THE KINETICS AND MECHANISMS OF THE IRIDIUM(IV) -

SULFUR(IV) REACTION IN A BUFFERED

CHLORIDE MEDIUM

by

ELDON LLOYD STAPP, B.S.

A THESIS

IN

CHEMISTRY

Submitted to the Graduate Faculty of Texas Tech University in Partial Fulfillment of the Requirements for

the Degree of

I4ASTER OF SCIENCE

Approved

Accepted

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ACKNOWLEDGMENTS

I am sincerely grateful to Dr. D. W. Carlyle for his

instruction and encouragement during the course of this

study. I am also indebted to Dr. Carlyle along with the

other members of my committee. Dr. John N. Marx and Dr.

Jerry L. Mills, for their helpful criticism of this thesis

I wish to acknowledge the financial support from the

Robert A. Welch Foundation (1972-73).

11

g7

CONTENTS

ACKNOWLEDGMENTS ii

LIST OF TABLES iv

LIST OF FIGURES V

I. INTRODUCTION SECTION 1

II. EXPERIMENTAL SECTION 11

Reagents 11

Stopped-Flow Apparatus 17

Kinetics Experiments 20

Stoichiometry 23

ESR 25

III. RESULTS SECTION 27

Stoichiometry of Iridium(IV)-Sulfur(IV) Reaction 27

Kinetics of the Iridium(IV)-

Sulfur(IV) Reaction 28

Iridium(III) Retardation 39

Copper(II) Effect 42

Iron(II) Effect 42

Radical Investigation 43 Equilibrium Study 46

IV. INTERPRETATION SECTION 50

Reduction of Hexachloroiridate(IV) by Sulfur(IV) 50

Copper(II) Catalysis of the Reduction of Hexachloroiridate(IV) by Sulfur(IV) 57

Conclusions 59

LIST OF REFERENCES 61

• • •

111

LIST OF TABLES

Table Page

1. Results of Stoichiometry Measurements Performed at 25°C and Ionic Strength 0.2, With 0.03M Acetate Buffer 28

2. Rate Constants for Iridium(IV) and Sulfur(IV) Reaction at Ionic Strength 0.2M, Buffer Strength at 0.03M, and Iridium(IV) at 1.3 X 10-4M 30

3. Acidity and k^j^^^ at 6.0 x 10~^M Sulfur(IV), 1.3 X lO-^M Iridium(IV), 0.03M Buffer, 0.2M Ionic Strength, and 25°C 33

4. Excess Iridium(IV) and Sulfur(IV) Observed Rate Constants at 0.2M Ionic Strength and 0.03M Sodium Acetate 37

5. The Effect of Iridium(III) on the Observed Rate Constant for the Iridium(IV)-Sulfur(IV) Reaction 40

6. Effect of Copper(II) Catalysis on the Iridium(IV)-Sulfur(IV) Reaction 44

7. Reaction of Sulfur(IV) with Various Compounds In the Presence of Aqueous Acrylic Acid 45

8. List of Absorbances for Various Concentration of Sodium Sulfite at 1.71 x 10"^M H+ and 0.06M Sodium Acetate 47

IV

|7

LIST OF FIGURES

Figure Page

1. Structure of sulfur dioxide 7

2. Structure of sulfite ion 7

3. Structure of pyrosulfite 7

4. A typical photograph of an oscilloscope trace of the disappearance of iridium(IV). This particular photograph is from experiment 16 of Table 2 32

5. A typical plot of Absorbance -Absorbanceg^ versus reaction time using the conditions given in Table 2 for experiment 16 32

6. Plot of acidity versus k , , with the data and conditions given m'Table 3 34

7. Plot of [S(IV)] versus k , ^ [H"*"] with iridium(IV) 9 x lO'^M, 0?2^i6nic strength, 0.03M sodium acetate buffer, 25°C, and the data points taken from Table 2 35

8. Plots of k , ^ versus excess iridium(IV) concentration using the data presented in Table 4. O, D, and A represent 6.80 x 10"^, 1.95 X 10"5, and 1.55 x IO'^M H+, respectively 38

9. A typical oscilloscope trace for experiment no. 1 from Table 6 for the copper(II) catalyzed iridium(IV)-sulfur(IV) reaction 43

10. Plot of Absorbance (corrected) versus total sulfur(IV) concentration squared taken from the data presented in Table 8 48

CHAPTER I

INTRODUCTION SECTION

The object of the study reported in this thesis is the

mechanistic study of the oxidation of sulfur(IV) in acidic

aqueous medium as shown by the two following equations.

2IrClg^~ + HSO3"" + H2O -^ 2IrClg^" + SO^^" + 3H"'" (1)

2IrClg^" + 2HSO2" -> 2IrClg^~ + S20g "' + 2H"^ (2)

Evidence will be shown for the possible existence of a radi­

cal intermediate. Catalysis by copper(II) for reactions 1

and 2 is also investigated in this study.

David w. Carlyle (1) has suggested that the reduction

of tris(1,10-phenanthroline)iron(III) by sulfur(IV) may 2-

occur by a mechanism involving pyrosulfite ion, S2OC- , as

the sulfur(IV) species. One of the purposes of this study

was to prove or disprove a mechanistic step involving

iridium(IV) and pyrosulfite. It is hoped that this thesis

will stimulate future study in the Cu(II) catalyzed reaction

of iridium(IV-sulfur(IV) reaction and the reaction of Cu(II)

with sulfur(IV).

Sulfur(IV)—The number of studies involving sulfur(IV)

reductions is huge and the discussion of past sulfur(IV)

reductions will be limited to inorganic molecules reacting

with sulfur(IV). As shown in equations 1 and 2, the

sulfur(IV) product may be sulfate, dithionate, or as in most

cases, a mixture (2). Dithionate has been shown to be un­

stable thermodynamically with respect to disproportionation

(3), but has been shown to be kinetically stable with re­

spect to both disproportionation and oxidation (2).

Sulfur(IV) has been shown to react with various 1-

equivalent and 2-equivalent inorganic oxidizing agents (4).

Bunau and Eigen (5) have studied the reaction of sulfur(IV)

with I2. Both references 4 and 5 have shown that the pri­

mary product for 2-equivalent oxidizing agents is sulfate.

Kinetic and stoichiometric data have been published for the

reaction of sulfur(IV) with Pd(II)(6), Fe(CN)g" "" (4, 7-10),

Cr(VI) (11, 12), Pt(II)-02 complex (13), a variety of manga­

nese complexes (14), aquoiron(III) complexes (14, 15, 16),

tris(1,10-phenanthroline)iron(III) (1), and other metal

ions (4). The above work and the stoichiometry work done by

Marshall and Higginson (4) show that one-equivalent oxidiz­

ing agents often, but not always, yield a mixture of sulfate

and dithionate when reacting with sulfur(IV). Considerable

effort has been given to understanding why some one-

equivalent oxidizing agents yield primarily sulfate while

others yield a mixture of sulfur products. The following

information has been stated by many workers (1, 2, 4, 10,

14, 17, 18). This information, which is a more general

application regarding the products of the reaction of a

metal ion with sulfur(IV), was summarized by Brown and

Higginson (17). Their generalized conclusions are that a

substitutionally inert metal ion plus sulfur(IV) yields pri­

marily the product sulfate, whereas a substitutionally labile

metal ion reacting with sulfur(IV) yields appreciable

dithionate.

As has been already mentioned, a large amount of kinetic

data has been gathered concerning reactions with sulfur(IV)

and metal ions (1, 6-16). A typical system is the oxidation

of sulfur(IV) by hexaquoiron(III) (14, 15, 16). These work­

ers ,suggest a sulfur(V) radical formation in their mechan­

isms. They have shown competition for the radical by Fe(II)

and Fe(III). This was suggested by evidence of marked in­

hibition of the Fe(III)-S(IV) reaction by Fe(II). The

workers also reported a strong catalytic inhibition effect

of oxygen on the Fe(III)-S(IV) reaction. Again they believed

this was possible proof for a radical as the oxygen may have

been scavenging any radicals present. Karraker (15) and

Carlyle (16) have proposed that the sulfur(IV) species is

in the inner coordination sphere of the iron(III) when the

first electron is transferred. Karraker (15) then proposes

that the sulfur(V) radical represented as -SO^H, reacts with

another Fe(III) to give sulfate. Carlyle (16) however, pro­

poses a HO~-SO-. radical intermediate which in separate

steps yields dithionate and sulfate on further reaction

with Fe(III). Carlyle's mechanism fits both his stoichio­

metric and kinetic observations.

Another typical reaction that has been studied exten­

sively is the Fe(CN)g -S(IV) reaction (4, 7-10). Swinehart

(8) favors an outer-sphere mechanism where the S(IV) is not 3-

coordinated to the Fe(CN)g . Veprek-Siska et al. (10) have

suggested that the mechanism involves an inner-sphere step 3_

where the S(IV)-Fe(CN)g are coordinated before the first

electron is transferred. Swinehart's proposed mechanism

involves radical formation, whereas Veprek-Siska's mechanism

does not.

Swinehart searched for the proposed sulfur(V) radical

via esr but had negative results. However the radical could

have been present in quantities sufficient for the mechanism

proposed by Swinehart but insufficient for detection by esr.

4-

Retardation by Fe(CN)g , as shown by Swinehart, was consis­

tent with a radical mechanism. An alternative to both mech­

anisms is that of Lancaster and Murray (7). Murray's

5-initial step involves the formation of a Fe(CN)^(CNSO^)

intermediate. This intermediate reacts with another

3- 4- 2-Fe(CN)g to give Fe(CN)g and SO^ . Murray used radio-

35 active S and a scintillation counter to prove the existence

of the intermediate. Murray (7) also reported that the pres­

ence of ferrocyanide retarded the rate of reduction of

ferricyanide by sulfite.

Because of the rich variety of mechanisms involved in

oxidations of sulfur(IV), reliable mechanistic generaliza­

tions have not been discovered. Another of the objectives

of this work was to add to the kinetic and stoichiometry

data for sulfur(IV) oxidations, and to try to improve the

understanding of these oxidations in general.

Because many other workers (1, 4, 8, 14, 15, 16) have

postulated the presence of sulfur(V) radicals when sulfur(IV)

is oxidized by metal ions, still another objective of this

paper was to get evidence concerning the existence of a

radical. The proposed radical" would account for the produc­

tion of dithionate when the sulfur(V) species dimerizes

before the radical can react with the one-electron metal

oxidant (4, 14, 16). Both spectrophotometric and esr data

have shown that free sulfur(V) radicals can exist in aqueous

solutions (19-24). Other work by esr has shown that coordi­

nation sometimes occurs between sulfur(V) and other radicals

and metal ions (25-28). Some workers (1, 4, 8, 14, 16) have

proposed that the sulfur(V) species exists as a free radical

in aqueous solution while others (7, 10) have theorized

sulfur(V) coordination to the metal ion.

A variety of sulfur(IV) species exist in aqueous solu­

tion. Sulfur(IV) used in this reaction is made by bubbling

gaseous SO2 into water. This solution is commonly considered

to be sulfurous acid hydrate, H2S02«6H20, but according to

Cotton and Wilkinson (29) the species is actually S02*7H20

which is a clathrate of the same type as other gas hydrates.

Other workers (30, 31) have found through uv, ir, and Raman

spectra that sulfurous acid is the minor species in aqueous

sulfur dioxide. This clathrate structure is believed to in­

volve sulfur dioxide molecules solvated by six medium cages

and two small cages containing a total of 46 water molecules,

Bisulfite and sulfite salts are well known. Even though

sulfur dioxide aqueous solutions are gas hydrates it takes

two equivalents of base to neutralize every sulfur dioxide

molecule. The following equilibria investigated by many

workers (21, 32-35) help explain the neutralization by a

base.

Q3 - + SO2 + H2O t HSO-^ + H^ (3)

- Q4 2-2HSO3 Z ^2°5 " 2° ^^^

Q3 = [HSO3"] [H"^] = 0.043 (32) (5)

[SO2]

^2^5^" Q. = - - 5- = 0.07 (35) (6) ^ [HS03-]'^

The structures for SO2 (29), SO3 ~ (29), and S^O^^" (36) are

shown in Fig. 1, 2, and 3, respectively. The point groups

for sulfur dioxide, sulfite ion, and pyrosulfite ion are

Cy , C3 , and C respectively. The molecules are resonance

s

Fig. 1.—Structure of sulfur dioxide

:t5: - -I s:

Fig. 2.—Structure of sulfite ion.

:o :

:p: Fig. 3.—Structure of pyrosulfite

8

stabilized, though only one resonance structure is shown

for each.

Iridium(IV)—Hexachloroiridate(IV) is a fairly strong

substitution-inert oxidizing agent as shown by equation

7 (37).

IrClg^" ^ IrClg^" + e" E'' = -1.017v (7)

The sodium salt of hexachloroiridate(IV) is extremely solu­

ble in water. It hydrolyzes slowly but is indefinitely

stable in dilute hydrochloric acid solution (38). Hexachlo­

roiridate (III) has been shown by Poulsen and Garner (39) to

aquate very slowly. The visible spectra for iridium(III)

and iridium(IV) have been studied (39, 40) and a molar ab­

sorb tivity -constant of 4050 M cm has been given for

iridium(IV) at 488 nm (39). The molar absorbtivity constant

for iridium(III) is 92 M'- cm"-"" at 418 nm (40).

The substitutional inertness of iridium(III) and

iridium(IV) is maintained during reductions of iridium(IV).

For example, a recent study (41) has shown that the reduc­

tion of the potassium salt of hexachloroiridate(IV) by

l", N02~, SO3 ", S ~, S2O3 ", and CgHgOg (ascorbic acid)

takes place without any change in the inner coordination

sphere for iridium(IV) or iridium(III). The product of

the reduction of hexachloroiridate(IV) was always hexachlo­

roiridate (III) . This of course indicates that the reduction

does not involve the replacement of an inner-sphere chloride

by an ion of the reducing agent (41).

Another recent study (42) has investigated the reduc­

tion of hexachloroiridate(IV) by ascorbic acid in 80% acetone-

water mixture. According to the workers' results the

reaction shows a first order dependence on iridium(IV) and

on ascorbic acid with a second order inverse acidity depen­

dence. However, Gupta, et al. (42), propose the species

H2li^Clg in their mechanism shown in equations 8-11.

Wo _ , H2lrClg ^ HIrClg + H (8)

^6«806 t^ ^6^706" + «^

HIrClg- + CgH^Og- -1° HIrClg2- + -CgH^Og

(9)

(10)

HIrClg" + -C^H^O^ ^^^^ HlrCl^^ + C^H^O^ + H"*" (11) '6"7^6 '6 6 6

The species is not required by evidence obtained in this

thesis and is not experimentally found to exist at apprecia­

ble concentrations. The CI is a weak base and coordination

to metal ions makes bases even weaker (43, 44). Therefore

2-

IrClg would be a very weak base and this has been veri­

fied by an experiment which will be presented later in this

thesis. Other references describing mechanisms of the reduc­

tion of hexachloroiridate(IV) did not mention any protonated

iridium(IV) complex. The reducing agents studied are

cobalt(II) complex (45), platinum(II) (46), iron(II) complex

(47), and tetraphenylborate (48).

X

10

Copper(II) Catalysis of the Oxidation of Sulfur(IV)—

Copper(II) has been shown to have a catalytic effect on the

oxidation of sulfur(IV) (4, 16, 49-51). Higginson (4) in-

terprets the catalytic effect of copper(II) as due to the

higher reactivity of copper(II) over other oxidizing agents

in oxidizing sulfur(V) radicals. Carlyle and Zeck (16, 51)

determined that as the copper(II) chloride concentration is

varied there is no marked change in the reaction kinetics.

Zeck (51) also studied the reaction of copper(II) and

sulfur(IV) and proposed a mechanism shown in equations 12-15

which also agrees with the observations of Basset and

Henry (52).

Cu(II) + SO^^ t^^ [CUSO3] (12)

[CUSO3] -^-^^ Cu(I) + -SO^'

- 14 2-SO3 + -503 - S20g

H2O + •SO3" + Cu(II) - ^ Cud) + so^^" + 2H

(13)

(14)

(15)

The copper(II)-sulfur(IV) reaction was found to be slower

than the iron(III)-sulfur(IV) reaction (52).

X

CHAPTER II

EXPERIMENTAL SECTION

Reagents

Reagent grade sodium hexachloroiridate(IV) from either

Ventron Corporation or Apache Chemicals Inc. was used with­

out further purification. A spectrum of an aqueous solu­

tion of iridium(IV) had an absorbance maximum at 488 nm,

with a molar absorbancy coefficient of 4050 in agreement

with the results of C. S. Garner (39). After each new con­

tainer of sodium hexachloroiridate(IV) was opened, it was

stored at a temperature of -5°C as a precaution, although

this was not shown to be necessary. A fresh iridium(IV)

solution was used for each experiment, although acidic

iridium(IV) solutions were shown spectrophotometrically to

be stable at least seven days. Neutral solutions hydrolyze

after several hours (38).

All solutions were made up from doubly distilled water

The solutions were bubbled with nitrogen. Oxidants were

removed from the nitrogen by passing it through a scrubber

solution of chromium(II). The chromium(II) was maintained

in the scrubber solution by contact with zinc-mercury

amalgam.

The procedure for making sodium hexachloroiridate(III)

was adapted from the method of Poulsen and Garner (39).

11

X

12

Ethanolic solutions of hexachloroiridate(IV) (0.03F) and

sodium nitrite, J. H. Baker Reagent Grade (0.17F), were

heated separately on a steam bath until hot. No water was

added to the 95% ethanol in making the solutions. These

hot solutions were mixed giving a green precipitate of

Na3lrClg'2H O which was filtered and washed with warm 95%

ethanol. The ethanol was vaporized by pulling air through

the fritted filter. The dried precipitate was kept as a

dihydrate by storage in a desiccator, using anhydrous cal­

cium sulfate as the desiccant. Just before each use, the

iridium(III) precipitate was added to doubly distilled water

and analyzed spectrophotometrically (39). Since dissolved

oxygen interferes with the iridium(III) spectrum, the

spectrophotometer cell was purged with nitrogen. The con­

centration of the iridium(III) solution was calculated from

Beer's law using the value of the molar absorptivity at 418

nm as given by Chang and Garner (40). Because iridium(III)

hydrolyzes in water the unused solutions were frozen at

-10°C and concentrations remeasured spectrophotometrically

before new experiments were run. Previous work has shown

that the rate of hydrolysis is small (39).

Reagent grade sodium acetate from J. H. Baker Chemical

Co. was used for the sodium acetate solutions. The sodium

acetate was recrystallized twice from ethanol. The ethanol

was vaporized by pulling air across the sodium acetate

13

crystals while they were in the fritted glass filter. An

aqueous solution of the salt was then boiled for five

minutes and analyzed by cation exchange. A 6 inch by 0.5

inch column of Bio-Rad 50W-X8, 50-100 mesh resin was used

with doubly distilled water as the rinsing agent. An

aliquot of the salt solution was loaded on the Bio-Rad 50W

resin which was in the hydrogen ion form. The acid liber­

ated was determined by titration with standard sodium

hydroxide using phenolphthalein as the indicator. The con­

centration in the stock solution was 1.29 M.

The sodium chloride solution used was that remaining

from an earlier study (16) prepared by David W. Carlyle.

The reagent grade salt was recrystallized once and an

aliquot of the salt solution was loaded on a Dowex 50W resin

column v/hich was in the hydrogen ion form. The liberated

acid was titrated with standard sodium hydroxide using

phenolphthalein as an indicator. The stock solution was

4.55 M (4).

The hydrochloric acid stock solution was prepared by

dilution from reagent grade Matheson Scientific hydrochloric

acid solution. The solution was analyzed by titration using

standardized sodium hydroxide with phenolphthalein indicator

The stock solution was 1.0 87 M.

Sulfur(IV) solutions were prepared by dissolving

reagent grade anhydrous sulfur dioxide gas from a Matheson

14

Gas Products lecture bottle into oxygen-free doubly dis­

tilled water. These solutions were stored under nitrogen

in glass screw-necked bottles with self-sealing rubber caps.

Withdrawals of the solutions were made via syringes with

steel needles. The solutions were stored at 10°C when not

being used. The sulfur(IV) solutions were analyzed iodo-

metrically (53). Freshly prepared sodium thiosulfate solu­

tion was made in CO2 free water from reagent grade crystals.

This sodium thiosulfate solution was standardized from a

primary standard solution of potassium iodate and potassium

iodide. Addition of acid to this solution liberates iodine.

The indicator was freshly prepared starch solution. A known

volume of sulfur(IV) was removed from its bottle via cali­

brated syringe. It was added below the surface of the

potassium iodide and potassium iodate acidic solution

(iodine solution) and allowed to react. The remainder of

•the iodine was quickly titrated against the standardized

sodium thiosulfate solution using the starch solution as

indicator.

A solution of aqueous sodium sulfite (Na2S03) was also

made. An amount of dry sodium sulfite from Fisher Scientific

Company, A.C.S. Grade was placed in a screw-necked bottle.

The bottle was sealed with a rubber serum cap and purged

with deoxygenated nitrogen. In another screw-necked bottle

some standardized hydrochloric acid was sealed and purged

15

with nitrogen. The hydrochloric acid was transferred via

syringe to the dry sodium sulfite bottle. The sulfur(IV)

concentration was analyzed in the same manner that aqueous

sulfur dioxide was analyzed for sulfur(IV), as has been

described.

The copper(II) chloride solution used in the experi­

ments described was that remaining from work done by Otto

Zeck (51). The solution was made by dissolving double

recrystallized Baker and Adamson Reagent Grade copper(II)

chloride in doubly distilled water. The solution thus ob­

tained was analyzed in the manner of the analysis for

sodium chloride already described.

The copper(I) chloride solution was prepared in a

glass screw-necked rubber sealed bottle. Copper(II)

chloride crystals were dissolved in 6N hydrochloric acid.

Copper turnings from "Chore Girl," an all copper scouring

pad, were added. The solution turned from a very dark green

to an almost black color when the copper turnings were added

The solution was purged with deoxygenated nitrogen. It was

stirred continuously with a magnetic stirrer. The solution

gradually became clear with a fine white and green precipi­

tate settling to the bottom within approximately 24 hours.

Each day for seven days the solution was filtered, placed

in a screw-necked bottle with fresh copper turnings, re-

purged with nitrogen, and stirring was continued. After

16

seven days the solution was colorless with no precipitate

present. The concentration of the copper(I) chloride solu­

tion was measured by titrating it with cerium(IV); an aliquot

of copper(I) chloride solution was withdrawn from the screw-

necked bottle via syringe and added to 50 ml of 3N hydro­

chloric acid. The resulting solution was titrated with

standardized cerium(IV) until the solution changed from a

light red to a green color, using ferroin indicator.

An earlier study (16) provided the iron(II) solution,

which was prepared by David Carlyle by reducing iron(III)

perchlorate with amalgamated zinc. The iron(III) perchlorate

was prepared (54) by adding concentrated perchloric acid

(Baker Analyzed Reagent) to an aqueous solution of iron(III)

chloride (Baker and Adamson Reagent). The solution was

heated strongly to drive off hydrogen chloride gas. When

•the chloride ions were no longer found in this solution, as

shown by testing with silver nitrate, the solution was fil­

tered and recrystallized from perchloric acid and -then re­

crystallized again. The iron (HI) perchlorate was then

dissolved in water. The amalgamated zinc was prepared by

adding granular zinc (Baker Analyzed Reagent - 50 mesh) to

1 M hydrochloric acid in order to remove any oxide that may

have formed on the zinc. The clean zinc was washed with

water and then placed in water along with one drop of

mercury.

17

The detailed procedure for reducing the iron(III)

perchlorate was to add the iron(III) solution to the amalga­

mated zinc in a screw-necked bottle which was purged wi-th

deoxygenated nitrogen in a manner previously described. The

solution was stirred until iron(III) could no longer be de­

tected by testing with potassium thiocyanate. The iron(II)

thus obtained was analyzed by titration with cerium(IV) in

the presence of sulfuric acid and ferrion, using a procedure

described by Vogel (55).

Stopped-Flow Apparatus

The stopped-flow apparatus used in the rapid reactions

studied was similar to the design of Nakamura (56). He re­

ports circuit diagrams for his equipment (56) and the pro­

cedure and equipment is as follows. The solutions were

mixed by causing a variable speed-controlled, electric motor

(Model NSH-55, Bodine Electric Co.) with a magnetic slip

clutch to push the plungers of two separate reagent syringes

These syringes force the reagents into a four-jet stream

splitting into an eight-jet Teflon mixing chamber, and then

into a 3 mm. inside diameter quartz observation tube. This

mixing procedure allowed the separate solutions to mix and

enter the observation tube within 3-5 msec. Having each

solution at about -the same density by having them approxi­

mately at •the same ionic strength, provided the most

18

efficient mixing. The solutions in glass flasks were

sucked into the syringes by reversing "the electric motor.

The stopped-flow apparatus was equipped with a constant tem­

perature water bath which kept the solutions in their re­

spective flasks, the quartz observation tube, and the

stopped-flow syringes at a constant temperature of 25°C.

The syringe barrel was encased in hollow brass through which

water from the constant temperature water bath flowed to

maintain the correct temperature. The same applied to the

quartz observation tube whereas the solution flasks were

set directly into the water bath.

The changing transmittance of the solution in the

observation tube was observed spectrophotometrically at

488 nm. The light source was a tungsten lamp in a Beckman

Model DU quartz spectrometer. The lamp was operated by a

5-35 V, 7 amp Sola power supply feeding 6 v DC and 4.1 amp

DC to the lamp. An EMI 6256 B photomultiplier tube, biased

by a continuously variable 0-1000 v., 0-20 mamp., regulated

power supply (Kepco No. ABC 1000 M) monitored the intensity

of the light passing through the observation tube. For all

of the experiments described the monochromator slit was set

open at 2.0 mm. The normally high level of noise was re­

duced by -the combination of the wide slit opening and low

bias voltage.

19

The signal from the photomultiplier tube, via an ampli­

fying and smoothing circuit (56), provided the A input of a

Tektronix 564B storage oscilloscope (56) . A constant bal­

ancing potential (5 6) was connected to the B input of the

oscilloscope, and the difference (A-B) was amplified. The

balancing potential taken from the schematic diagram as

given by Nakamura (56) permitted amplification of changes

as small as 1 mv to the full vertical scale of the oscil­

loscope without amplifying the total signal of up to 6 v.

Turning a stopcock halted the flow of the mixed solution

causing the slip clutch on the electric motor drive to

activate. This clutch activation prevented the breakage

of the syringes and prevented any fresh solution from being

pushed into -the observation tube. As the stopcock closed,

a triggering circuit initiated a single sweep of the oscil­

loscope. A second sweep could be provided manually. A

Tektronix Camera Base Model C-12 using a Polaroid camera

photographed the oscilloscope traces.

The transmittance change is represented by the ordinate

of the recorded trace. Absorbance and transmittance are

very nearly directly proportional if the absorbance change

is limited to less than 0.1, and so the trace may be treated

as an ordinary concentration-time curve. This particular

oscilloscope was equipped with a variable adjustment time

scale wi-th sweep rates ranging from 1 x 10 to 5.0

20

_2 sec./cm. , but only sweep rates slower -than 5 x 10 sec./cm

were used.

Kinetics Experiments

The rate measurements were determined on either the

stopped-flow apparatus or the Beckman Acta V recording

spectrophotometer. During the reactions when sulfur(IV)

was in large excess the stopped-flow as used because of

•the rapid rate of reaction. When sulfur (IV) was in slight

excess, hence slower reaction rates, the Beckman Acta V

was used.

Rate measurements were initiated on the stopped-flow

by mixing iridium(IV) solution with sulfur(IV) solution.

The stopped-flow reservoir for iridium(IV) consisted of an

open-necked volumetric flask being constantly purged wi th

deoxygenated nitrogen. Hydrochloric acid to adjust -the so­

lution to the desired pH, the calculated amount of sodium

acetate solution (57) to buffer the solution, and sodium

chloride solution to maintain ionic strength of 0.2 had pre­

viously been added to make up the iridium solution. The

flask was placed in the constant temperature 25°C water bath

undernea-th the stopped-flow apparatus. It was connected via

a glass joint, glass tubing, and three-way stopcock for with­

drawal by the stopped-flow syringes. Nitrogen was bubbled

through the solutions for approximately 15 minutes prior to

21

initiation of the experiment to flush out •the oxygen.

Bubbling of the iridium(IV) solution was continued until

conclusion of the experiment. Just before the experiment

was initiated enough solid Na2lrClg was added to make the

-4

solution about 1 x 10 M. The sulfur(IV) reservoir con­

sisted of a small-necked, stoppered bottle (7). Glass tubes

•through the stopper enabled purging with nitrogen and addi­

tion of aqueous sulfur dioxide solution. Before adding the

aqueous sulfur dioxide, the reservoir was charged with

appropriate amounts of hydrochloric acid solution, sodium

chloride solution, and water. The bottle was placed in the

water bath and purged with deoxygenated nitrogen for approx­

imately 15 minutes. Bubbling was discontinued after the

aqueous sulfur dioxide was added via syringe and Teflon

syringe tube by plugging the gas outlet with a small glass

rod. However sufficient nitrogen pressure was maintained

so when the solution was withdrawn by the stopped-flow

syringe via glass tubing and three-way stopcock, nitrogen

would bubble into -the sulfur (IV) bottle to maintain a posi­

tive pressure. 'The total sulfur(IV) concentration in the

reservoir varied for each experiment and was calculated

from dilution from the stock solution.

The rate measurements on the stopped-flow were measured

at 488 nm. Four to six measurements were done with each set

of reactant solutions; within each such group of measurements,

the average deviation was typically about 7%.

22

The disappearance of iridium(IV) when sulfur(IV) was

in low concentration, and when iridium(IV) was in excess,

was monitored on the Beckman Acta V recording spectrophoto­

meter. The 10 cm. spectrophotometer quartz cells were

charged with the appropriate amounts of hydrochloric acid,

sodium acetate buffer solution, sodium chloride solution,

sodium hexachloroiridate(IV), and water. The cells were

then sealed with rubber serum caps and secured with copper

wire. Nitrogen was bubbled through the solution while the

cell and solution were being brought to constant tempera­

ture by immersion in a thermostated 25°C water bath. The

aqueous sulfur dioxide solution was added via syringe, -the

cell was shaken vigorously by hand, placed in the spectro­

photometer, and the absorbance recorded at 488 nm. The

cell was held in •the spectrophotometer by a specially-

designed hollow brass cell jacket positioned in the cell

compartment. The inside surface of the jacket was V-shaped

so the cell was in contact with the jacket along two lines.

Water from a constant temperature bath circulated through

the hollow brass cell jacket. The contact between the cell

and the brass jacket was sufficient to maintain temperature

uniformity for the duration of the experiments.

When checking for Cu(II) catalysis a known volume of

aqueous CUCI2 was added to the iridium(IV) solution pre­

viously described. An aqueous solution of iridium(III) was

23

also added to the iridium(IV) solution when checking for

retardation by iridium(III).

Stoichiometry

Stoichiometry experiments were run with iridium(IV) in

excess for one set of experiments and in another set of

experiments sulfur(IV) was in excess.

The iridium(IV) solutions at pH's of 4 and 5, buffered

with sodium acetate, with sufficient sodium chloride added

to maintain ionic strength at 0.2 were made up in 100 ml.

quantities. A portion was placed in a 10 cm. spectrophoto­

metric quartz cell, thermostated at 25°C, and purged with

nitrogen. The iridium(IV) solutions were then titrated with

standardized iron(II) solution by addition in increments of

0.05 ml of iron(II) via syringe. The disappearance of

iridium (IV) was monitored by •the Beckman Acta V recording

spectrophotometer. The change in absorbance was noted on

•the spectrophotometer chart paper after each addition of

iron(II) solution. Eventually there was very little or no

change in absorbance. The volume of iron(II) was plotted

on graph paper versus the absorbance values for the

iridium(IV) solution. Two lines were drawn tangential to

the points, and at the lines' intersection the amount of

iron(II) was read showing the end-point of the titration.

Ano-ther cell was -then filled with the now standardized

24

iridium(IV) solution, purged, and thermostated in the manner

previously stated. A known amount of sulfur(IV) was then

added via syringe and the reaction allowed to react to com­

pletion in the cell. The resulting solution contained un-

reacted iridium(IV). This iridium(IV) was analyzed by

spectrophotometric titration against iron(II) as just

described. The volume of iron(II) versus absorbance values

for iridium(IV) were plotted on graph paper. The end-point

was determined. The amount of iridium(IV) consumed in the

reaction with sulfur(IV) was calculated. The amount of

iridium(IV) remaining from the reaction with sulfur(IV) was

subtracted from the amount of iridium(IV) before reaction

with sulfur(IV). The amount of iridium(IV) consumed was

compared to the amount of sulfur(IV) consumed to obtain the

stoichiometry of the reaction.

Stoichiometry experiments were also run with sulfur(IV)

being in excess of iridium(IV). A known concentration of

aqueous iridium(IV) was placed in a 5 cm. spectrophotometer

cell with the respective amounts of hydrochloric acid,

sodium acetate solution, sodium chloride solution, and water

and capped wi-th a rubber serum cap. The solution was purged

with nitrogen, and the cell thermostated to 25°C in a water

ba th. An excess of aqueous sulfur dioxide was added via

syringe and the reaction allowed to go to completion. A

known amount of solution was then withdrawn via syringe

25

from the cell and analyzed for sulfur(IV) as already de­

scribed. The amount of sulfur(IV) remaining was subtracted

from the amount added to obtain the amount of sulfur con­

sumed in the reaction with iridium(IV). This was compared

with the amount of iridium(IV) reacted for the stoichiometry

of the reaction.

ESR

A Varian Associates Model V-4502 Electron Spin Resonance

Spectrometer was used for the possible detection of a radical

intermediate proposed for the iridium(IV)-sulfur(IV) reaction

This spectrometer system included the ESR Control Unit Model

V4500-10A, X-Band Klystron, Model V4250B Sweep Unit, Model

V-FR250 3 Field Regulated Magnetic Power Supply, Model V4560

100 Kc Field Modulation and Control Unit, Model V-4532 Dual

Sample Cavity, Model V-4270 Output Control Unit, Model

V-4260-B Power Supply Unit, Model G-22A Dual Channel Graphic

Recorder, and magnet all made by Varian Associates. The

system also included a Hewlett Packard Oscilloscope Model

120 AR for visual readout and tuning of the instrument.

Drawn quartz tubes with outside diameter of 0.2 to 0.4 mm

were used as sample cells. After the esr had been tuned

using a solution of Fremy's salt in both sample cavities

(g 2.005) (58), one of the tubes was replaced with a cell

containing a nitrogen purged sample of iridium(IV)-

-'rm

26

sulfur(IV) reaction mixture. An unsuccessful search was

made for a signal near the g value of 2.0030 for '303

radical (19) and near 2.0058 for the "302" radical (19).

Unsuccessful attempts were also made to observe the esr

signal for hexachloroiridate(IV) in aqueous solution.

Values for g were found in the literature for solid hexa-

chloroidate(IV) (59).

Equilibrium Studies.—To check the earlier work of

R. M. Golding (35) spectrophotometric runs of various con­

centrations of the analyzed sodium sulfite solutions were

investigated. Absorbances of the nitrogen purged, acidic,

constant ionic streng-th, sodium acetate buffered solutions

were measured on the Beckman Acta V at 255 nm, the absorp-2-

tion maximum for S2OC. . The absorbances were measured in

0.1 mm quartz cells.

CHAPTER III

RESULTS SECTION

Stoichiometry of Iridium(IV) -Sulfur(IV) Reaction

Prior work (4, 14, 15, 51, 52, 60) has shown that the

oxidations of sulfur(IV) solutions usually give the products

of sulfate or dithionate or a mixture of bo-th. Other work

by W. C. E. Higginson (18) has shown that both dithionate

and sulfate are formed in the oxidation by hexachloroiri­

date (IV), suggesting net reactions 16 and 17.

2IrClg^" + HSO3" + H2O ^ 2IrClg^" + SO^^" + 3H''" (16)

2IrClg^" + 2HS03~ -^ 2IrClg" " + S20g^' + 2H'^ (17)

Higginson's stoichiometry measurements were limited in

number, however, and were done within narrow concentration

ranges and did not include experiments with sulfur(IV) in

excess. Additional stoichiometry measurements were made in

this work in order to learn whe-ther the stoichiometry is

dependent on concentration conditions.

A summary of the series of stoichiometry experiments

is shown in Table 1. These results are in approximate

agreement with Higginson's figure of 1.87 (18) for [Ir(IV)]/

[S(IV)], although the ratios are consistently lower; within

the apparently large uncertainty, the stoichiometry does

not appear to be strongly affected by the concentration

27

28

variations listed. The method for the analysis of the

excess iridium(IV) and excess sulfur(IV) has already been

described.

Table 1.—Results of Stoichiometry Measurements Performed at 25**C and Ionic Strength 0.2, With 0.03 M Sodium Acetate Buffer

Expt. No.

1

2

3

4

5

6

7

8

9

ref .4

ref .4

a.

b.

1

1

1

1

1

1

1

1

1

[H+] Molar M Exces

X

X

X

X

X

X

X

X

X

10-5

10-"

10-4

10-"

10-"

10-5

Average

10-"

10-5

10-"

Average

3.98 X 10"^

2.95 X lO"^

Iridium(IV) is

Sulfur(IV) is

' Concentration of ;s Reagent Deficient Rea

a

a

a

a

a

a

of above:

b

b

b

of above:

.02M

.02M

[Ir(IV)]

[Ir(IV)]

the reagent in

the reagent in i

.003

.003

.009

.009

.009

.009

.0072

.0072

.0127

.02M

.02M

[Ir(IV)] igent [S(IV)]

[S(IV)]

[S(IV)]

excess, at 0.02M

excess, at : 0.05M

1.22

1.49

1.27

1.48

1.50

1.37

1.39

1.20

1.38

1.48

1.35

1.87

1.81

Kinetics of the Iridium(IV) -Sulfur(IV) Reaction

A summary of the data giving the experimental rate

measurements for the reaction of iridium(IV) and sulfur(IV)

29

at various acidities, buffer strengths, and ionic concentra­

tions is shown in Table 2. For the experiments with

iridium(IV) held constant at 1.3 x lO" M, ionic strength at

0.2M, sodium acetate buffer at 0.03M, and excess sulfur(IV)

varied from 5 x lO" -2.4 x lO" M, a pseudo first-order rate

constant could be found. The experiments listed in Table 2

except for experiments 11, 17, 21, 32, 44, and 46 conformed

to the following rate law.

I^ILII^) =^obsd.t^^(^^)] (18)

dt

The k , , rate constants were determined in the following

manner. Absorbance-Absorbanceop versus time could be plotted

from an oscilloscope trace photograph of the disappearance of

iridium(IV) with time. A typical trace photograph for experi­

ment number 16 is shown in Fig. 4. The absorbance points

versus time were plotted on semi-logarithmic graph paper. A

typical plot is shown in Fig. 5 for experiment 16 from Table

2. Since the plot is linear and is pseudo-first order then k , -, can be calculated from equation 19. obsd.

As evident from Table 3 the observed rate constant decreases

with increasing [H"^] at constant iridium(IV) and sulfur (IV)

— 3 + (6 X 10 M) concentrations. A plot of [H ] vs. k , ^ shown

ODSd.

in Fig. 6 gives a slope of -1, implying that k , , is

30

Table 2.—Rate Constants for Iridium(IV) and Sulfur(IV) Reaction at Ionic Strength 0.2M, Buffer Strength at 0.03M, and Iridium(IV) at 1.3 x 10"^M

E x p t . No .

1

2

3

4

5

6

7

8

9

10

1 1

12

1 3

14

15

16

17

18

19

20

2 1

22

23

24

25

26

27

1

3 . 9

4 . 6

5 . 3

7 . 5

8 . 7

8 . 9

9 . 1

9 . 6

9 . 6

1 .0

1 .2

1 . 4

1 .7

1 .7

1 .7

2 . 0

2 . 0

2 . 7

2 . 7

3 . 4

3 . 5

3 . 6

4 . 1

4 . 7

5 . 0

5 . 2

5 . 6

[H+] M

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

1 0 " ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - ^

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

1 0 - 5

[S (IV) ]

4 . 5

6 . 0

7 . 1

6 . 0

3 . 0

6 . 0

6 . 4

6 . 0

6 . 0

4 . 5

5 . 0

1 .25

4 . 5

1 .0

5 . 0

2 . 0

1 .0

6 . 0

6 . 0

4 . 5

9 . 8

2 . 1 1

6 . 0

6 . 0

2 . 0

1 . 1

1 .0

M

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

1 0 " ^

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 4

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 4

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 2

1 0 - 3

1 0 - 3

1 0 - 3

1 0 - 2

1 0 - 3

o b s d .

1 7 . 1

1 7 . 5

1 9 . 7

1 2 . 4

5 . 1

1 1 . 2

9 . 4

1 0 . 1

1 0 . 1

6 . 6

0 . 6 5 ^

1 .5

3 . 6

0 . 9 0

0 . 4 0

1 .3

0 . 6 7

3 . 5 ^

3 . 2

1 .9

2 . 6 ^

1 4 . 2

2 . 3

2 . 0

0 . 5 7

6 . 2

0 . 2 2

c a l c d .

1 6 . 8

1 9 . 0

1 9 . 7

1 1 . 8

5 . 1

9 . 9

1 0 . 4

9 . 2

9 . 2

6 . 6

0 . 6 0

1 .3

3 . 9

0 . 8 7

0 . 4 3

1 .5

0 . 7 4

3 . 3

3 . 3

1.9

4 . 1

8 . 5

2 . 1

1 .9

0 . 5 9

3 . 1

0 . 2 6

31

TABLE 2—Continued

E^Pt. [H+] [S(IV)] k ^ _ k^^^^^

No. M M

28

29

30

31

32

33

34

35

36

37

38

39

40

41

42

43

44

45

46

5.8

5.8

6.8

6.8

6.8

6.8

6.8

6.8

6.8

6.8

6.8

7.1

8.6

8.6

2.0

1.0

1.0

1.0

2.0

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-5

10-4

10-3

10-3

10-3

10-2

4.0

4.0

2.4

1.1

6.0

6.0

6.0

6.0

2.0

1.0

5.0

3.0

2.0

1.0

6.0

1.0

1.0

5.0

6.0

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

X

10 3

10-3

10-2

10-2

10-3

10-3

10-3

10-3

10-3

10-3

10-4

10-3

10-3

10-3

10-3

10-3

10-3

10-4

10-3

1.0

0.92

7.0

2.9

1.1^

1.6

1.4

1.5

0.43

0.19

0.11

0.60

0.38

0.15

0.48

0.035

0.017^

0.008

0.189®

1.0

1.0

5.2

2.3

1.3

1.3

1.3

1.3

0.43

0.22

0.11

0.62

0.34

0.17

0.44

0.015

0.015

0.007

0.003

^lonic strength 0.1.

^Sodium acetate is 0.015M, sodium chloride is 0.185M.

^Sodium acetate is 0.06M, sodium chloride is 0.14M.

Oxygen gas present in solution.

®The k , , takes into account -the Q„o^ ~ = -043 (32). calca. nbU-

32

mm,fii^ ^ L?*^*!^ij^: ^J&uaatmM

•\ X riB«ii«««!!'-^•B;^«r?.'-;5'g/j 9ir«p> innamaisr-imiiVr-Tiis-

i r • 1

V i.«.: '.jH'-i'm ':S»-:«»BlMa^.

. i i i i ' ' l a a s i i l l

* "*• { - ; •'? .<» J | »

S'^f--'

I iii s -: •:• :•! a t ?• - i • •'. H B B .--"K.-- "•tfaK,-.,,;aw«i.-

= a a '

N 1

, ; , 1 ••

3 1~ .!:•'•.»•'•*m

ian

- •sst:i«amV'^. " -" '.-.'vjnnKitufc

I-;"!'*!'••HIM ism •! ••• •iil'Jf"''! -; 1

"• ' • " 3~".:;s •5'3>«.'*«3a2r.T; ,;':iai!; • - ;:

V 'qr r s s i s s s a R i B e a e H B * . : : ^ ; ^ : - - -1 : 'iaisum^mmvmsmmunHmnnnmm

1 iBiaaaMBBBeBBHUIiBliiiiiaB

porj-.ij • t : : '

^^^-SBKBHIIBiiiiil § ^ l ! ^ ! ' 3 B " i i ( > l 1 H ^ ^ ^ ^ ^ « i s a B * " B B i i B a B n « a H i i i i M B

tan Jti

i *»•:?-j '_a«Ba; r.'3ii!;%aEss»af i i i»i;M3H£;jsai i4;«;'«iiBana3B( iMMti-siisaaBaL'

r ^ ^ a a B e a a l 1 v;5r :-><i«ic«aaB! '• ^ >-• S ' T W « >f O « >/'

'<^4fi£3B3aaaj^;i

T!m><sae!aBasr

B!Sa^B93CS«! •iafiaBRssBfliFs! e i s s a a c a s a : ] ' ^ s:sa = S B 3 C 9 a N : : : : s s s 9 s a a t : : : s = = s a s 3 a ! • i i i a a f i a i i k k • • • • • • • - " - ^ -

Fig. 4.—A typical photograph of an oscilloscope trace of •the disappearance of iridium(IV) . This particular photo­graph is from experiment 16 of Table 2.

2 4

time (seconds)

Fig. 5.—A typical plot of Absorbance-Absorbance^ versus reaction time using the conditions given in Table for experiment 16.

33

inversely proportional to [H ]. A plot of k wg^ [H ]

versus sulfur(IV), shown in Fig. 7, gives a non-linear least

squares computer (61) calculated slope of 0.0147 +_ .0004.

The above information gives

^ _ .0147 [HSO,-] ^obsd. - ^ (20)

[H+]

which gives the complete rate law in equation 21.

-d[Ir(IV)] = 0.0147 [Ir (IV) ] [S(IV)] (21) dt [H" ]

Table 3.—Acidity and k j g at 6.0 x 10-3M Sulfur(IV) 1.3

X 10-4M Iridium(IV), 0.03M Buffer, 0.2M Ionic Strength, and 25°C

M

2 4.63 X 10-^ 17.5

4 7.5 X IQ-^ 12.4

6 8.9 X 10-^ 11.2

8 9.6 X IQ-^ 10.1

19 2.7 X 10-^ 3.2

23 4.12 X 10-^ 2.3

24 4.7 X 10-^ 2.0

34 6.8 X 10-^ 1.4

42 2.0 X 10-4 0.48

46 2.0 X 10-^ 0.19

Various series of experiments were run to test the

effect of some changes in the standard conditions.

^

34

o •

o o o

^ o i H

O •

o o o .H

S VO 1 o H

X

1—1

+

o •

o o f H

O •

o H

ffi "—'

c •rH

C <u >

•H tr>

(0 C 0

•H

ta

and

co

nd

it

(d -d

0) ^ + j

^ -P •H >

•

CD XI 0

M

W :3 to J <u >

>1 -p •H TJ •H O td

M-l 0

4-> 0 H Cl< 1

Fig

. 6

.-a

ble

3

.

EH

35

o 0) CO

vo I o

X

ffi

CO

O

500.0

100.0

10 .0

L I I I M 1 1 I I M I

» I I i I I t I t I I i

0.001 0 .01 [S(IV)]

F i g . 7 . — P l o t of [S(IV)] v e r s u s k

T—rn

J L

o b s d . [H-^] with iridium(IV) 9 x IO'^M, 0.2 ionic strength, 0.03M sodium acetate buffer, 25°C, and the data points taken from Table 2.

36

Comparing -the k t , expression in Table 2 for experi­

ments 11 and 21 where the total ionic strength is 0.1

instead of 0.2 to the calculated k , , gives 0.60 and 4.1

respectively.

In Table 2 experiment 46 shows a large difference in

^obsd ^^^ ^calcd * " ^ reason is unclear but may be due

to a different mechanistic step at such a high acidity.

When the sodium acetate is reduced to half of its normal

concentration of 0.0 3 M the observed rate constant increases

as shown in Table 2, experiment 18. However, when the sodium

acetate concentration is doubled the observed rate constant

decreases as shown in Table 2, experiment 32. The changes

in the rate are quite small especially when compared to cal­

culated values. Therefore the effect is small, although the

data is not inconsistent with a small effect. The procedure

for measuring the kinetic experiments has already been de­

scribed, but when no nitrogen purges the solution, as seen

in Table 2 experiments 43 and 44, the observed rate constant

decreases.

A series of experiments were run to measure the rate

constant when iridium(IV) was in excess. These observed

rate constants along with the respective hydrogen ion and

iridium(IV) concentrations are shown in Table 4. In this

particular series of experiments, •the decreasing absorbance

of iridium(IV) was measured. However, the absorbance

37

Table 4.—Excess Iridium(IV) and Sulfur(IV) Observed Rate Constants at 0.2M Ionic Strength and 0.03M Sodium Acetate

Expt.

No.

1

2

3

4

5

6

»

1.95

1.95

6.80

6.80

1.55

,1.55

[H+]

M

X

X

X

X

X

X

^[S(IV)]

10-5

10-5

10-5

10-5

10-4

10-4

small

See text.

[Ir(IV)]

M

5.46 X

1.61 X

4.22 X

2.37 X

6.25 X

1.67 X

in this

10-5

10-4

10-5

10-4

10-5

10-4

Study

obsd.

0.017^

0.052^

0.005^

0.045®

0.005^

0.013^

Average

•

calcd.

0.041

0.121

0.009

0.051

0.006

0.016

calcd.

^obsd.

2.49

2.33

1.98

1.14

1.20

1.21

1.73

Waveleng-th 488 nm.

wavelength 520 nm.

Wavelength 525 nm.

decreases were proportional to decreases in sulfur(IV) con­

centration, so that •the pseudo-first order rate constant

also described the disappearance of sulfur(IV) as expected

from equation 21. Because the sulfur(IV) and iridium(IV)

are not used in a 1:1 ratio, k^j^^^ for the disappearance

of sulfur (IV) is not equal to •the k , . for the disappear­

ance of iridium(IV). As shown in Fig. 8 plots of k^j^^,

versus iridium(IV) concentrations give straight lines with

slope 1 also as expected from equation 21. These plots fit

the pseudo-first order rate law in equation 22. If the

50.0-

I

o (U CO

CO I o l-l X 10.0—

CO

O 5.0 -

[Ir(IV) ] X 10-^M

Fig. 8.—Plots of k obsd. versus excess iridium(IV) concentration using the data presented in Table 4. O, n, and A represent 6.80 x 10"5, 1.95 X 10"5, and 1.55 x lO'^M H+, respectively.

39

same rate law is assumed to hold in excess iridium(IV) as

-d S(IV) = k t)s i, [Ir(IV)] [S(IV)] (22) dt

holds in excess sulfur(IV) the differences in the k , , obsd.

values can be used to estimate the stoichiometric ratio.

As can be seen from Table 4 the observed rate constants are

each smaller than the value calculated from equation 21.

This difference is in agreement with the stoichiometry of

the reaction, which is 1.35 - 1.87 Iridium(IV) for every

sulfur(IV) reacted. The average of the column of k^^T^j / caxcQ.

k , J is 1.73 in compliance wi-th both the earlier sto-ODSd. ^

ichiometry values given in Table 1 and Higginson's (18)

value. The large range in values of k , , /^obsd ^^ ^^®

in part to -the fact •that iridium(IV) had to be in relatively

high concentration. At •these concentrations •the sulfur (IV)

concentration was very low and any more dilution would have

compounded the error. But because of the high iridium(IV)

concentration and the small change in absorbance there was

a high pen noise factor which contributes to -the wide range

in values.

Iridium(III) Retardation

The possible effect for iridium(III) retardation of the

reaction between iridium(IV) and sulfur(IV) was investigated.

A series of experiments measuring •the effects of iridium(III)

are shown in Table 5. The iridium(IV) was held constant at

(D Xi -P

U O

cd • P CO c o u 0

+J (d

(D > U O CO

o o -P O

• H C -P o o

fd ^ 0 H Pi H H ^--- - > g H •H }

•H ^ M .H H :3

W m I o fd -p > O H

M-l g M-l ^ W -rH

a -H

H H

CO JQ O

I

X

H

H H H S

LO I o

X S

W

40

Xi

00

• in

•^ o r~-

CN ro rr

(N

CO

vo in CO

'^ • « ^

CO

00 t

CN

VO '^

CO

00 r»

cvj

m ^

c^

n H

CN]

•<;f r>-

rH

CO ^

CM

00 nH

CN

»;J« rj< ^ 'i^f T3< T^r

CN CN CN CN OJ CNJ

in 1 o rH

X

o • in

•^ 1 o iH

X m CN •

• > ! l '

1 o iH

X LO •

CN

'^ 1 o iH

X

o • in

ro 1 o rH

X

o • H

n 1 o fH

X

a\ •

fH

00 00 00 00 00 00

vo vo vo vo vo vo

0) :3 fH

(d >

Q) tP (d iH <D > rt^

in o H

in o iH

in o

H

in o rH

in o iH

0) :3 fH fd >

Q) en fd u 0) > <

ID •^ T^ en CO I I I I I o o o o o

X X X X X o in o o o • CN • • •

in • in H 00

00 00 00 00 00

vo vo vo vo vo

in <U fH

EH

O

•P

X fH CN n - in vo 00 <J o fH

41

Xt CO

o

0) 0

-H 4J C O U I I

in

W

9 EH

CN I O

X

> H

H H H S

H

in I o

X S

W

o

4J

X

vo CN

•^ ''"''

CN VO

in

r* • ^

f H in

• ^

00 CN

"<;r

00 r-"^

in in LO lo lo o o o o o

O rH fd >

tn fd u (D

lO ' ^ ' ^ I I I o o iH fH

CO CO I I

o o o

X X

o

lO

i n CN

X X

o o . •

i n H

iH

X

CN CO

i n

T^ ';r ^ ^ ^ vo vo vo vo vo

ro m CO CO ro

CN CO in vo

in I o fH X •

fH i n c i

00

+) fd

+J

fd •P CO

O o tH a) x:

>

H

-H

o •H +J fd

g o u M-l

CO <D

iH fd >

a) +J (d

H :3

TJ O •H fH M fd H U

fd Xi

X f<M:-.-ti igfgg.

42 -5

8.9 X 10 M with ionic strength of 0.2M and buffer strength

of 0.03M. As can be seen in Table 5 the increasing of the

iridium(III) concentration causes a random scattering of

values for the observed rate constant. Therefore it is

concluded that iridium(III) has not more than a small

effect on "the proposed rate law.

Copper(II) Effect

A series of experiments shown in Table 6 were done to

investigate the possible effects of copper(II) catalysis of

the reaction between iridium(IV) and sulfur(IV). The pro­

cedure for these experiments has previously been described

in the experimental section. The extent of the catalysis

was difficult to determine because of the sudden break of

the oscilloscope trace (see Fig. 9 and compare with Fig. 4

where no copper (II) was present). As shown in Table 6 the

break becomes more pronounced as the copper(II) concentra--4

tion increases. A reaction between copper(I) (1 x 10 M)

and iridium(IV) (4 x 10- M) under the same conditions of the

iridium(IV)-sulfur(IV) reaction was too fast for observation

by stopped-flow. Assuming that 12.5% of the reaction could

have been observed a rate constant could be calculated,

having a lower limit where k >: 1.0 x 10 M sec .

Iron(II) Effect

An attempt was made to study the possible effects of

iron(II) on the iridium(IV)-sulfur(IV) reaction; however.

X

43

-3, an experiment showed that an iron(II) (5 x 10 M) reaction

-5 with iridium(IV) (9 x 10 M) was too fast for stopped-flow.

Again assuming only a small portion of the reaction could

be observed (12.5%) then the lower limit of the rate con-

4 -1 -1 stant could be calculated where k > 8.3 x 10 M sec

iif^ipiLijiijfw «iik"ij!fy rm •mm-.tm-.immi

" • " ^ 3 « « « " • : ?.1l!S5f

' :. .-^-SBaj

iX^BSSiaBIBB&r ifTBBUiaaaa Hu Si 8) w'«t , j£

-;;\a«iu<iaa< iaiajiiBSiii2;£i!aHiiiaaa«i'aaJiidifti««^&if'-'JKiuuuiBssaHHBuaaHaiiHaiiaaaaBflir

!liS£id

Fig. 9.—A typical oscilloscope trace for experiment no. 1 from Table 6 for the copper(II) catalyzed iridium(IV)-sulfur(IV) reaction.

Radical Investigation

A series of experiments were run to detect the presence

of a radical in the iridium(IV)-sulfur(IV) reaction. As

shown in Table 7 various compounds were allowed to react

with aqueous sulfur dioxide in the presence of acrylic acid,

a known radical scavenger which is soluble in water. The

compounds used were at approximately the same concentration

44

as the iridium(IV) concentration in •the kinetic experiments.

Aqueous sulfur dioxide was added to the nitrogen purged

solutions. The qualitative observations which were made are

recorded in Table 7. The possibility of radical formation

is evident from the acrylate polymer believed to be formed

in •those qualitative reactions when a milky turbid substance

was visible. There were no attempts made to characterize

the polymers formed.

Table 6.—Effect of Copper(II) Catalysis on the Iridium(IV)-Sulfur(IV) Reaction

Expt.

No.

1

2

3

4

H-

X 10"

6.8

6.8

6.8

6.8

• ^ M

[Cu(II)]

M X 10-3M

3.5

7.0

14.0

21.0

[S(IV)]

X 10"3M

6.0

6.0

6.0

6.0

Results^

Mild break of curve

Mild break of curve

Pronounced break of curve

More pronounced break of curve

i

See text.

Search For Radical via ESR.—Attempts were made by •the

procedure described earlier in the experimental section to

examine the iridium(IV)-sulfur(IV) reaction for a radical

intermediate through detection by esr of any radical formed

Scanning in •the expected region of the radical showed no de­

tectable signal. Because of the rapid reaction rates and

45

extremely low concentration of any reactive radical inter­

mediate, a signal may have been present, but so weak that

it was hidden in pen noise. (Increasing the reaction rate

would increase the amount of radical proposed present. A

flow system using this factor has been built by R. 0. C.

Norman (62) for the detection of radicals produced in fast

reactions.) Another worker is currently investigating for

•the presence of radicals in sulfur (IV) systems.

Table 7.—Reaction of Sulfur(IV) with Various Compounds In the Presence of Aqueous Acrylic Acid

Compound Results

Iodine

Hydrogen Peroxide

Cerium(IV)

Iron(III)

Tris (orthophenanthraline) iron(III)

Copper(II)

Iridium(IV)

The yellow color of iodine dis­appeared immediately upon addition of sulfur(IV) but no polymer formed.

There were no visible results.

Upon addition of sulfur(IV) a white milky cloud formed.

A white milky cloud formed a few seconds after addition of sulfur(IV).

The iron(III) was oxidized, but no evidence of a polymer appeared.

A white cloud formed.

A pale yellow milky cloud formed a few seconds after addition of sulfur(IV).

Sulfur(IV) There were no visible results

46

Equilibrium Study

The equilibrium constant for the following reaction

shown in equation 23 was briefly studied to determine whe ther 2-

^2^5 ^^ ^ prominent species in the iridium(IV)-sulfur(IV)

reaction.

^23 ^ ^ 2-2HSO2 ^ - 2^5 ••" ^2^ ^^3)

A series of experiments described in Table 8 give the cor-

2-rected absorbance values for S^O^ at •their respective

HSO^ concentrations. The absorbance values are plotted

versus concentration of sulfur(IV) in Fig. 10. The slope of

the line drawn in Fig. 10 is equal to the quantities in the

following equation,

slope = Q e 1 (24)

where Q is the equilibrium constant for reaction 23, e is 2-

the molar absorptivity for S20^ , and 1 is the cell path

length (0.1 mm). From the slope of the line, the value

Ke-__ = 560M- was determined which is not in good agreement 255

with the previous value, which was 280M (21, 35). Taking

•the point for experiment 7-A in Table 8, the value of Q may

be estimated by taking into consideration the difference of

its X-axis distance away from the line drawn in Fig. 10.

The calculated value of Q was found to be 0.05 which agrees

wi-th Golding (35) . However, the calculation for Q is approxi­

mate and is based on one line and one point. An upper limit

47

Table 8.—List of Absorbances for Various Concentrations of Sodium Sulfite at 1.71 x 10~5M [H+] and 0.06M Sodium Acetate

Expt. No

2-A

4-A

6-A

7-A

8-A

[HSO3 ] initial M

Abs.

.2

.3

.4

.7

.4'

Abs. corr.

0.229

0 .543

0 .938

2 . 4 1

0 .936

0 .219

0 . 5 2 8

0 .918

2 . 3 8

0 .916

.12M Sodium Acetate.

of Q < 0.11 was calculated, using the reasonable assumption

that -the error was no more than 5%. The method used in this

calculation is given in the following calculation.

2-The actual amount of S2OC is

(.49) 1/2 _ (.43) 1/2 _ = .700 - .656 = .044 .022 2 2 2

where (.49) '' is the greatest formal bisulfite concentra-

1/2 tion used, giving an absorbance of 2.4, and where (.43)

is the value that would have been sufficient bisulfite con­

centration had the points been on a straight line.

Since Q = [S2O5 "]

[HS03-]^

2- - 2 t h e n s u b s t i t u t i n g i n t h e v a l u e s f o r [32©^ ] and [HSO^ ]

g i v e s Q = 0 .022 = 0 . 0 5 0 , 4 3

^ X

48

5.0

-P o a) u u o o 1.0

(D O c (d

- 0.5 o CO

T 1 I M I I I 1 1 I I I I J

I I I I I I I J ' I I I t i 0.05 0.1 0.5

[HSO3]

Fig. 10.—Plot of Absorbance (corrected) versus total sulfur(IV) concentration squared taken from the data presented in Table 8.

49

The rate law for the iridium(IV)-sulfur(IV) reaction

(found as equation 21) shows that the reaction of iridium(IV) 2-

with SyOc is not important. The estimated Q ^ 0.11 shows

that the HSO^ concentration is not significantly reduced

under the conditions of the experiments investigated in 2-

•this study by the formation of S2OC . Again these calcula­

tions are only approximations. The deviation from Golding's

calculation (35) could be due to a medium effect even though

experiments 6-A and 8-A seem to refute any medium effect

vida infra.

Examination of the data for experiments 6-A and 8-A

from Table 8 is evidence that sulfur(IV) does not react with

the sodium acetate buffer. This is suggested by •the absor­

bance value for experiment 6-A, which was 0.4M sodium

sulfite with 0.06M sodium acetate buffer, being equal to the

absorbance value for experiment 8-A, which was 0.4M sodium

sulfite with 0.12M sodium acetate buffer.

CHAPTER IV

INTERPRETATION SECTION

Reduction of Hexachloroiridate(IV) by Sulfur(IV)

The features of the empirical rate law given in equation

21, Chapter III for the uncatalyzed reaction of hexachloro­

iridate (IV) in excess sulfur(IV) are consistent with the

following mechanisms: A (reactions 25-28), B (reactions 29-

32), and C (reactions 33-36).

Mechanism A

HSO3- Z"^^ ^^3^" + ^ ^ ^ ^

IrClg^- + SO^^- - ^ IrClg3" + .30^' (26)

IrClg^- + -303- ^^^ IrClg3" + S03(S04^-) (27)

2 ^503- ^^^ S20g^- (28)

Mechanism B

HSOr - ^ SO^^- + H^ (29) 3 H- i

2- 2- " 0 4-

IrClg" + SO3 - ^ [IrClg-S03]^ ^30)

[IrClg-S03]4~ + IrClg^- ^^|^ 2IrClg3" + S03(S04^-) (31)

2[IrClg-S03]4" ^^^^ 2IrClg3- + S20g^- (32)

Q33 „^ „, - . „+ Mechanism C

H2lrClg z'^'^ HIrClg -f H' (33)

50

51

HIrClg" + HSO3- - 34 HIrClg^" + '303- + H" (34)

HIrClg- + -303- 5^^ HIrClg^- + 303(30^^-) (35)

2 .303" 5^^ S20g^- (36)

Though •the empirical rate law provides no information re­

garding the fast steps after the rate determining steps,

mechanisms A, B, and C are some reasonable possibilities

-that are consistent with the stoichiometry. The complex

shown in equation 30 for mechanism B could have an oxygen

or sulfur linkage to the iridium species. Also for this

complex the electron could be transferred across to the

iridium(IV) and the radical could remain attached ( [IrClg-

•SO-.] ) . This radical complex could then dissociate giving

the products found in equation 26 which would react further,

equations 27 and 28. The proposed radical complex could

also react giving the products in equations 31 and 32. The

4-formation of [IrClg-'S03] will be called mechanism B'

henceforth in this thesis. Mechanism C equation 33 shows

the dissociation of H2lrClg. This equation could also be

the dissociation of HIrClg which will be called mechanism C

Mechanism C fits the rate law but is believed to be of

a minor significance because the reacting species is be-

2- -lieved to be SO3 and not H3O3 . The species H2lrClg has

been proposed (42) recently in a mechanism involving hexa­

chloroiridate (IV) and ascorbic acid, as has been previously

52

discussed. Mechanism C , using the proposed HIrClg- disso­

ciation, would fit the rate law; however, the arguments that

have been previously described and with the experiment de­

scribed in the following sentences, it is the contention that

neither H2lrClg nor HIrClg exists in any appreciable amount

under -the conditions used in this work. This was verified

by adding to an aqueous hydrochloric acid solution at pH =

4.15 enough solid hexachloroiridate(IV) to make a calculated -3

1 X 10 M solution. The pH remained at 4.15 as read on a pH

meter. More iridium(IV) was added to make about a IM solu­

tion and the pH remained at 4.15. The proposed basicity of

2- 2-

IrClg (42) would have caused the pH to increase as IrClg

protonated, which experimentally did not happen. Therefore

it is believed that mechanism C and C could be deleted be­

cause H2lrClg and HIrClg are not believed to exist at sig­

nificant concentrations.

Mechanisms A and B' are free radical mechanisms whereas

mechanism B is not. The difference between mechanism A and

B' is whether the product of the first electron transfer is 4-

a -SO-j- radical (A) or a [IrClg--S03] radical (B'). Both

mechanisms are in accord with both the stoichiometric and

kinetic observations and so both could be credible mecha­

nisms. To differentiate between mechanisms A or B' and B

the complex [IrClg-S03]4- or the free radical -303 must

be shown to be present. The complex was not detected

y^

53

experimentally. However the complex may not absorb in the

visible region or the concentration may have been too low

for detection.

As has been previously mentioned the free radical was

not detected by esr. Again it is possible that a free radi­

cal was present in small concentrations with an upper limit — 6

of about 10 M, or the lifetime was very short.

As has already been mentioned sulfur(V) radicals can

exist in aqueous solutions (20-24) . There are several possi­

ble pathways for further reaction of the sulfur(V) radical.

The sulfur(V) species could react with oxygen, but this is

a negligible reaction since all solutions were deoxygenated

by nitrogen. The radical might react with -the solvent but

it has been shown that the decay of sulfur (V) radicals by

reacting with water is insignificant (21, 22). It might

disproportionate or dimerize. Disproportionation of the

sulfur(V) species would give sulfate and sulfur dioxide.

However, present evidence by Hayon, et al. (21) suggests

•that dimerization of sulfur (V) radicals is the most impor-o

tant decay route. A rate constant of k = 5.55 x 10

M" sec- (21) has been measured for the dimerization of

sulfur(V) radicals. Of course the radical could also react

with another iridium(IV) species. Both dimerization and

reaction with iridium(IV) are proposed in mechanism A.

trn 4\r:^iiU •;

54

As already noted in Chapter III toge-ther with -the data

given in Table 5, the observations point to -the fact that

iridium(III) has at best a very small effect on the rate

constant for the iridium(IV)-sulfur(IV) reaction. This

small effect observed could be medium effects only. The

lack of an inhibition effect shows that no step in the mech­

anisms with iridium(III) as a product is appreciably revers­

ible. As noted in Table 7, Chapter III the results for the

scavenging of any possible free radical are given. Acrylic

acid has the properties of both being soluble in water and

a good radical scavenger. Iridium(IV) reacting with

sulfur(IV) in aqueous acrylic acid under the same conditions

used for the kinetic runs formed a pale yellow milky cloud.

This cloud formed into a gelatinous ball believed to be an

acrylate polymer. The acrylic acid solution was shown not

to react with either iridium(IV) or sulfur(IV) alone, there­

fore no acrylate polymer could have resulted from these

sources. Several workers (1, 4, 14-16) have postulated that

-the hexaaquoirion (III) reaction with sulfur (IV) takes place

via radical mechanism. The experimental results in Table 7

also showed the presence of a polymer when iron(III) reacted

with sulfur(IV). This evidence strongly suggests, although

is not conclusive by any means, that mechanism A, a radical

mechanism, plays an important role in -the iridium(IV)-

4-sulfur(IV) reaction. The complex [IrClg-S03J seems to be

53

unlikely to cause the polymerization of acrylic acid so

mechanism B is believed to be an unimportant mechanism for

the iridium(IV)-sulfur(IV) reaction.

D. W. Carlyle (1) proposed that tris(l,10-phenanthro-2_

line) iron (III) might react with the 3^0^ species because

he found a second order rate constant for the bisulfite ion

concentration in his rate law. An alternate proposal of

Carlyle (1) was that in a complex mechanism the two bisulfite 2-

ions arrive separately, rather than together as the 320^

species, and that an iron-bisulfite ion pair is •the inter­

mediate which reacts wi th the second bisulfite ion. One of

the objectives of this •thesis was to determine whether the

pyrosulfite ion was a dominant species. As shown in the

rate law given in equation 21, Chapter III, this work shows

a first order dependence for sulfur(IV) which is inconsis-2-

tent with the 3^0^ species. As previously mentioned in

•this •thesis, -the sulfite concentration is not significantly

reduced by the formation of pyrosulfite under the conditions

investigated in this work. The most prominent species under

the conditions of this study is the bisulfite ion.

It has already been noted that when oxygen is present

in the reaction mixture of iridium(IV) and sulfur(IV) that

the rate of disappearance of iridium(IV) with respect to

time decreases. This could be explained by the fact -that

sulfur(V) radicals may have been scavenged by the oxygen

56

present. As shown in the following equations, flash photo­

lysis of sulfite ions in the presence of oxygen forms radi­

cals which react very fast with oxygen to form sulfate (21).

SO3 hv -303 + e _ (37)

•SO3 +02"^ 'SO^- (38)

•SOg- + HSO3- -> H3O5- + -303- (39)

HSO^- + 303^- -> HSO^- -f 30^^- (40)

Since '30^ would have the tendency to be reduced rather than

oxidized, then 'SO^ probably would not react with iridium(IV)

The above mechanisms does show a competition for the pro­

posed '303 radical. Therefore the retardation by oxygen

is possible evidence for the presence of radicals though not

conclusive evidence.

Even though the presence of any sulfur(V) radical has

not been conclusively proven; the possibility of the species

being a free radical, '303 , or being bonded to the metal

4-ion, [IrClg-*303] is worth some discussion. As has been

mentioned, some workers have postulated the presence of a

free radical (1, 4, 8, 14-16) while others (7, 10) have the­

orized the metal-sulfite radical complex. Higginson and

Brown (18) together with Novoselov and Muzykantova (41),

and this work suggests that with a fairly strong oxidizing

agent which is substitutionally inert, the preferred route

of the reduction of sulfur(IV) by iridium(IV) is believed

57

to be the direct formation of the free radical by an outer

sphere mechanism. Higginson and Brown's (18) small amount

of stoichiometric data for the reaction of hydrazine with

hexachloroiridate(IV) and sulfur(IV) with hexachloroiri­

date (IV) indicates the likelihood of an outer-sphere mecha­

nism. They suggested that the nonstoichiometric oxidation

of sulfite completely to sulfate and hydrazine to nitrogen

is a consequence of an alternate path of reaction for the

free radicals produced (18). This could possibly be the

proposed dimerization as suggested in equation 28. As has

been previously mentioned, Novoselov and Muzykantova (41)

indicate that hexachloroiridate(IV) is substitutionally

inert during reduction by sulfite. A more complete set of

stoichiometric work together with the kinetic data presented

in this -thesis substantiates Higginson and Brown and

Novoselov and Muzykantova. The attachment of the sulfur(IV)

to the iridium(IV) via a chlorine bridge must not be dis­

counted as a means of initial electron transfer. Since

4-neither '30^ nor [IrClg-.SO-] has been concretely iden­tified, the possibility of either being present is evident.

Copper(II) Catalysis of the Reduction of Hexachloroiridate(IV) by Sulfur(IV)

Al-though only a few data were collected in this work

on the copper(II) catalyzed iridium(IV)-sulfur(IV) reaction,

a mechanism is postulated using the data found in Table 6,

y^

58

Chapter III and other observations of the oxidation of

sulfur(IV) with copper(II) as a catalyst (4, 16, 49, 50).

Mechanism D is shown in equations 37-42.

Ir(IV) + S(IV) •>3'7 ir(iii) + s (V) (37)

Cu(II) + S(IV) --^ Cu(I) + 3(V) (38)

Ir(IV) + S(IV) ^5^^ Ir(III) + S(VI) (39)

Cu(II) + S(V) ^40 ^^^jj ^ g^^jj 4Qj

Cu(I) + Ir(IV) H^^ Cu(II) + Ir(III) (41)

23(V) ^5^^ S20g^- (42)

This mechanism is quite similar to the mechanism of

Higginson and Marshall (4) for the reaction of hexaaquoiron-

(III)-sulfur(IV) catalyzed by copper(II). According to

Higginson and Marshall's (4) data the increase in copper(II)

concentration caused an increase in the rate of disappear­

ance of iron (III). As shown in Table 6 an increase of •the

copper(II) concentration causes a sharper break in the dis­

appearance trace for iridium(IV). This sharper break is

believed to be caused by the increased competition for the

sulfur(IV) species by copper(II). Because of the steep dis­

appearance of iridium(IV) as shown in Fig. 9, Chapter III

with the sharp break, it is difficult to determine the rate

determining step. Whether the rate determining step is

equation 37 or 38 would entail further study. An

59

experiment investigated the reaction between copper(I) and

iridium(IV), shown in equation 41. The iridium(IV)-

copper(I) reaction was too fast for detection at stopped-

flow speeds.

Conclusions

This kinetic and mechanistic study of the hexachlo­

roiridate (IV) -sulfur (IV) reaction in a buffered chloride

medium has shown that the reaction falls into the previously

described pattern (18), where the reduction of substitution-

ally inert metal ions yield a mixture of dithionate and sul­

fate as shown in equations 1 and 2, Chapter I. Under the

conditions previously described, the following rate law was

obtained:

-d[Ir(IV)] _ 0.0147 [Ir (IV) ] [S (IV)]

dt [H+]

The rate law refutes the previous suggestion that the re-2-

active species would be 320^

Several mechanisms were postulated with all being dis­

carded except for mechanisms A and B'. Mechanism A had a

•SO^ radical intermediate, whereas mechanism B' had a

4-[IrClg-*303 ] radical intermediate. Though neither radical

proposed was detected by esr, a radical scavenger, acrylic

acid, was shown to polymerize when in the presence of the

reactants. Mechanism B' could be an inner sphere mechanism.

60

where the electron is transferred through a bond, or an

outer sphere mechanism, where •the electron is transferred

through space and not through the bond. Mechanism A is an

outer sphere mechanism.

Though the copper(II) catalyzed iridium(IV)-sulfur(IV)

reaction was only briefly investigated, it is concluded

that •this reaction is possibly worth an in depth study. An­

o-ther project could be a more elaborate investigation of

the kinetics and stoichiometry of the reaction between

iridium(IV) and hydrazine. This hydrazine study would

possibly add more information to the mechanism of the re­

duction of iridium(IV) as presented in this thesis.

LIST OF REFERENCES

1. D. W. Carlyle, J. Amer. Chem. Soc., 94, 4525 (1972).

2. H. Russell, Jr. and D. M. Yost, "Systematic Inorganic Chemistry." New York, New York: Prentice-Hall, Inc., 1944, pp. 358, 359.

3. W. M. Latimer, "Oxidation Potentials," 2nd ed. Engle-wood Cliffs, New Jersey: Prentice-Hall, Inc., 1952.

4. W. C. E. Higginson and J. W. Marshall, J. Chem. Soc., 447 (1957).

5. G. V. Bunau and M. Eigen, Zeit. fur Phys. Chem., Nene Folge, 7, 108 (1956).

6. G. A. Earwicker, J. Chem. Soc, 2620 (1960).

7. J. M. Lancaster and R. 3. Murray, J. Chem. Soc., A,

2755 (1971).

8. J. H. Swinehart, J. Inorg. Nucl. Chem., 27, 2313 (1967).

9. I. I. Malik and B. Singh, J. Indian Chem. Soc, 14, 435 (1937).

10. J. Veprek-Siska, et al., Colin. Czech. Chem. Commun. 30, 1390 (1965).

11. G. P. Haight, Jr., et al., J. Amer. Chem. Soc., 87, 3835 (1965).

12. W. D. Bonner and D. M. Yost, Ind. Eng. Chem., 18, 55 (1926).

13. J. J. Levison and 3. D. Robinson, J. Chem. Soc., A,

762 (1971).

14. H. Bassett and W. G. Parker, J. Chem. Soc, 1540 (1951).

15. D. G. Karraker, J. Phys. Chem., 67, 871 (1963).

16. D. W. Carlyle and O. Zeck (submitted for publication by Inorg. Chem., 1973).

17. A. Brown and W. C. E. Higginson, Chem. Comm., 725 (1967)

61

62

18. A. Brown and W. C. E. Higginson, J. Chem. Soc, Dal ton Transactions, 166 (1972).

19. R. O. C. Norman and P. M. Storey, J. Chem. Soc., (B) , 1009 (1971).

20. S. 0. Nielsen, K. Schested, and Z. P. Zagorski, J. Phys. Chem., 75, 3510 (1971).

21. E. Hayon, A. Treinin, and J. Wilf, J. Amer. Chem. Soc, 94, 47 (1972).

22. T. Kwan, T. Ozawa, and M. Setaka, Bull. Chem. Soc. Jap., 44, 3473 (1971).

23. B. D. Flockhart, K. J. Ivin, R. C. Pink, and B. D. Sharma, Chem. Comm., 339 (1971).

24. L. Dogliotti and E. Hayon, J. Phys. Chem., 72, 1800 (1968).

25. R. W. Brandon and C. 3. Elliott, Tetrahedron Lett., 4375 (1967).

26. R. O. C. Norman in "Essays in Free-Radical Chemistry," ed. R. O. C. Norman, Chem. Soc. Special Publ., 24, 1970, ch. 6.

27. Y. 3. Chiang, J. Craddock, D. Mickewich, and J. Turkevish, J. Phys. Chem., 70, 3509 (1966).

28. L. Omelka, A. Tkac, and K. Vesely, ibid., 75, 2575 (1971).

29. F. A. Cotton and G. Wilkinson, "Adv. Inorg. Chem.," Interscience, 1962, p. 545.

30. M. Falk and P. A. Giguere, Canad. J. Chem., 36, 1121 (1958).

31. L. H. Jones and E. McLaren, J. Chem. Phys. 28, 99 5 (1958).

32. M. Frydman, G. Nilsson, T. Rengemo, and L. G. Sillen, Acta Chem. Scand., 12, 878 (1958).

33. K. 3. Pitzer, J. Amer. Chem. Soc, 59, 2365 (1937).

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34. M. Eigen, K. Kustin, and G. Maass, Z. Phys. Chem.

(Frankfurt), 30, 130 (1961).

35. R. M. Golding, J. Chem. Soc, 3711 (1960).

36. A. W. Herlinger and T. V. Long, Inorg. Chem., 8, 2661 (1969).

37. F. P. Dwyer, H. A. McKenzie, and R. 3. Nyholm, J. Proc Roy. Soc. N. S. Wales, 81, 216 (1947).

38. F. A. Cotton and G. Wilkinson, "Adv. Inorg. Chem." New York: John Wiley and Sons, Inc., 1966, p. 1011.

39. C. 3. Garner and I. A. Poulsen, J. Amer. Chem. Soc , 84, 2032 (1962).

40. J. C. Chang and C. 3. Garner, Inorg. Chem., 4, 209 (1965).

41. Z. A. Muzykantova and R. I. Novoselov, Zh. Neorg. Khim., 15(11), 3084 (1970).

42. M. C. Agrawal, K. C. Gupta, and 3. P. Mushran, Indian J. Chem., 10, 642 (1972).

43. M. M. Jones, "Ligand Reactivity and Catalysis," Academic Press, 1968.

44. E. L. King and J. C. Templeton, J. Amer. Chem. Soc, 93, 7160 (1971).

45. P. Abley, E. R. Dockal, and J. Halpern, ibid., 94,

659 (1972).

46. J. Halpern and M. Pribanic, ibid., 90, 5942 (1968).

47. J. Halpern, R. J. Legare, and R. Lumry, ibid., 85,

680 (1963).

48. P. Abley and J. Halpern, J. Chem. Soc, 20, 1238 (1971).

49. J. Veprek-Siska, Ann. Genie Chim., 3, 126 (1967); Chem. Abst. 70, 23272Z (1969).

50. A. Hasnedl, K. Madlo, and J. Veprek-Siska, Collect. Czech. Chem. Comm., 36, 3096 (1971).

64

51. O. Zeck, Ph.D. Unpublished Ph.D. dissertation, Texas Tech University, 1972.

52. H. Bassett and A. J. Henry, J. Chem. Soc, 914 (1935).

53. A. I. Vogel, "A Textbook of Quantitative Inorganic Analysis." New York: John Wiley and Sons, Inc., 1961, p. 370.

54. D. W. Carlyle, Inorg. Chem., 10, 761 (1971).

55. A. I. Vogel, "A Textbook of Quantitative Inorganic Analysis." New York: John Wiley and Sons, Inc., 1961, p. 322.

56. 3. Nakamura. Unpublished Ph.D. dissertation. Univer­sity of Chicago, 1964, p. 128.

57. D. A. Robinson and R. H. Stokes, "Electrolyte Solu­tions." 2nd ed. London: Butterworths, 19 65, p. 339. The value of 1.71 x 10"^M for the acid dissociation constant of acetic acid was obtained from data from this reference.

58. J. Q. Adams, S. W. Nicksic, and J. R. Thomas, J. Chem. Phys., 45, 654 (1966).

59. J. H. E. Griffiths, J. Owen, and I. M. Ward, Proc. of Roy. Soc. of Lon., A 219, 526 (1953).

60. I. N. Kuz'minykh and T. B. Bomshtein, J. Appl. Chem., U.S.S.R., 24, 497 (1951).

61. This program is based on Report LASL-2367+ Addenda, Los Alamos Scientific Laboratory, Los Alamos, N. M., 1959.

62. D. J. Edge and R. 0. C. Norman, J. Chem. Soc, (B) , 182 (1969).

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