they are stable toward moderately strong oxidizing agents and toward reducing agents.21-23 EDTA is stable when it is boiled in strong

21 Ibid, p. 3.

22R. L. Pecsok, J. Chem. Educ., 29, 597 (1952). 23 "Keys to Chelation," The Dow Chemical Company, p. 7*

acid or alkaline solutions, but it is not stable above pH 12.

Preparation of EDTA

Ethylenediaminetetraacetic acid was first made commercially avail­

able in 1936 by the I. G. Farbenindustrie in G e r m a n y . ^ It was prepared

2^I. G. Farbenindustrie, German Patent 638,071 (1936).

by treating ethylenediamine with HCN and CH2O to form CNCH2NHCH2CH2NHCH2CN,

which was saponified with acid or alkali to yield HOOCCH2NHCH2CH2NHCH2COOH. 11

The treatment of this product with HCN and CH2O was followed by saponifi­ cation to give ethylenediandnetetraacetic acid.

In 1938 the General Aniline Works patented^-* a process for prepar-

#F. Miinz, U. S. Patent 2,130,505 (Sept. 20, 1938).

ing EDTA by the reaction of monochloroacetic acid on ethylenediamine.

By using the Strecker synthesis, Ulrich and Ploetz^ made EDTA by

2^H. Ulrich and E. Ploetz, U. S. Patent 2,200,995 (June 25, 19Uo).

the reaction of HCN with CH2O to produce HOCH2CN, which was allowed to

react with ethylenediamine,

A method was devised by Bersworth^7 in 19U5 for the preparation of

27F. C. Bersworth, U. S. Patent 2,387,976 (Oct. 30, 19U5).

the tetrasodium salt of EDTA. A solution of H 2NCH2C 00Na and C3H7NH is

refluxed as long as NH3 is evolved. The resulting solution of C3H7N-

(CH2C00Na)2 is treated with (03^ ) 2!® to give (C3H7 )NCH2COONa, which is

treated with ethylenediamine to yield (Na00CCH2)2N (CH2)2N (CH2COONa)2 .

Bersworth obtained a patent for another method which later was

28F. C. Bersworth, U, S. Patent 2,387,735 (Oct. 30, 19U5). 12 9 9 called the method of carboxymethylation of amines. The process

Smith, J. L. Bullock, F. C. Bersworth, and A. E. Kartell, J. Org. Chem., ll*, 355 (19U9).

involves the reaction of ethylenediamine with NaCN in basic solution at

10°C. during the slow addition of CH2O. Yields of the tetrasodium salt of EDTA amounting to 96 per cent of the theoretical were obtained. This method was preferred to the Strecker synthesis, which required handling large quantities of HCN, and to the reaction of chloroacetic acid or its

salts with ethylenediamine, rtiich results in the formation of mars­ hy -product s.

EDTA as a Chelating Agent

EDTA is potentially a hexadentate ligand because it has four

carboxyl groups and two tertiary amine groups which are capable of

coordination. If coordination occurs at all six positions a structure

containing five 5-membered rings would result. Evidence for the forma­

tion of hexadentate EDTA complexes was found by Schwarzenbach.3® He

Schwarzenbach, Helv. Chim. Acta, 32 , 839 (191*9).

showed that the reversible change from red [C0YH 2O] ” to blue [CoYOH} “

during the increase in pH was prevented by boiling the monoaquo complex 13 to form the hexadentate fCoY] ", Some doubt was still expressed31 that

31R. L. Pecsok, J. Chem. Educ., 29, $91 (1952).

the hexadentate actually existed. However, all doubt was removed when

Busch and Bailar3^ confirmed the hexadentate character of [Ccilj “ by

32D. H. Busch and J. C. Bailar, Jr., J. Am. Chem. Soc., 75, k$7k (1953).

examining its infrared spectrum and by resolving it into its optical

isomers. The pentadentate EDTA complex, also previously assumed to

exist, was confirmed by the same methods.

Palladium(ll) and platinum(II) complexes of EDTA, shown by

MacNevin and Kriege33 to exist in solution, were prepared by Busch and

33W. M. MacNevin and 0. H, Kriege, Anal. Chem., 26, 1768 (195U).

Bailar3** who proved their existence as both bidentate and tetradentate

3**D. H. Busch and J. C. Bailar, Jr., J. Am. Chem. Soc., 78, 716 (1956). “ “

complexes. Hence, it is established that EDTA uses all or part of its

six coordinating positions to form complexes. When it forms a complex

that is not a hexadentate with a metal whose coordination number is 6, ah the remaining positions are coordinated with other donors, such as

Hj>0, NH3, CN”, , etc.

Because EDTA always acts as a polydentate ligand forming 5-membered fused ring structures, the "chelate effect" described by Schwarzenbach^

35q. Schwarzenbach, Helv. Chim. Acta. 35, 23hh (1952).

exists in all its complexes. The magnitude of the formation constants

of a number of EDTA complexes shown in Table III^ ia quantitative

3&g <> Schwarzenbach, R. Gut, and G. Anderegg, Helv. Chim. Acta. 37, 937 (195U).

evidence for their unusual stabilities. The ethylene bridge between

the atoms contributes greatly to the stability of the complexes.

Pfeiffer and Simons^ showed that the complex of EDTA is some-

37p# Pfeiffer and H. Simons, Ber., 76b , 8U7 (191*3) *

what stronger than the corresponding methylaminediacetic acid complex.

This difference can be attributed only to the linkage through the

ethylene bridge which is lacking in methylaminediacetic acid. Further,

Schwarzenbach and Ackermann^® made a study of the stability of the

3 % . Schwarzenbach and H. Ackermann, Helv. Chim. Acta, 31, 1029 (19U8). 15

TABLE III

LOGARITHMS OF FORMATION CONSTANTS OF EDTA COMPLEXES

(20°C,yU=0.1)

Complex Log. K Complex Log. K

CaT 10.96 t 0.1* NdY 1 6 .6 1 t 0.05

CoY 16.31 ± 0.05 SiriY 17.11* t 0.05

nit 1 8 .6 2 ± 0 .0 6 EuY 17.35 t 0 .0 6

CuY 18.80 - O.lU GdY 1 7 .3 7 1 0 .0 5

ZriY 16.50 ± 0.02 TbY 17.93 t 0.05

CdY 16.146 ± 0.02 DyY 1 8 .3 0 ± o.o5

PbY 18. Oh 1 O.lli ErY 18.35 ± 0 .0 6

vor 18.77 1 0.05 TiriY 1 9 .3 2 - 0 .0 6

GaY 20.27 1 0.1 YbY 19.51 ± 0 .0 7

A1Y 16.13 ± 0.05 LuY 19.83 1 0.07 +

YY 18.09 ± 0.0k % (D 2 5 .1 ± 0 .1

MriY U 4.OI4 ± 0 . 2 TKY 2 3 .2 * 0 .1

LaY 15.50 ± 0.05 ScY 23.1 ± 0.15

CeY 15.95 ± 0.05 IriY 2U.95 ± 0 .1

PrY I6.I4O - 0.05 conqplexsis of the series (HOfXJC^^NCC^^NCCHgCOOH^, varying n from

2 to 5. They found that as n increases the complex becomes less stable.

If the two nitrogen atoms coordinate with the metal atom when n is 1* or

5, unstable 7 or 8-membered rings would have to form. For this reason the two ends of the ligand tend to act independently, forming MgY-type complexes.

EDTA forms complexes with practically all stable bivalent and polyvalent metal ions. In these complexes metal atoms and EDTA exist

in a 1:1 molar ratio. At least a 10 per cent excess of EDTA is usually

required to chelate the metal ion quantitatively, and in some practical

situations as much as 100 per cent excess is desirable. The formation

of the 1:2 complex has been observed only with tetravalent cations. It

is always very small compared with the quantity of 1 :1 complex, as

comparison of the formation constants of the two species of thorium(lV)

demonstrates.-^

39G. Schwarzenbach, R. Gut, and G. Anderegg, Helv. Chim. Acta, 37, 937 (199U).

The order of chelation among the metal ions depends upon the rela­

tive stabilities of the complexes which are expressed quantitatively by

their stability constants. Since the optimum pH for the EDTA complexes

varies with the metal ion, the order of chelation varies somewhat with

pH. For example, whereas the EDTA chelates of most metal ions have

their greatest stability in alkaline solution, ferric iron precipitates 17 as the hydrated , and trivalent aluminum and are converted to unchelated anions when acidic solutions of their EDTA complexes are made alkaline.

Because of the unusually high stability of the metal chelates of

EDTA, many of the usual chemical reactions of the metal ions are inhibi­ ted by chelation with EDTA. These chelates are also extremely soluble

in water. The combination of these two properties of the metal chelates

of EDTA makes this ligand unique in its sequestering action. For this reason the investigation of EDTA for a variety of uses has been in prog­

ress in recent years. Hundreds of papers in the literature are a wit­

ness to this fact. In addition to its variety of uses in industrial

processes and products, EDTA is used in water softening, medicine, agri­

culture, chemical separations, and analytical determinations.

Preparation of EDTA Complexes

There are several methods that lend evidence to the existence of

chelates or other complexes in solution. If metal ions do not undergo

their usual chemical reactions in a solution which contains a complex-

ing agent, they are assumed to be "tied up" in a stable complex. The

absorption spectra of solutions of complexes are different from those

of their aquo ions. Job rs method, or the method of continuous vari­

ations,^UO,Ul is based on this fact, may be used to show the molar

^°P. Job, Ann. chim. (). 9 (10), 113 (1928).

^W. C. Vosburgh and G. R. Cooper, J. Am. Chem. Soc., 63, U37 (1910-). ------18 ratio of metal to ligand in the complex. Electrical-conductance.methods show the disappearance of ions which are consumed in complex formation.

Metal ions react with a chelating agent by displacing one or more weak acidic protons of the chelating agent, producing a decrease in pH with chelation. Optical activity, solubility, oxidation potentials, magnetic susceptibility, polarographic, and other measurements are also useful tools. However, there is perhaps no more convincing proof of the exist­ ence of a complex than its preparation and confirmation by elemental analysis.

Mapy metal chelates of EDTA have been prepared. Pfeiffer and

Offermann^ isolated the first EDTA complexes in 19U2. They reported

^P. Pfeiffer and ¥. Offermann, Ber., 7$B, 1 (19U2).

that they heated a solution containing EDTA, NaOH, and Cu(0H )2 to boil­ ing and collected the unreacted Cu(0H )2 on a filter. By allowing the filtrate to stand after the addition of alcohol, Na2 [CuY] separated in the form of green crystals with a blue tinge. Pfeiffer and Offermann also reported that they prepared the corresponding salt

Cu/Cuf] *1^20 by heating a solution containing EDTA and Cu(0H )2 and allowing the product to precipitate from the hot solution. It was in the form of blue needles with a violet tinge. They prepared

K2[CaY] and ^[MgY] *5 ^ 0 by boiling aqueous solutions of EDTA and

KOH with the corresponding carbonate and by precipitating the complex by the addition of alcohol, A solution of the calcium complex gave no 19 precipitate 2k hours after the addition of oxalate, and the magnesium was not precipitated from a solution of its complex with sodium phosphate and *

Brintzinger and Hesse^ prepared the copper complex H2 CuY and

Brintzinger and G. Hesse, J. anorg. allgem. Chem., 2k9» 113 (191*2).

the complex Hp[NiY] by the addition of solid HJjY to solutions of

the metal salts. (U02)‘2 H 2Y *H20 was also prepared. In these com­ pounds the failure of copper and nickel to undergo their usual chemical

reactions confirmed conqplex formation. The U 02++ ion gave its usual

precipitation reactions, which showed that it formed a salt with H2Y* rather than a complex.

A number of EDTA complexes were prepared by Brintzinger, Thiele,

and Muller^- by adding solid H^Y to hot or boiling solutions of the

^ H . Brintzinger, H. Thiele, and U. Miller, Z. anorg. u. allgem. Chem., 251, 285 (19li3).

metal salts. These included NafCoY] #UH2 0, H[brY], HjFeY] , NHj^/Fel? ,

H[NdY] , and a complex containing a 2:1 molar ratio of to ligand.

Since its solution gave the usual reactions of Pb++, it was probably a

lead salt of the lead complex, Pb PbY . These workers obtained

crystals of HLaY and UY*2H20 which were either normal complexes or

tertiary amine salts, since the metal ions responded to their usual 20 reagents. A thorium compound of a similar nature was given the formula

Th2f 2*^2^ as a double molecule because three-fourths of the water was more easily removed than the remainder.

Schwarzenbach and Biedermann^ isolated the complex H(CrYH20] from

k-’G. Schwarzenbach and W. Biedermann, Helv. Chim. Acta, 31, U59 (19W).

which the mole of water could not be removed in high at 100°C.

This work established the existence of complexes in which EDTA serves as a pentadentate ligand with water coordinating in the sixth position,

Schwarzenbach^ produced considerable evidence for the pentadentate

Schwarzenbach, Helv. Chim. Acta, 32, 839 (19U9).

nature of EDTA by preparing a series of cobalt complexes having the

general formula MfCoYxVn^O, in which M stands for H+, NH^+, or one of

the alkali metals, and X stands for H2O, CNS", Cl", or Br". When Br" was removed from its conplexes by the addition of Ag+, M[CoY] *nH20 type

complexes were formed.

Calcium, strontium, barium, , bismuth, zinc, and cadmium

complexes of EDTA were prepared by Brintzinger and Munkel^, starting

^H. Brintzinger and S. Munkelt, Z. anorg. Chem., 256, 65 (19U8),

with EDTA and either the carbonate or the hydroxide of the metal. The 21 zinc and the cadmium complexes were found to be weakly coordinated with

EDTA. In attempts to prepare the EDTA complexes of lithium, beryllium, and yttrium, they found the simple salts as established by the usual reactions of the metal ions.

Among other EDTA complex preparations that may be mentioned are the preparation of the rare complexes by Moeller, Moss, and

Marshall^-® and the platinum and complexes by Busch and Bailar.^

^T. Moeller, T. A. J. Moss, and R. H. Marshall, J. Am. Chem. Soc., 21, 3182 (1955). ^D. H. Busch and J. C. Bailar, Jr., J. Am. Chem. Soc., 78, 716 (1956).

EDTA in Qxidation-Reduction Reactions

EDTA is said to be attacked slowly by strong oxidizing agents, such 50 5l as permanganate and chromic acid, to form cyclic ureides. 3

5o "Key to Chelation— The Versenes," Bersworth Chemical Company, Farmington, Massachusetts, p. 5.

^ F. J. Welcher, "The Analytical Uses of Ethylenediaminetetraacetic Acid," D. Van Nostrand Coupany, Inc., Princeton, New Jersey, 1958, p. 2.

It is generally considered stable and inert toward most chemical com­ pounds and sufficiently resistant toward moderately strong oxidizing agents to be used in peroxide bleaching baths and to be included in powdered perborate bleaching products. There is some evidence in the

52 "Sequestrene," Geigy Industrial Chemicals, p. 3. 22 literature that EDTA is more easily oxidized than is usually assumed*

Schwarz enbacir-' found that EDTA is decomposed with the formation of

^ G . Schwarzenbach, Helv. Chem. Acta. 32 . 839 (191*8).

formaldehyde, and some of the cobalt (III) is reduced when rather concen­ trated solutions of complex [Col] ” are heated for a rather long time.

Gold compounds are reduced to the metal by EDTA. This reaction serves as a spot test for , and it is capable of detecting 0 .1 per cent

gold in solution.In photochemical studies of riboflavin, Merkel and

pt4W. A. Hynes, K. L. Yanowski, and J. E. Ransford, Mikrochemie ver Mikrochim. Acta, 35, 160 (1950).

Nickerson-^ observed the photoreduction of methylene blue in the presence

J. R. Merkel and W. J. Nickerson, Biochim. et Biophys. Acta, ll*, 303 (19510. ------

of EDTA. A further study of this behavior with near ultraviolet light

by Os ter and Wotherspoon^ showed that EDTA was consumed in the reaction.

-^G. Oster and N. Wotherspoon, J. Am. Chem. Soc., 79, 1*836 (1957).

They concluded that, although EDTA is resistant to reducing agents

under normal conditions, it appeared that it was oxidized in that reaction. They suggested that EDTA could possibly be oxidized to amine , since amine oxides are formed by the oxidation of tertiary amines by hydrogen peroxide.

Analytical Uses of EDTA

Since the first use of EDTA in an indirect determination of metal ions by Schwarzenbach in 19b$j more than a thousand papers on the use of EDTA in inorganic analysis have appeared in the literature. Accord­

ing to W e l c h e r , ^ in addition to the use of EDTA as a primary standard,

d 7 ^ ‘F. J. Welcher, "The Analytical Uses of Ethylenediaminetetraacetic Acid," D. Van Nostrand Company, Inc., (1958), p. 9.

all of those uses may be classified under three general principles:

1. The titration of metallic ions with standard EDTA solutions

2. Colorimetric determinations based on measurement of the color

of some EDTA complexes

3. The use of EDTA to prevent interference in the determination

of various metals

EDTA has a number of properties that make it a suitable reagent

for use in volumetric methods. In the first place it reacts rapidly

and stoichiometrically with a large number of metal ions without

successive reactions or side reactions. The use of EDTA in direct

titrations, in which a standard solution of EDTA is added to a solution

of metal ions, is quite common. The equivalence point may be determined

in a variety of ways. Metal indicators, which are organic compounds

that form weak complexes with metal ions, are commonly used for this 2k purpose in a solution buffered at pH 10 with an ammonium hydroxide- avnmonium chloride buffer. Some of the common metal indicators are

Eriochrome Black T, , Metalphthalein, Tiron, and Pyrocatechol

Violet. "When visual observation of the end point is difficult, the titration is often followed spectrophotometrically. If an unbuffered solution of metal ions is titrated with a standard solution of tri­ sodium or tetrasodium salt of EDTA, there is a sudden rise in pH at the equivalence point when hydrolysis of EDTA occurs according to the following reactions:

HT3” + H2O » H 2Y “ + 0H~

+ 2H20 = H 2Y “ + 20H“.

Metal ions and the EDTA conplex of the same metal have different oxida­ tion potentials. For this reason there is a sudden change in the poten­ tial of an indicator electrode at the equivalence point in the titration

of a solution of metal ions with EDTA. This makes potentioraetric titra­

tions possible. Amperometric titrations may also be made by titrating

at a constant applied potential until the diffusion current drops to

the residual level. Because the electrical conductance of a solution

of metal ions is changed by complexing the metal with EDTA, it is often

convenient to do conductometric titrations. High frequency titrations

are also possible.

Sometimes it is convenient to add an excess of EDTA to a solution

of metal ions and back-titrate the excess of EDTA. This approach is

applied to the determination of metals which react too slowly with EDTA

for a direct titration or which precipitate as hydrated oxides at the

pH required for the titration. Another method involving the addition of an excess of standard EDTA solution is the alkalimetric method.

-^G. Schwarzenbach, ¥. Biedermann, and F. Bangerter, Helv. Chim. Acta, 29, 811 (I9l|6).

A standard base is used to titrate the protons released in the reaction,

+ H2I" - Mr(n”k) + 2H+.

Some metal ions whose EDTA complexes are much more stable than the magnesium complex can be treated with an excess of standard solution of

[Mgl]= releasing a quanitity of Mg++ which is equivalent to the unknown metal ion M++, according to the following reaction:

+/Mg$= - + Mg++.

Then the Mg++is titrated with a standard EDTA solution by using a metal indicator. It is possible to determine many anions indirectly by titrat­ ing an equivalent quantity of cations with a standard EDTA solution.

If the ions of a metal are colorless, the EDTA chelate of that metal is also usually colorless. However, there is usually a deepening of the color of the ions of the transition metals on complexing them with EDTA. This makes it possible to determine many metals spectro­ photometrically with EDTA.

It is well known that EDTA reacts with ions of many metals to form complexes that are so stable that the ions do not respond to their usual reagents. In such cases EDTA finds frequent use as a masking agent to prevent the interference of many cations during the determination of

others. This principle is applied in almost all types of analytical methods. 26

Chemistry of Platinum Group Metals

Oxidation Potentials of Rhodium and Iridium

Latimer^ summarizes the potentials of rhodium and iridium as

^ M. Latimer, "The Oxidation States of the Elements and their Potentials in Aqueous Solutions,” 2nd ed., Prentice-Hall, Inc., Hew York, N. Y., p. 219.

follows:

Acid Solution

Rh Rh+— Rh++— Rh3+— Rho++—Lrlgfl, Rho^ (-0.8)______

-o.UU RhCl6c ca»-l»2 RhClA8

<-1.1 Ir++ >-1.1

Tr (-1.15) - Tr3+ (-0.7) Ir02_ >-1.3 IrO^ 1

-0.93

-0.77 -1.017 IrC16‘ IrCl6

IrBr£' -0.99 IrBr'

Basic Solution

-O.Oli Rh- Rh^O^— -0.9 Rh02- -0.9 RhO^

- 0.1 - 0.1 Ir- - Ir203- -Ir02------Ir0^= 27

The stability of +h iridium is well known from the number of compounds and complexes that exist in this state, of which IrCl^” is a common example. However, it is a moderately strong oxidizing agent, comparable to bromine, stronger than , and capable of oxidiz­ ing ferrous iron. These facts may be confirmed by considering the E° values for the couples involved.

IrCl6= = IrCl6= + e~> E° = -1.01760

^®F. P, Dwyer, H. A. McKenzie, and R, S, Nyholm, J, Proc. Roy. Soc, N. S. Wales, 81, 216 (19U7).

2Br~ = Br2 (l) + 2e“, E° = -1.065261

^G. Jones, and S, Baeckstrom, J, Am, Chem. Soc., 5>6, 1J?2U (193U).

NO + 2H20 = NO3" + UH+ + Ue", E° = -0.9662

^^W. M. Latimer, "The Oxidation States of the Elements and their Potentials in Aqueous Solutions," 2nd ed#, Prentice Hall, Inc., New York, N. Y., p, 93.

Fe + Fe + e , E° = -0.77163

63Ibid. 3 p. 223.

The IrCl^3, IrCl6= couple, being a 1-electron change, is rapid enough to

be used for titrations. The recent literature includes numerous exam­

ples of this. Separations and Analyses of Platinum Group Metals

The method of Gilchrist and Wichers^ is the one that has had the

^R, Gilchrist and E. Wichers, J. Am. Chem. Soc., 57 3 2565 (1935).

most use during the last twenty years for the separation of the platinum

group metals from each other and for their subsequent determination. If

the is present as an alkaline osmate or as a bromo-osmate, the

osmium can be distilled in about one hour from a solution containing

10 per cent nitric acid. If it is present as a chloro-osmate, concen­

trated sulfuric acid should be substituted for the nitric acid in order

to reduce the seven- or eight-hour distilling time. If is

absent, a mixture of sulfuric and nitric is used. The OsO^ is

distilled into a 1 :1 solution saturated with

dioxide. It is customary to convert the platinum group metals remain­

ing in the distilling flask to their chloro acids and fume them with

sulfuric acid to convert them to their sulfates. After the solution is

diluted with water, sodium bromate is added, and the ruthenium is

distilled as RuO^ into a 1 :1 hydrochloric acid solution saturated with

. If iridium is present, it is important to do the oxida­

tion with bromate when the solution is sufficiently acid in order to

avoid complications which lead to incomplete separation of the ruthenium.

The osmium is recovered by precipitating it with sodium bicarbonate as

the hydrated oxide from a hot solution at pH U, as indicated by the

change from yellow to blue of broraphenol blue. The oxide is collected

in a Monroe platinum and ignited to the metal in a stream of 29 hydrogen, in the presence of . The metal is allowed to cool in a stream of and weighed. The hydrated ruthenium oxide is also precipitated hot with sodium bicarbonate, but in this case the pH is controlled with bromcresol purple. It is ignited to the metal

in a manner similar to osmium and cooled in a stream of hydrogen. The next step is the separation of palladium, rhodium, and iridium from platinum by a double precipitation as the hydrated oxides in the presence

of bromate. The oxidizing medium decreases the extent of hydrolysis of

the platinum, and it is effective in producing a more satisfactory pre­

cipitate of the other oxides and doing it more quickly. The precipi­

tation is carried out hot with sodium bicarbonate at pH 8. The platinum

is precipitated as the , ignited, and weighed as the metal.

After the combined precipitates of the hydrated oxides of palladium,

rhodium, and iridium are dissolved in hydrochloric acid, the palladium

is precipitated as a dimethylglyoximate and weighed as such. The

organic matter is destroyed in the filtrate containing the rhodium and

iridium by adding sulfuric acid and nitric acid and evaporating several

times to fumes of . T5.tanous chloride is used to reduce

the rhodium to the metal, which is collected on a filter and dissolved

in nitric acid. The rhodium is precipitated as the sulfide, which is

ignited in air and then in hydrogen, and cooled in hydrogen. It is

weighed as the metal. The is removed from the filtrate con­

taining the iridium with . The iridium is now precipitated

again as the hydrated oxide, which is ignited in air and in hydrogen.

The metallic iridium which is formed is cooled in a stream of hydrogen

and weighed. BACKGROUND OF THE PROBLEM

In 195U Kriege^ carried out a study of the reactions of EDTA with

6*0. H. Kriege, Ph. D. dissertation, Ohio State University (195U).

the platinum group metals in chloride solution. He presented spectro- photometrie, complexomatric, potentiometric, and chemical evidence for a chelate of divalent palladium with EDTA. He used Job's method to confirm a 1 :1 ratio of metal to ligand in the complex and showed by migration studies that the complex is an anion. The mean value of the

logarithm of its stability constant was determined as l8.£ - 0 .6 at

2S>°C. and at an ionic strength of 0.2 in the pH range 3*72-8.95.^ It

66W. M. MacNevin and 0. H. Kriege, J. Am. Chem. Soc., 77, 6lU9 (19#).

served as the basis for a coraplexometric titration^ and a spectro-

^W. M. MacNevin and 0. H. Kriege, Anal. Chem., 27, £35 (1955?)•

photometric determination of palladium. At about the same time this

68Ibid., 26, 1768 (1951*).

complex was prepared as the bidentate and converted to the tetradentate

30 31 by Busch and Bailar,^ Kriege concluded that tetravalent palladium was

6% . H. Busch and J. C. Bailar, Jr., J. Am. Chem. Soc., 785 716 (1956).

reduced to the bivalent state before complexing with EDTA, because the ultraviolet and visible absorption spectrum of a solution containing

PdCl^1* and EDTA was similar to that of a solution of the divalent palladium-EDTA complex.

Kriege found complex formation between trivalent ruthenium and

EDTA based on spectrophotometrie, eomplexometric, and chemical evidence.

He used the results of Job's method and eomplexometric evidence in support of a 3:1 ratio of metal to ligand in the complex. He found no complex formation in chloride solutions between EDTA and trivalent rhodium, trivalent iridium, tetravalent osmium, or divalent or tetra­ valent platinum. However, Busch and Bailar^ prepared a bidentate

chelate of divalent platinum and converted it to a tetradentate in a manner similar to that for divalent palladium,

A complex of tetravalent iridium and EDTA formed in chloride solu­

tion was also reported by Kriege, again supporting his conclusion with

spectrophotometric, eomplexometric, and hydrolytic evidence. He showed

that the ultraviolet and visible absorption spectra of chloride solu­

tions of tetravalent iridium changed radically with the addition of EDTA. 32

By adding EDTA to an acid solution of tetravalent iridium and back- titrating the unreacted EDTA with a standard zinc sulfate solution, and by using Eriochrome Black T as indicator, he confirmed a reaction between the tetravalent iridium and EDTA. He showed that the addition of EDTA to an acidic solution of tetravalent iridium prevents the imme­ diate formation of the hydrated iridium oxide when it is neutralized or made mildly basic. He found that the addition of EDTA to a solution containing tetravalent iridium prevents the formation of iodine when potassium iodide is added. Kriege interpreted these four phenomena as evidence of complex formation between tetravalent iridium and EDTA.

When Kriege added a large excess (at least 3:1 molar ratio) of

EDTA to a solution of tetravalent iridium, the reaction went stoichio- metrically, consuming two moles of iridium for every mole of EDTA.

This behavior was interpreted as the formation of a 2:1 complex

(iridium:EDTA), and it was used as the basis for a eomplexometric method for iridium. However, when the quantity of EDTA in the reaction mixture was decreased, the ratio of iridium to EDTA reacting increased. When the ratio of EDTA to iridium in the reaction mixture reached 0.5:1, the ratio of iridium to EDTA reacting reached 2.6:1. This was interpreted by Kriege as probable formation of a mixture of 3il and 2:1 (iridium to

EDTA) complexes, for which he postulated the formulas f^I^IClgj] = and

When a large excess of EDTA was added to tetravalent iridium solu­ tions, and the pH was adjusted to the range ll.U-12.6 a maximum occurred 33 at 313 njp. which obeyed Beer's law. This was used by MacNevin and

Kriege?1 as the basis for a spectrophotometric method for iridium.

71 W. M. MacNevin and 0. H. Kriege, Anal. Chem., 28, 16 (1956).

Dunton7^ studied the electromigration rates of the complexes of

?^M. L, Dunton, Ph. D. dissertation, Ohio State University (1956).

the platinum group metals with EDTA reported by Kriege. A separation of platinum, palladium, rhodium, and iridium by paper electrochromatography was obtained as a result of complex formation of tetravalent iridium and divalent palladium with EDTA.It73 was reported that the migration

73W. M. MacNevin and M. L. Dunton, Anal. Chem., 29, I 806 (1957).

of tetravalent iridium toward the anode was increased greatly by the addition of EDTA, and its mobility became greater than that of any other platinum group metal species studied. Heating the tetravalent iridium-

EDTA solutions to 8o°~90°C. caused a decrease in the mobility of the

iridium. This behavior was explained as a result of converting the

iridium-EDTA complexes to their hydroxy or aquo forms. Trivalent rhodium

in the presence of EDTA was reported to form a precipitate at the point

of application in electrochromatographic studies. EXPERIMENTAL

Reactions of Iridium(IV) with EDTA

Job’s Method Applied to Iridium(IV) and EDTA

It was necessary to obtain additional evidence for the molar ratio

of the metal to EDTA in the tetravalent iridium-EDTA complex, since the

2:1 and 3 si ratios reported by Kriege were unusual for EDTA complexes.

It was also strange that at least a 3:1 molar ratio of EDTA to iridium was required in the reaction mixture for the quantitative formation of

a conplex containing two moles of iridium per mole of EDTA.

Job’s method, or the method of continuous variations, for determin­

ing the formula of a conpound or complex in solution is applicable to

reactions of the type

A + nB * ABn ,

in wich the dissociation constant is not very large. The method is

based on the fact that if a suitable property of the complex and the

same property of the dissociated forms are measured and their differ­

ence Y is plotted as a function of the composition X, then Y is a maxi­

mum at the composition corresponding to the formula of the complex.

The value of n is expressed by the formula

where n is the number of moles of B combined with one mole of A, X is

the mole fraction of B, and 1 - X is the mole fraction of A. The

property is measured on eleven solutions, each containing the same

total concentration of A and B and varying by tenths from pure A to

3k pure B. If at least one of the species absorbs light, the solutions may be analyzed spectrophotoraetrically.

A series of eleven such solutions of IrCl£,= and EDTA was prepared and allowed to stand at room temperature for equilibrium to be estab­ lished, as shown by constant absorbance readings. The absorbance of each solution was then measured at k3% and h90 mji. The results, plotted in Figure 1, show that two moles of tetravalent iridium react with one mole of EDTA. This is consistent with Kriege’s observations of a 2:1 complex. Moreover, the straightness of the lines and the sharp point that they form indicate a reaction that is nearly complete.

Other Preliminary Studies

After the additional evidence for the formation of a 2:1 complex

of tetravalent iridium with EDTA in chloride solution was provided by

the results of Job’s method, it was deemed expedient to prepare the

complex. However, some preliminary experiments were carried out first

in order to determine the most suitable conditions for the crystalli­

zation of the complex in order that large quantities of expensive rea­

gents would not be lost in unsuccessful attempts. Some of the first of

these experiments were done on a micro scale by combining drops of the

reactants on a microscope slide and watching for the formation of crys­

tals. This approach was soon abandoned because of the high solubility

of the product and of the impossibility of controlling concentrations,

temperature, and pH.

The next series of experiments was carried out with very small

quantities of concentrated solutions of H2lrCl6 and of EDTA. At any Enhancement, 4 . 0 - - - - - 0.2 -1.4 0.6 0.8 1.0 0 iue . o' Mto fr DA n Iiim (E7) Iridium and EDTA for Method Job's I. Figure 0.2 , oe rcin f EDTA of fraction mole x, 0.4 0.6 o 4 9 0 mju. 0 9 4 o A 4 3 5 m/j. 5 3 4 A 0.8 1.0

36 37 concentration, tetravalent iridium precipitates quantitatively as the hydrated oxide at pH U and above in the presence of bromate. Partial precipitation occurs above pH 3, especially in rather strongly concen­ trated solutions and warm solutions. As EDTA is slightly soluble below pH 2 and least soluble at pH 1.5, it precipitates as the tetra acid below pH 2. However, its solubility increases again below pH 1.5, but

only to a limited degree. For these reasons, it was necessary to com­

bine the reactants in the pH range 2-3 where EDTA exists as a mixture

of mostly H3Y“ with smaller quantities of H^Y and H 2Y“,7U or to work in

^R. L. Pecsok, J, Chem. Educ., 2£, 597 (1952).

rather dilute solutions at pH 1 or less, where EDTA exists only as H[,Y.

Complex formation at such low pH values is expected to be slow, because

protons compete with the metal ions for the ligand and because mass

action effects in dilute solutions are small.

When a nearly saturated solution of EDTA was added to a freshly

prepared solution of tetravalent iridium so that the pH of the mixture

was about 2.5, there was an almost immediate change in color from the

opaque, deep red-brown of IrCl6= to a very pale yellow or straw color.

If the iridium solution was aged, it turned pale green upon the addi­

tion of EDTA. This phenomenon is explained by the partial hydrolysis

of the IrClg®. In both cases, effervescence was observed. At pH 1 or

less the reaction was slow, even when the solution was heated on a

steam plate at 65-75°C., and the final color was usually cherry red at

maximum concentrations for those conditions. When the red solutions 38 were diluted they turned yellow, and their absorption spectrum was almost identical with that of reaction mixtures of IrCl^ and EDTA at pH 2-3. The absorption spectra of solutions of H2lrC16 and of solu­ tions containing ^IrCl^ and EDTA at two pH values are shown in

Figure 2. The H2lrCl6 absorption spectrum has bands at 305> Ul7, UUO, i;90, and 580 nyi. As it absorbs very strongly in the region 385-330 nju, it has a very deep red-brown color, and except for very dilute solu­ tions, it is opaque. A solution containing the same concentration of

H2lrCl6 and an excess of EDTA absorbs very little above 280 mp, but it has low peaks at 333 and 393 mp. It is evident from Figure 1 that the presence of a very small quantity of unreacted IrCl^*5 is expected to have a great effect on this spectrum and to shift the peaks from 333 and 393 tji.

Attempts to Prepare Iridium(lV)-EDTA Complex

In experiments with small quantities, 0.0010 mole of IrClj^ was added to 0.0010 mole of the disodium salt of EDTA in 5.6 ml. of water.

When a dark colored precipitate of hydrated iridium oxide began to form after some of the IrClj^ was added, a few drops of concentrated HC1 were added to dissolve it. When part of the IrCl^ was added the solution changed to yellow. When the remainder of the IrCl^ was added, the solu­ tion turned red from the presence of the IrCl£= in solution. During the first 20 minutes on a steam plate at 65°-75°C. its color changed to cherry red and remained that color while the heating was continued for the next hour. Three fractions of impure NaCl were precipitated by evaporating this solution under an air stream. Further evaporation Absorbance (I cm. cell) 0.0 0.8 0.4 2.0 200 iue . bopin uvs o HICg n HIC6 rae wt EDTA with Treated H2IrCI6 and H2IrClg for Curves Absorption 2. Figure J a-a-A-Arnwwvn/yvA-rv- —J— 400 aeegh myu Wavelength, 600 2rI pu ET, H 1 . 0 pH EDTA, 2.8 plus pH EDTA, H2IrGI6 plus i6 H2IrC H2IrCI6 0.00078 M 0.00078 iw a a w i - p — A 800 k jJ V VO Uo resulted, in the formation of a brown syrup which dissolved readily on the addition of a little water. The addition of alcohol to this solu­ tion precipitated a red-brown oil which was removed from the mother liquor* When more alcohol was added to the oil, it became a tar-like

semisolid. Since the properties of this substance were unlike those

of either reactant, it was apparently the reaction product. However,

the quantity of it was too small for convenient analysis.

Because the removal of NaCl in several fractions was not only

time-consuming but wasteful of material, it was desirable to avoid con­

tamination with NaCl by starting with the tetra acid of EDTA rather

than the disodium salt. Then much of the HC1 could be removed by vola­

tilization while the volume was reduced. In using this approach, it

was found that the total color change occurred in less than 30 minutes

when the reaction was carried out in a solution as acidic as 1:1 HC1

and containing a 1:2 molar ratio of H 2lrCl£, to EDTA. Heating was con­

tinued one and one-half hours longer. After the solution was evaporated

to dryness under an air stream, the addition of distilled water produced

a solution with a pH of 1.5 at which EDTA is least soluble. Nearly half

the original EDTA added was precipitated as H^Y, and it was collected on

a filter. When the filtrate was concentrated under an air stream, three

fractions of red-brown tar-like material were precipitated by the addi­

tion of alcohol. The determination of iridium by X-ray fluorescence^^>

7*W. M. MacNevin and E. A. Hakkila, Anal. Chem., 29, 1019 (1957).

7^E. A. Hakkila, Ph.D. dissertation, Ohio State University (1957). Ill in the three fractions which were dried over Drierite yielded 20.6,

23.1, and 26.0 per cent iridium, respectively. The percentages of iridium in the first two fractions corresponded closely to the theoret­ ical 2 2 .1 per cent iridium in a complex containing one mole of iridium to two moles of EDTA. Such a complex could be expected to have the formula /HglrYjj) ++ in acid solution, and the structure ++ o-8. ch2 H0-6-CH^ CH2-CH2 n \ Q :id n - cho - 6-oh HO-C-CH2

Ir( I VY 0 ! /,CH2-C-0H H 0-$-CH2lT (SCH2-CHi'X NCH2-6'-OH ch2 - 0 5

This ratio was not expected from theoretical considerations nor from the usual behavior of EDTA in complex formation. It is not consistent with Kriege’s work?? or with the results of Job’s method in this study

7?0. H. Kriege, Ph.D. dissertation, Ohio State University (195k) •

in which two moles of iridium(IV) were found to react with one mole of

EDTA. Hence, there was positive evidence for both a 2:1 complex and a

1:2 complex of iridium(lV) and EDTA obtained under similar conditions.

Larger quantities of the tar-like product were prepared from reac­

tion mixtures which contained a 1:2.U molar ratio of IrCl5= to EDTA and

were adjusted to pH 1 with HC1. The procedure was similar to that k2 described above involving precipitation by the slow addition of alcohol to aqueous solutions. After the product was re-dissolved and re-precipi­ tated three times, it was dried over ^2®$ an

Galbraith Microanalytical Laboratories for elemental analysis. Iridium was determined by X-ray fluorescence. Atomic ratios of the elements in a complex containing four nitrogen atoms, based on typical results, are shown in line 1, Table IV.

TABLE IV

ATOMIC RATIOS OF ELEMENTS BASED ON ELEMENTAL ANALYSIS OF PRODUCT

CHN Cl Ir

Original 1 6 .8 Uo.U U.oo 6.8o 1.13

First Re-purification 15.3 37.5 U.oo - 1.18 Second Re-purification 16.5 3 6 .0 U.oo 6.78 -

The product was re-purified twice more in the same way and it was

returned for analysis. The atomic ratios from these results are shown

in lines 2 and 3, Table IV. It is seen that the ratio of carbon to

nitrogen in the product is not $tl as it is in EDTA, but Usl. Since no

nitrogen other than that in the EDTA was added to the reaction mixture,

and since a decrease in the carbon to nitrogen could occur only by loss

of carbon, there was decomposition of the EDTA in ■vhich two carbon

atoms were lost from each EDTA molecule. Fragments of two EDTA mole­

cules and about seven atoms were associated with each iridium atom. Since iridium(IV) has a coordination number of 6, it is impossi­ ble for it to coordinate at more than six positions. If it formed a complex with fragments of two EDTA molecules by using four of the six positions, a maximum of two positions would be left for chloride. This suggests a possibility of some chloride in the solution as an inpurity.

The product was extremely soluble in water, but no other solvent was found that would dissolve it. This behavior is consistent with that of many EDTA conplexes. Its aqueous solution varied between straw color and yellow. Its absorption spectrum in a solution in which the iridium was 0.0073M is shown in Figure 3. By comparing Figure 2 and

Figure 3 and considering the difference in concentration, it is seen that a solution of the product has a spectrum identical with the spec­ trum of the reaction mixture of H2lrCl6 with EDTA. Both have absorp­ tion maxima at 333 and 393mp.

Titration of Iridium(IV)-EDTA Product

When a solution of the product isolated from the reaction of l^IrCl^ and EDTA is treated with solutions of AgNC^, Hg(NC>3)2, or

Pb(1103 )2 , a tan precipitate containing all the iridium is formed, leav­ ing a clear, colorless supernatant liquid. It was feasible to make use

of this reaction by titrating such solutions with a standard solution of AgN0 3, in order to determine whether or not the extra chlorine existed as an impurity in the form of chloride. These titrations were

carried out potentiometrically by using a -silver chloride elec­ trode. Two end points were obtained. The first, involving a reaction which was too slow to be suitable for a titration reaction, occurred at 2 . 4

2.0

0) o oe

0) o c o X) S 0.8 X3 <

0.4

0.0 200 3 0 0 4 0 0 5 0 0 6 0 0 7 0 0 8 0 0 Wavelength, fi m

Figure 3. Absorption Curve for Product Isolated from Reaction of H2IrCI6 and EDTA. Ir = 0.0073M +0.13>U v4 vs. the S. C. E. The second, involving a rapid reaction, occurred in the range +0.2U to +0.29 v. vs. the S. C. E. The volume of standard AgNO^ solution required to reach the first end point was quite small compared with that required for the second. For comparison, solu­ tions of H2lrC16 and solutions of NaCl were titrated with end points also in the range +0.2U to +0.29 v. vs. the S. C. E. When NaCl solu­ tions were titrated, the end point fell in the same range. When NaCl t was added to solutions of the product, the first end point required the same volume of standard AgNO^ solution as when NaCl was absent, and it occurred, as usual, at +0.15U v. vs. the S. C. E. However, more stand­ ard AgN03 solution was required for the second end point which occurred, as usual, in the range +0.2U to +0.29 v. vs. the S, C. E. Hence, the presence of chloride could not be determined by this experiment, and no quantitative information was obtained; but it did suggest the presence of at least two species in the product.

Chromatographic Separation of the Product

Chromatographic separation of the product prepared from the reac­ tion of IrCl^= and EDTA was suggested by the titration with AgN03 solu­ tion, which indicated that the product contained at least two species.

A column 2.5 cm, X 20 cm. was prepared with powdered, Catalyst

Grade, Activated, Chromatographic Alumina obtained from the Harshaw

Chemical Company. When some of the product was placed on this column and developed with distilled water, two bands were observed. The one, which was easily eluted with water, was colorless and possessed an amine

odor. The other was a straw to tan colored band which contained all the iridium and which was held tenaciously at the top of the column. This substance could be eluted with a 1 per cent solution of nitric acid.

A quantity of this eluate was collected and evaporated by a Rinco evap­ orator. When the volume became rather small, the color turned from straw to deep red. As IrCl^” is deep brown-red and IrCl^2 is very light in color, it was considered possible that the change in color was caused by the oxidation of trivalent iridium to IrCl^- by the concen­ trated nitrate solution. This was confirmed by treating the straw- colored eluate with various oxidizing agents, such as Ce(S0^)2> CI2, and HNO^. In each case the solution turned deep red immediately. The original straw color was restored by adding a reducing agent, such as hydroquinone, dihydrochloride, or hydroxylamine hydrochloride.

When the reaction mixture of H 2lrCl6 and EDTA was treated with oxidiz­ ing agents, the color changed from straw to deep red. Here again, the straw color was restored by treating the solution with reducing agents.

These phenomena suggest that the reaction of IrCl6~ with EDTA is an oxidation-reduction reaction in which IrCl6= is reduced to IrCl65, rather than complex formation.

Reactions of Iridium(lV) with Na?H^Y

In considering the probability that a complex would form more

easily with the disodium salt of EDTA than with the tetra acid, it was decided to attempt to prepare the complex in solutions at higher pH.

It was hoped that a complex could be formed and isolated in this way without the decomposition of EDTA and the reduction of IrCl£=. These reactions were run by adding the disodium salt of EDTA to chloride solutions of tetravalent iridium and then adjusting the pH to the range

3-5 at room temperature. Kriege observed that after IrCl^“ reacted with

EDTA, the pH of the solution could be increased without precipitating the iridium as the hydrated oxide. This observation was confirmed in the present work. The product was precipitated by the slow addition of alcohol to the reaction mixture. In one preparation the product was re-dissolved and re-precipitated seven times. After it was dried over

P20^ in vacuo, the elemental analysis gave results similar to those of

the product obtained from more acidic solutions. The atomic ratios of

the elements in a compound containing four nitrogen atoms, based on the

elemental analysis, are shown in Table V.

TABLE V

ATOMIC RATIOS OF THE ELEMENTS BASED ON ELEMENTAL ANALYSIS

OF PRODUCT PREPARED FROM SOLUTIONS AT pH 3-5

c HN Cl Ir

16.6 U3.6 U.oo 6.92 1.07

In this case the C:N ratio is close to U:1 as before, and there is

nearly a 7il ratio of Cl:Ir. This is interpreted as a IrCl^" ion and

a Cl“ ion associated with two fragments of decomposed EDTA. U8

Reduction of Iridlum(IV) with EDTA

Previous experiments have suggested that IrClg*" is reduced to

IrCl^ 5 by EDTA* and no complex is formed. The speetrophotometric evi­

dence that this is the reaction that occurs is most conclusive.

The change in the absorption spectrum of a solution containing

0.00078M IrCl£= -when EDTA was added has already been discussed and

shown in Figure 2. The electrolytic reduction of O.OOO78M IrCl^" to xa IrCl^- was carried out on a Sargent-Slorain Electrolytic Analyzer by

using platinum electrodes and a 1 :U sulfuric acid solution containing

hydroxylamine hydrochloride as a depolarizer for the anolyte. The

catholyte was contained in a porous porcelain cup which was placed in

a beaker containing the anolyte. The absorption spectra of this solu­

tion before and after electrolytic reduction of the IrCl^ 3 are shown in

Figure U. It is seen that after electrolytic reduction to IrCl^-, the

same absorption spectrum is obtained as that of a solution of IrCl^*

reduced with EDTA, shown in Figure 2. When the solution of trivalent

iridium prepared by electrolytic reduction of tetravalent iridium was

treated with EDTA and heated, the small effect of heating was the only

change observed in the absorption spectrum. This is interpreted as evi­

dence for the absence of complex formation between trivalent iridium

and EDTA in chloride solution. Since a complex of iridium(lV) or

iridium(III) with EDTA would have absorption spectra very much unlike

that of IrCl^5, the speetrophotometric evidence is considered conclusive

proof that iridium(lV) is reduced to iridium(HI) by EDTA. Absorbance (I cm. cell) 2.0 2.4 0.4 0.8 0.0 200 Figure 4. Absorption Curves for H2IrCI6 and H3IrCI6 Formed Formed H3IrCI6 and H2IrCI6 for Curves Absorption 4. Figure y lcrltc euto o H IrCI6 H2 of Reduction Electrolytic by aeegh mp. Wavelength, r 0. M 8 7 0 0 .0 0 = Ir 600 04 0 8 0 0 U9 50

In order to add evidence to this fact, a solution of 0.0073M IrCl£= was reduced with hydroxylamine hydrochloride, and its absorption spec­ trum was compared with that of a reaction mixture of 0.0073M IrClg= and

EDTA. In preparing these solutions, the reactants were combined and

the solutions were heated in a water bath and cooled to room temperature.

The volume was adjusted to 0.0073M iridium, and the pH was simultaneously

adjusted at 3.U. The absorption spectra of these solutions are shown in

curves 1 and 2, Figure 5. The similarity of these curves shows that the

absorbing species is also IrCl^- in the solution containing EDTA. The

absorption spectrum of a solution of IrCl^- is affected by the amount of

heating it has undergone. The presence of any unreacted IrClg,- has a

great effect on the IrCl6= absorption spectrum since IrCl6= absorbs

strongly in the region 3 8 5 -5 3 0iiji where IrCl^" absorbs only slightly.

The oxidation products of EDTA and of hydroxy lamine hydrochloride inter­

fere at 300 mp and below. The differences in curves 1 and 2 are attrib­

uted to these three interferences.

Curve 3, Figure 5, shows the absorption spectrum of a solution pre­

pared by dissolving the product isolated from a H2lrCl6-EDTA reaction

mixture. Comparison of this curve with curves 1 and 2 shows that the

absorbing substance in the isolated product is essentially IrCl^S,

The evidence for the reduction of IrCl6= to IrCl6= with EDTA is

summarized as follows;

1. A solution containing IrCl£ 5 formed by the reduction of IrCl6=

with hydroxylamine hydrochloride or by electrolytic reduction

had the same absorption spectrum as a reaction mixture of Absorbance (I cm. cell) 200 Figure 5. Absorption Curves for IrCI^ IrCI^ for Curves Absorption 5. Figure H 2 IrCI 6 ih DA n Hdoyaie Hydrochloride Hydroxylamine and EDTA with 0 600 400 aeegh mji Wavelength, .a H a 2. . o rdc fo H from Product 3.o . H □1. Ir = 0.0073 M 0.0073 = Ir 2 2 lrCI IrCl 6 6 ET rato mixture reaction EDTA - HONH _ omed Form 3 I ecin mixture reaction CI 2 IrCI

by 6 euto of Reduction n EDTA and 800 51 52

IrCl£= and EDTA and as a solution of the product isolated from

such a mixture*

2* All products containing iridium -which ware separated from reac­

tion mixtures of IrCl^*3 and EDTA were shown by elemental analy­

sis to consist of deconposition products of EDTA in which the

ratio of carbon to nitrogen was changed from 5 :1 to U:l, con­

firming decomposition of EDTA.

3. The gas evolved when EDTA was added to solutions of IrCl^° was

identified as CO2. It was formed by the oxidation of EDTA.

it. In order to obtain a stoichiometric reaction in which two moles

of IrCl^- reacted with one mole of EDTA, Kriege78 found that

7 80. H. Kriege, Ph.D. dissertation, Ohio State University (195U).

three moles of EDTA were required for each mole of IrCl^~ in

the reaction mixture. In this study two moles of tetravalent

iridium were found to react with each mole of EDTA by Job's

method. Yet a product was separated from this mixture which

contained one mole of iridium to two moles of decomposed EDTA.

These strange ratios are indicative of an oxidation-reduction

reaction, especially since a possible couplex was expected to

contain a 1:1 ratio of iridium to EDTA,

By adding an excess of the disodium salt of EDTA to solutions of

IrCl£~ and titrating the unreacted EDTA with a standard ZnSOj^ solution ■with Eriochrome Black-T as an indicator, Kriege7^ was able to determine

79Ibid.

the quantity of EDTA that reacted with IrCl£ , and thus the molar ratio

of the reactants. He found that when the excess of EDTA was large,

(at least 3:1 molar ratio of EDTA to IrCl^") 0.5 mole of EDTA reacted

with each mole of IrCl^". But as the ratio of EDTA to IrCl^** in the

reaction mixture reached 0.5:1, the ratio of iridiura(IV) to H)TA that

reacted was 2.6:1. He interpreted this reaction as the formation of a

possible mixture of 2:1 and 3:1 ir idium(IV) -EDTA complex.

In order to investigate this strange behavior, a series of reac­

tions was carried out in this study with IrCl^** in large excess. The

IrCl6= was varied from two-fold to sixteen-fold, and the reaction mix­

ture was heated for 12 minutes on a steam plate. An iodometric method

for the determination of tetravalent iridium, was developed by Delepine 8o

80M. Delepine, Ann. Cham. (Paris). 1 (9), 277 (1917).

in which an excess of potassium iodide was added to the acidic solution

of IrCl^-. The equivalent quantity of iodine released was titrated with

a standard sodium thiosulfate solution. By using this method to deter­

mine the unreacted tetravalent iridium in the reaction mixtures, the

quantity of tetravalent iridium entering into the reaction, and conse­

quently the molar ratio of iridium to EDTA, was calculated. The data

in Table VI show that approximately four moles of IrCl^** react with su each mole of EDTA when an excess of about four moles of IrCl^** is present per mole of EDTA.

TABLE VI

MOLAR RATIO OF IRIDIUM(IV) : EDTA IN THE PRESENCE

OF EXCESS IRIDIUM(IV)

Molar Ratio Taken Molar Ratio Reacted EDTA : IrCl6= Iridium(lV) : EDTA

1 :2 2.0 :1

1 :U 3.7:1

1 :8 U.2 :l

1 :1 6 U.Otl

In reaction mixtures containing EDTA and IrCl6= in the molar ratio of

1:2, all of the IrCl^“ present reacted. Figure 6 shows these data plotted rith Kriege*s for reactions in which EDTA was in excess. This plot shows that tetravalent iridium reacts with EDTA in the molar ratio of 2:1 when EDTA is in large excess and in a molar ratio of approxi­ mately 2j.:l when IrCl5 3 is in large excess. Moles iridium (E) reacting per moleEDTA 2 3 4 F igure 6. Dependence of Molar R atio of R eactants on Moles Moles on eactants R of atio R Molar of Dependence 6. igure F Reacting oa rto rI : EDTA : IrCI« ratio Molar IrClg in excess excess in IrClg o oa rto DA IrClc EDTA: ratio Molar EDTA rdu (ET): Iridium n excess in EDTA 56

Postulated. Mechanism

The following equations are written to explain the reactions in terms of the data.

Excess EDTA in the reaction mixture

H00CCH2 CH2C00”

1. 2IrClA“ + NCH2CH2N = 2IrCl63 / \ -oocch2 ch2cooh

hoocch2 gh2

+ NCH2CH2N + 2C02

•ch2 ch2cooh

free radical

ch2cooh HOOCCHo CHo \ 2 / 2 ch2 ch2 or dimerizes 2. nch2ch2n I I / \ CHg -CEr 2 or polymerizes •CHr c h 2c o o h CH2C00H

CH2C00H ch2coo“

CH2 CH2 CH2 CHo 3. | | = 1 I + 2H+ ch2 ch2 ch2 ch2 •N ch2cooh CH2C00‘ Excess IrClA= in the reaction mixture

CH2COO" CH2 /N n ch2 ch2 ch2 ch2 in 2IrCl6s + I I - 2IrCl62 + j 1 + 2C02 CH2 CH2 CHo .CHo N X N 2 ch2coo" qh2

These equations account for the experimental facts which were

observed up to this point in this study, and they correlate data which were anomalous when complex formation was assumed. Equation 1 shows

the reaction of two moles of IrCl^*8 with one mole of the disodium salt

of EDTA, which was shown to be quantitative when EDTA was in large

excess. It accounts for the evolution of C02 by decarboxylation of

the EDTA at the carboxylate ions, resulting in the formation of free

radicals which can combine in various ways to yield an assortment of

products. Whether these combine by cyclization, dirnerization, or

polymerization, they are expected to be strong acids. As dissociation

occurs according to equation 3, or more likely by successive reactions,

the reaction shown in equation I4. is expected to proceed. In a large

excess of IrCl6“, reactions 1-U are expected to take place, both as

shown and stepwise, resulting in an assortment of products in which the

molar ratio of carbon to nitrogen is between U :1 and 5>:1 , but closer to

U:l. This is substantiated by the lnl3>:l ratio of carbon to nitrogen

in the isolated product. 58

Additional Evidence for the Mechanism

If equation 1 is correct for the reaction in a large excess of

EDTA, it should be possible to confirm it by the evolution of one mole of CO2 for every mole of IrClgT present. Also, if the series of reac­ tions goes to completion according to equations 1 -U in the presence of a large excess of IrCl6=, four moles of CO2 should be evolved for each mole of EDTA present, and pH-lowering should be measured that corre­ sponds to two moles of H+ ion for each mole of EDTA present.

An ordinary CO2 train that is commonly used for the gravimetric determination of CO2 in limestone was constructed. The H2lrCl6 and the

EDTA were combined in a C02~free atmosphere in the reaction flask.

After the reaction was allowed to proceed for at least 10 minutes while the reaction flask was warmed in a water bath, 10 ml. of 1:1 HC1 was added and the solution was boiled for 5 minutes. The CO2 was absorbed

in Ascarite and weighed. The results showed that when at least three moles of EDTA per mole of IrCl^ 3 were present in the reaction mixture,

0.96 mole of CO2 was recovered per mole of IrCl^1* present. The devia­

tion from unity is mostly experimental error. These results confirm

the quantitative evolution of CO2 from a reaction mixture containing a

large excess of EDTA according to equation 1.

When the reaction mixtures contained a large excess of IrCl5“, a

maximum of 3.25 moles of CO2 was recovered per mole of EDTA present.

In four experiments 3.25, 3.16, 3.00, and 2.77 moles of CO2 were recov­

ered per mole of EDTA present. The pH drop in similar reaction mixtures

during the reaction of IrCl£,= and EDTA corresponded to the release of

approximately 1.7 mole of H+ ions per mole of EDTA present. Apparently, in the presence of excess 11 *01 5 ", reactions 1 —14. do not go to completion when the pH drops low enough to prevent complete dissociation of the acid according to equation 3.

Discussion

Speetrophotometric proof has been presented that tetravalent iridium is reduced to trivalent iridium by EDTA. This reaction pro- ceeds instead of the formation of a complex a^ previously reported, x

0 « t 0. H. Kriege, Ph.D. dissertation, Ohio State University (19!?U).

The occurrence of the oxidation-rednction reaction is supported by-

several other types of experimental evidence. The reaction proceeds with the evolution of CO2 whose quantity depends on the molar ratio of

the reactants. It was shown by the evolution of CO2 and the consumption

of EDTA that the reaction proceeds quantitatively when EDTA is in large

excess, but the additional reaction that occurs when tetravalent iridium

is in large excess does not always go to completion. The product that

was isolated was shown by elemental analysis and speetrophotometric meas­

urements to contain one mole of IrCl£2, two moles of decomposed EDTA and

excess chlorine. This product contained a ratio of carbon to nitrogen

that accounted quantitatively for the CO2 evolved. Although the postu­

lated mechanism predicts the possibility of several organic species in

this product, it appears to be essentially an amine salt, such as I or II. 60 CHgCOOH

CHo CHp 1 I , IrCl(<2, Cl" ch2 ,ch2 NH i c h 2c o o h

I

H+ H+ HOOCCHoN -CHo -CHp -NCHpC 0 0 H \ \ ch2 ch2 1 | , Irci65, Cl' ch2 ch2

HOOCCH2 N-CH2 -CH2 -NCH2COOH H+ H+

II

The cation in formula I is the product of equation 2 in acid solution formed by cyclization of the product of equation 1. The cation in formula II is the dimerized product of equation 1 in acid solution.

Salts of this type containing a tertiary amine and a metal chloride com­ plex are not uncommon.

Since the reaction between tetravalent iridium and EDTA was shown to be an oxidation-reduction reaction, it is necessary to explain in terms of oxidation-reduction the experimental results that Kriege ftp interpreted as evidence of coirplex formation. The change in absorption

82Ibid. 6l spectrum of a solution of IrCl^“ by the addition of EDTA does not necessarily serve as evidence for complex formation of iridium(lV) with

EDTA, The new absorption spectrum was shown in this study to be identi­ cal with that of IrCl^ 2 formed by the reduction of IrCl^“ by other meth­ ods, He showed that in the presence of a large excess of EDTA, half a mole of EDTA disappeared for every mole of iridium present. It was shown in this study that the EDTA was consumed by oxidation rather than complex formation, Kriege observed that a solution of tetravalent iridium treated with EDTA became more resistant to hydrolysis when the pH of the solution was increased. It is well known and easily verified by a simple experiment that trivalent iridium is more resist.**^ than tetravalent iridium to hydrolysis.

Assuming a lsl iridium(lV)-EDTA complex Hakkila reported^ that it

^E. A. Hakkila, Ph.D. dissertation, Ohio State University (1957)*

behaved as an anion when prepared and kept in acid solution, but in basic solution it changed to a cationic form which did not revert to the anion upon re-acidifying. Hakkila also reported that when the pH of a solution containing IrCl£- was increased to 10 and re-acidified, much cationic migration occurred. The results of this study, which

show that EDTA reduces IrClg- to IrCl6=, explain the similar behavior

found by Hakkila in a solution containing IrCl6= and a solution pre­ pared from IrCl£= and EDTA, since iridium exists as IrCl£= in each

solution before each is made basic. 62

Evidence of Rhodium(III)-EDTA Complex

Kriege^ reported that he found no evidence for the formation of

H. Kriege, Ph.D. dissertation, Ohio State University (19E>U)*

an EDTA complex with rhodium(III) in chloride solution. If the chloride complex is so stable that it prevents the formation of the EDTA complex, it should be possible to form the rhodium(III)-EDTA complex by starting with the rhodium in the form of its sulfate. Since sulfate has a much less tendency than chloride to form complexes, the rhodium-EDTA complex should form more easily in sulfate solution.

The rhodium was obtained for this study from The American Platinum

Works as RhCl^. It was necessary to convert some of it to the sulfate.

One milliliter of a 0.028M solution of RhCl^ in hydrochloric acid was treated with a solution until it reached pH 8, and it was warmed in order to precipitate the rhodium as Rh2°3« The Rh2°3 was dissolved in sulfuric acid and the pH was adjusted to about 1 with

sodium hydroxide. The resulting yellow solution was treated with 1,0 ml,

of 0,0f>M solution of the disodium salt of EDTA, and it was heated on an

electric hot plate for one and one-half hours. There was no appreciable

drop in pH, Its volume was adjusted to the mark in a 10 ml, volumetric

flask as its pH was adjusted to 3,0 with sodium hydroxide solution. Its

absorption spectrum is shown in Figure 7> curve 1, In a similar experi­

ment the excess sulfate was removed from solution by treating it with barium hydroxide rather than sodium hydroxide to adjust the pH to 1, and Absorbance (I cm. cell) 0.6 0.2 0.4 0.8 1.8 r iue 7. Figure bopin uvs o Roim H) DA Complex EDTA - (HI) Rhodium for Curves Absorption 0 5 3 0 0 4 vlnt, mp. avelength, W . h + hoie EDTA + chloride + Rh3 + 2. . hl + EDTA + RhClT 3. . h+ • uft + EDTA + sulfate +• Rh3+1. 0 550 500 0 5 4 the BaSO^ precipitate was removed by collecting it on a filter. Then when the EDTA was added and the solution was heated, there was a drop in pH which served as evidence of complex formation with the EDTA.

Apparently, the buffering action of the more concentrated sulfate solu­ tion prevented a pH-drop during complex formation.

Since those experiments, Hakkila 8*5^ found speetrophotometric evi-

®^E. A. Hakkila, Ph.D. dissertation, Ohio State University (1957).

dence for the formation of a rhodium(III)-EDTA complex by using freshly o/r prepared solutions of rhodium chloride. McKay showed that the freshly

^E. S. McKay, Ph.D. dissertation, Ohio State University (1956).

prepared red solutions contain anionic rhodium and the aged solutions containing cationic rhodium are yellow. McKay also formed the yellow cationic rhodium by dissolving RI12 O3 in acid. The red solutions con­ tain RhCl6“> and the yellow solutions contain the aquo Rh^+ion. Par­ tially aged solutions probably contain all the intermediate species.

In this study it was found that the formation of the rhodium(III)-

EDTA complex in a 0.0028M RhCl6 5 red solution is accompanied by a pH drop from 3*55 to 2.28. This occurred in a 30-minute heating .

By starting with the yellow form of rhodium in chloride solution, the pH dropped from 3.38 to 2.58 during complex formation. The absorption

spectra of these two solutions are shown in Figure 7j curves 2 and 3, respectively. All of these curves have absorption maxima at 350 mn and are similar to the absorption spectrum of the rhodium(III)-EDTA reported by Hakkila. They show that apparently the same EDTA complex forms whether EDTA is added to a red solution of RhCl^5, a chloride solution of the aquo ion of Rh^+, or the sulfate solution of that ion. However, comparison of the magnitude of absorbance at the absorption maximum,

3f>0 rap, shows that the reaction goes nearest to completion when start­ ing with the RhCl^2 solution. This is consistent with the fact that the aquo ion is more stable than the chloride complex, since the aquo ion forms when the chloride solution is aged. The drop in pH after the addition of EDTA adds evidence for complex formation in each solution. CONCLUSIONS

A study of the reaction of EDTA with iridium(lV) in chloride solutions has indicated the following:

1. The reaction of EDTA with iridium(IV) in chloride solutions

is an oxidation-reduction reaction, and no EDTA complex is

formed.

2. The IrCl5~ coiqplex is reduced to IrCl6=.

3. The oxidation of EDTA proceeds by a decarboxylation reaction.

li. If EDTA is in j.-;vge excess (at least 3:1 molar ratio), the

reaction is stoichiometric, reducing two moles of IrCl^- to

IrCl£S per mole of EDTA and evolving two moles of CO2.

5>. If IrCl^ is in large excess, as many as four moles of IrCl^*

are reduced to IrCl<$s per mole of EDTA. This reaction does

not usually go to completion.

6. The product isolated from EDTA-H2lrCl6 reaction mixtures was

apparently a mixture of amine salts containing IrCl^=, Cl",

and degradation products of two moles of EDTA.

7. After the reduction of IrCl^” to IrCl£“ by EDTA, the EDTA in

excess does not form a complex with the iridium(III).

A study of the reaction of EDTA with rhodium(lll) has indicated the following:

1, EDTA forms a complex with rhodium(IIl) in sulfate solution.

2. The rhodium(III)-EDTA complex that forms in sulfate solution

is the same one that forms in chloride solution.

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Werner, A., Z. anorg. Chem., 3, 267 (1893). AUTOBIOGRAPHY

I, Harold Dale McBride, was born near Mansfield, Ohio, August 18,

1918. After receiving my secondary education in the public schools of

Madison Twp., Richland County, Ohio, I was employed as a factory worker in Mansfield, Ohio, for three years, and attended college for two years at The Ohio State University and Otterbein College. I served from

January, 191+2, to December, 19U5, in the Army and in the Air Force where

I was commissioned Second Lieutenant. During World War II, I served as a Radar Observer-Night Fighter in the South Pacific. In January, 191+6,

I returned to The Ohio State University for three quarters. I was em­ ployed as office manager at the Cement Products Company in Mansfield,

Ohio, and as a mason and factory worker in Ashland, Ohio, until Septem­ ber, 1950, when I returned to The Ohio State University which granted me the Bachelor of Arts degree in December, 1951. I attended the gradu­ ate school there from January to July, 1992, when I became employed as a Research Assistant in chemical research by the Kettering Foundation.

I held that position for a year full-time and then part-time until receiving the Master of Science degree from The Ohio State University in June, 195U. I served alternately as Assistant and Research Assistant until October, 1956, when I was appointed Assistant Instructor. I served in that capacity until July, 1957, when I received a National Science

Foundation Fellowship which I held until completing the requirements for the degree Doctor of Philosophy.

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