Electrochemical study of xenon trioxide

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Authors McClelland, Paul Harold, 1943-

Publisher The University of Arizona.

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ELECTROCHEMICAL STUDY' OF XEHOU TRIOXIDE

byPaul.Harold.McClelland

A Thesis Submitted to the Faculty of theDEPARTMENT OF CHEMISTRY

In Partial Fulfillment of the Requirements For the Degree of ' -

; MASTER OF SCIENCE.In the Craduate CollegeTHE UNIVERSITY OF ARIZONA

1970

STATEMENT BY AUTHOR

This thesis has been submitted in partial fulfillment of requirements for an advanced degree at the University of Arizona and is deposited in the University Library to be made available to borrowers under the rules of the Library.

Brief quotations from this thesis are allowable without special permission, provided that accurate acknowl­edgment of source is made. Requests for permission for extended quotations from or reproduction of this manuscript in whole or in part may be granted by the head of the major department or the Dean of the Graduate College when in his judgment the proposed use of the material is in the interests of scholarship. In all other instances, however, permission must be obtained from the author.

APPROVAL BY THESIS DIRECTOR This thesis has been approved on the date shown below:

L. RAMALEY Assistant Professor of Chemistry

* Date

TABLE OF CONTENTS

PageLIST OF ILLUSTRATIONS ,. . . . . 0 = . . » o . , . •» ivLIST OF TABLES © © © © © © © © ©. © © © © © © © © © "W

A 3ST3 A0T 0 © © & © © © © © © © © ■ . © © © © © © © © © "9*jL1» IETE0LUCTI0N . ©. © „ ■© © © © © © © © © © © © © ' © © © 12© EXPERIMENTAL • ©' © © © © © © © © © ©. © © © © © ' © © © 11

Ma"fce37ials ■ © © © © © © © © © © © © © © © © © © 11S 01*1x1/3*023, s © © © © © © © © © © © © © © © © © © 11Equipment; © © © © © © © © © © © © © © © © © © 1*3

3 © POLAROGRAPIIY © © © © © © © © © © © © © © © © © © © 164© STATIONARY ELECTRODE VOLTAI#IETRY © © © © © © © © © 265„. ROTATING DISK VOLTAMMBTRY © © © © © © © © © © © © . 46:6 © SUMMARY © © © © © © © © o-o © © © © © © © o © © © © 64-

APPENDIX A; POLAROGRAPHY © © © © © © ©' © © © © © . © 67APPENDIX Bs STATIONARY'ELECTRODE. TOLTAMECETRY © © © 70APPENDIX C; ROTATING DISK VOLTAMMBTRY © © © © © © 76REFERENCES © © © © © © © © © © © © © © © © © © © © 79

iii

LIST OF ILLUSTRATIONS

o o o o o o o

O O O O 0 . 0 o

O O O 0 - 0 0 - 0 o o o

Figure.1« Cell Design Used During this Study „20 A Typical Polarogram with Maximum „3= Concentration Study at pH 1 ..0 4» "Apparent vs pH of Xe(YI) Solutions 5» Composite of Polarograms at Various.pH’s 6« Stationary Electrode Voltammogram at 3.0 mve/sec0, 7o Stationary Electrode Voltammogram at 9.3 mv0/seCo, 8. Stationary Electrode Voltammogram at 45.7 mv./sec,

o o o o o o

o O O O O

Ip ^ 1+0.167 V vs C o o o o o o o o o 6 o o o o

o o o o o o o o10. Log ip. and Log i+0 # 1 Y vs Log v11. Current Functions vs Scan Rate & . . . . . .,. . .12. Corrected l/E Curves Compared to Theoretical

C U r V e O . O O O . O . ". 0 . 0 O . 0 0 O O O

13o Peak Potential Shift with Scan Rate14. Rotating Disk Voltammogram at 1.67 mv./sec

' 15. Rotating Disk Voltammogram at 3.34 mv./sec16. Rotating Disk Voltammogram at 8.35 mv./sec

O O O O o o o

o o o o o o o o o o. o o o

O 0 . 0 0 O O 0- 0 o o o o o o o o o

17. Ip and I ^ ^ vs Scan Rate18. : Ip and ^ vs C19. RDE Voltammogram Corrected for Background20. I/E Curves- Showing the Effects of 1^ and I^21. l/E Curves for 1 X. 10"% XIO, at Three Scan

Rates in 0.065M H2SO4 4: it-

0 0 " 0. 0 Q O 0 , 0 0 0 O

Page.. 14, 17 ;. 18 o 21

. . 22

. 27

. 28

. 29

. 34

. 40

o 48 v o 49

. :54 57

« 63

LIST OP TABLES

Table Page.to Some of the Simple Xenon Compounds 0 = « * * «, 22e Results of Polarography « »."«■ <» » «. o = » „ v o 693o Data for Figure 9 » -» o <, = o o <, o 0 = o 6 » = 704o Data for Figure. 10 ©.o p p p p p p p p p p p p p _715 p Data for Figure 11®. © p p p p © © © © © © © © © 726o Data for Figure 12 © © ©, © © © © © © © © © © © © 13

7© Data for Figure 13 © © © © © © © © © © .© © © © © 748© Value of D Related to Scan Rate andConcentration;© © © © © © © © © © © © © © © 75

. : ' -i9© Data for Figures 17 and 18 © © © © . © © © © © © © 7610© Variation of D with Scan Rate and.

Concentration ©:© © © © © © © © © © © © © © 78

v

ABSTRACT

The polarographic reduction of Xe(VI) was investi­gated, over the concentration range 5„3 X 10“^M to 1.05 X 10""%? and over the pH range of 1.0 to 13.0. Thorough ex- amination of the system at pH:1.0 showed a maxima above 2.10 X 10^4% concentration, and a diffusion coefficient of 8.0 X.10" sq= cm./sec. The results of Jaselskis (Science 143. .1324$, 1964.) were verified except for the slope of

vs pH.. The reduction of Xe(VI) at an oxidised platinum

electrode was studied using linear sweep voltammetry at a spherical electrode and a rotating disk electrode. Xe(VI) reduction was initiated by PtO reduction and continued after PtO removal to form a current plateau. Covering the electrode with adsorbed iodide or iodine caused the; *" catalysis at PtO reduction potentials to stop, and forced the irreversible Xe(VI) reduction to more cathodic poten­tials. Calculation of the diffusion coefficient using the Levich equationp yielded an average value of 8.0 X 10 sq cm./sec. with or without adsorbed iodine or iodide.

CHAPTER 1

. INTRODUCTION

Before 1962, noble gases had not been known to form true chemical compoundso Because of their filled valence shells and the. consequent difficulty of adding or removing ah electron, it had become a text book axiom that the noble gases were chemically inert. According to Selig, Malm and Claasen (1964), Dr. Linus Pauling noted in 1933, that the noble gases might react with fluorine or oxygen to form XeFg KrFg or H^XeOg. Experimental efforts, to produce these compounds failed, adding to the conviction that the noble gases were indeed inert, with the only known compounds being clathrates (enclosure compounds)

These convictions were destroyed by Hell Bartlett at the University of British Columbia. While working with oxygen and platinum hexafluoride, Bartlett (1962) reasoned that since the oxidation potential of xenon is lower than that of oxygen (1,2.13 V vs 12.20 7), xenon should form a compound similar to 02PtPg. , Amazingly, xenon and platinum . hexafluoride did react in the gns phase to form a yellow powder. Much work was then done;at Argonne National Laboratory and Velsewhere to synthesize noble gas Compounds.

2

Research with RuRg showed that the metal hexafluoride did not form an ionic compound with xenon? but fluorinated it.

As perplexing as this seemed to the traditional textbook mind, no special concept of bonding is needed to describe the noble gas compounds. Chemically the xenon fluorides are similar to the halogen fluorides; xenon tri- oxide is iso-electronic with I0^~9 and the ionization poten­tials of xenon and radon are lower than nitrogen, oxygen, chlorine, and fluorine.

Of the-noble gases, the chemistry of xenon is the most abundant. Radon chemistry is sparse due to the diffi­culty of handling the radioactive and short-lived element. Krypton chemistry is limited by the instability of the fluorides. Unlike the halogens, xenon possesses all.of the even oxidation, states from 2 to 8. Table 1 lists some of. the simple xenon compounds.

Table 1. Some of the Simple Xenon Compounds. .

FLUORIDBS 0XY7LU0RIDES OXIDES ADDUCTS . SALTSXeP2 XeOP4 Xe02* . XeP2:28bP^ Ba2XeOgXeP^ Xe02P2 Xe05 XeP.6sSbP5 Ea^XeOg

' ■ . - Xe04 : .

*Xe02 has only been seen as the ion in mass spectroscopyf

Except for the xenon-metal adducts, all of the xenon compounds are formed via the fluorides. Reactions 1, 2 and 3 describe xenon fluoride formation, XeEg a^d Xe?^ can be stored in dry quartz? however, XeEg will react to form SiF^ and XeOE^, Xenon to fluorine bonds have a strength of roughly 30 Kcal/mole, and readily form whenever fluorine free radicals are present in the elemental mixture (Gunn 1963)-

Xe + F2 -- — ^^;.7,l.,atm.-» xeP2 . (crystals) (1)Xe + 2F2 X e F 4 (crystals) (2)xe + % (crystals) (3)

' ' When the fluorides are hydrolyzed in acid or neutral solutions, the surface of the crystals turns momentarily yellow, and a stable, xenon containing solution is formed.Below are the hydrolysis reactions as described by Appiemanand Malm (1964)o

XeP2 + H2.0 — — -~-^Xe + 1/2 02 + 2HF. : (4)3XeF4 + 6H20 -— — ^-XeO^ + 2Xe +.3/2 02 + 12HF (5);XeF6 + 3H20 — — — ^-XeO, + 6HF (6)

In these reactions, the principal product was xenontrioxide. Xenon trioxide, xenic acid as it has been wrongly called, appears to be molecular in solution, A 0,78M sol­ution of GCe(VI), 10“^M in HCIO^, had a corrected net molal freezing point depression of 1,95+0,15° C, The molal freezing point constant for Water is 1.85° C, indicating

that the Xe(VI) solution was essentially molecular,Appleman and Malm (1964) also noted that a 0o02M solution of Xe(Vl) in 10*"'% HCIO^ had a molal conductance of less

. than 0,04 mho-cm^. The conductance of 10“% HCIO^ alone isO0,0417 mho-cm , The Xe(YI) did not contribute to the con-.

ductivity of the solution. Measurements of the Raman spec­tra of. Xe(YI) .solutions by Claaseri and Knapp (1964) suggested that they contained principally molecular XeO^,

Xenon trioxide crystals have a structure that is a regular array of XeO^ molecules equivalent to and isoelec- tronic with solid HIO^ (Koch and Williamson 1964), Molec-

. ular XeO? is a trigonal pyramid with the O-Xe-O bond angle of 103° and a length of 1,76 A, (Appleman" and Malm 1964),

Xenon trioxide is not thermodynamically stable, having a heat formation pf +96 Kcal/mole, However, it does form very rapidly and.is chemically stable indefinitely in acidic solutions. It is explosive when dry.

As mentioned before, the term xeniC acid is a mis­nomer, Appleman and Malm (1964) noted that the solutions ofXe(Yl) had a weakly acidic behavior with a pKa of roughly 10,5, At a pH of 10, a reaction occurs with species in­volved probably being XeOy + 0H“ -=5= HXeO^“ , The measured K .forl-the titration of XeO^;with NaOH was 6,7 X110"*% Dispro- portionation also starts to occur at that pH to the perxenate structure of Xe0g"% (Koch and Williamson 1964), The fact, that Xe0j5 has no appreciable basicity was indicated by the .

5less than 1# loss of hydrogen activity in a 0.1M HCIO^ solu­tion which is 0.05M in Xe(VI). A further analogy can be made to halogen chemistry considering the disproportionation of C103~ to C104~ in strongly basic solutions.

Base hydrolysis of the fluorides or base addition to Xe(VI) yields a solution containing Xe(VIIl). Except in the case where ozone is present, base hydrolysis is accompanied by loss of xenon according to reaction 7. XeF^ and XeF^

2XeFg + 4Na+ + 160H~ — >Na.XeOg + Xe + 02 + 12F- + BHgO (7)

both form Xe(VIIl) (Appleman and Malm 1964), and XeF. decom- poses as before. Koch and Williamson (1964) postulated a mechanism for the disproportionation, but Appleman and Malm found with concentrations of 1 X 10“ or lower that no mechanism satisfactorily explained the results. Appleman and Malm also noted that the presence of fluoride catalyzed the disproportionation to a great and variable extent.

The perxenates, Xe(VIII) salts, are stable in air and alkaline solutions. Sodium perxenate precipitates quantitatively from 1M ITaOH and is only soluble in dis­tilled water to the extent of 0.025M. At pH 7, Xe(VTII) decomposes instantaneously to Xe(VT) and oxygen. Sodium perxenate was found by Appleman and Malm (1964) to become anhydrous at roughly 100°C and decompose at 360°C. X-ray diffraction studies of Hamilton, Ibers and Mackenzie (1963)

and Zalkin et al. (1963) show the presence of octahedral XeOg"* ions in the crystal lattice of Na^XeO^ • GE^O and • 8H20.

Xenon solutions are very powerful oxidizers. The estimated electrode potentials for the acidic and basic solutions were reported by Appleman and Malm (1964) as given below.

pH 1 H^XeOg + 2H+ + 2e — » XeO, + 3H20 E° = 3.0 VpH 1 XeO™ + 6H+ + 6e — ► Xe + 3H20 B° = 1.8 VpH 11 HXeOg-5 + 4H+ + 2e— > HXeO ~ + 2H20 E° = 0.9 VpH 11 HXeO “ + 7H+ + 6e — » Xe + 4H20 E° = 0.9 V

The oxidizing power of these solutions had drawn some attention. From the literature the following applica­tions of the noble gas compounds were noted: 1) The oxi­dation of Pu+3 to Pu+4 to facilitate separation from the lanthanides (Cleveland 1965). 2) Oxidation of Am+ to Am+^for a similar reason (Holcomb 1965). 3) Determination ofprimary alcohols by oxidation to organic acids (Jaselskis and Warriner 1966). 4) Rapid spectrophotometric determi­nation of manganese in steel (Hyman 1966). 5) Slow oxi­dation of vicdiols (Jaselskis and Vas 1964).

To date only two electrochemical references have appeared in the literature. Jaselskis (1964&) polarograph- ically reduced Xe(VT) in buffered 0.1M K^SO^ over a pH range 4.6 to 8.0. The half-wave potential varied from

7vs N.H.E. (i.e., a slope of 0.077 V/pH unit). A well formed cathodic wave equivalent to a six electron reduction was obtained over the concentration range 4 X 10~^ to 2 X 10"" M. Jaselskis* presentation of a typical polarogram indicated no maximum and mentioned no addition of any maximum suppressor. The only anomaly was a small anodic prewave that coalesced into the cathodic wave with increasing Xe(Vl) concentration. The diffusion current constant, I, for the Xe(VI) reaction was reported to be between 9.93 and 11.05. The anodic pre­wave was stated by Jaselskis (1964a) to be similar to the anodic wave of periodate and due to the formation of an unstable mercurous xenate. The presence of such an inter­mediate seems unlikely with an oxidizer as strong as Xe(VI). Buffers for the system were 0.02M acetate at pH 4.6 and 0.2M phosphate at pH 8.0.

Jaselskis (1964b) polarographically reduced XeFg,XeO^ and Na^XeO^ in 0.1M K^SO^, buffered with 0.001M NaOH, 0.04M acetic acid-0.02M sodium acetate with and without 0.1M NaF, or 0.005M HagCO^-O.OO^M NaHCO^. The reduction of Xe(VIII) at pH 10 and 11 produced a curve with a maximum and an I of 9.8. Maxima on the XeF^ reduction curves were sensitive to F” concentration, and completely suppressed by 0.1M HaF buffer. No mention was made to any distortion of the Xe(VI) waves. Reduction curves of Xe(VI) were similar to those obtained by Jaselskis (1964a) except that I was 7.7 instead of 9.8. Xe(VI) in Jaselskis (1964b) was prepared

by the. decomposition of sodium perxenate in the reaction cell by adding one drop of 1o0M while in Jaselskis(1964a) the "xenic acid" solution was used as received from Appleman and Malm (Argonne National Laboratories) $, with no mention of the-method of preparation. All reduction cur­rents were proportional to concentration and the number of electrons in the reduction.

Prom Jaselskis‘ accounts (1964a?b) the electro­chemistry of Xe(¥I) does not . appear to be exciting or . unusual. The mercury electrode reactions» however, are only a small part of the totality of applied electrochemis­try. The reactions of Xe(VI) have to be studied over a wider. pH and coneentration range and at different electrodes before it can be dismissed without further comment.

In Xe(¥I), the. chemist may have an Aladdin1 s lamp. High oxidation potential and apparent reactivity^ open the possibility of.catalytic analysis.: Catalytic waves depend upon homogeneous chemical reactions for at least part of . their limiting current „■ In some cases catalytic currents allow analysis of very small amounts of electro-active ma­terial, or analysis of a substance which is not electro- active. In a catalytic reaction sequence such as below, it is Step I which is the basis of the scheme. Steps I and II need not be reversible, but Step II must be at least as fast as Step I, for the system to be useful analytically. In the ideal catalytic scheme; the electro-reduction of Ox

should be very rapid and reversible, regeneration of Ox in

Step II should be fast, thus allowing determination of Y and Z should not contribute to the limiting current.For the case where Y is being determined catalytically, it need not be electro-active, but in the case where Ox is being determined, both Y and Z must not be electro-active.If Step II is slow, no noticeable catalytic current will flow. If Step II is fairly fast, catalytic current will flow and it will be possible to determine Y and K^. When Step II is very fast, it will not be possible to differ­entiate between the diffusion and catalytic currents in the presence of Y. Only the unusually high current flow and the subsequently larger diffusion coefficient will indicate that the reaction is truly catalytic.

Hydrogen peroxide solutions are rich in catalytic reactions, particularly the HgOg - heavy metal solutions.A very small amount of molybdate (10“%) catalyzes the rapid reduction H^Og at roughly +0.65 V. H202, itself, is not even completely electro-active at much more cathodic potentials. HgOg has a reduction potential of +1.77 V, reducing to H^O and 0H~. Xenon trioxide has a higher reduc­tion potential and reduces to Xe, H^O and sometimes 02, none of which will permanently contaminate the solution.

Ox + ne — ^ Red Step IStep II Step 11Ox + ne” Red

/ . 10

Provided that XeO^ reacts with electrochemically reduced species with enough speed9 many catalytic reactions are possibleo

By studying the electrochemistry of XeO^ solutions it was hoped that more would be learned of the basic electro­chemistry of the noble gas compounds? and that it might be feasible to apply, xenon compounds to trace catalytic anal- ysis. The study, was accomplished by the use of dropping mercury$ stationary electrode, and rotating disk electrode polarography at various pH's and concentrations« Reaction - reversibility and electrode kinetics were studied using the solid spherical platinum electrode and cyclic volt- ammetry,.

CHAPTER 2

EXPERBESNTAI,

MaterialsAll chemicals were analytical reagent grade

(Mallinckrodt Chemical Works) and used as received, except for the potassium iodate and. the sodium perxenate <, Primary standard potassium iodate v/as obtained from G. P. Smith Co., and the sodium perxenate was purchased from Peninsular Chem Research Inc., G-ainsville, Florida „ All solutions were made from doubly distilled water.

Solutions .Xenon trioxide stock solutions were prepared di­

rectly from the perxenate, It was first attempted to dis­solve the perxenate directly in the test buffers. The perxenate dissolved well upjto pH fi, then troubles arose.An attempt was made to dissolve the perxenate in acid and neutralize to the desired pH and buffer. As a consequence of over-neutralization on the first attempt, there was a low xenon titer in some of the solutions due to the dis- proportionafion of Xe(Vl) at higher pH?s„ Because of the pH influence» it was decided to do the rest of the work at lower pH’s where the solutions were more stable,

: / ' . "■ 11 : : ' . . ' •

12For the final runs, the stock XeO^ solutions were

made by weighing 36.4 mg. of Ha^XeOg • 25^0 into a 100 ml. volumetric flask and adding 0.065M H^SO^ to the mark. Dis­solution was accompanied by a momentary yellow coloration and the evolution of bubbles. Such solutions were stablefor at least two months. Solutions of varying concentra­

tetions, as low as 5 X 10 VM, were prepared by dilution of the stock solution.

Analytical determination of oxidizing power in the Xe(VI) solutions was done by adding a test aliquot to a volumetric flask containing nitrogen saturated 0.1M H^SO^. Enough solid KI was added to bring the final solution to 0.5$ KI. The amount of generated was read at the 353 mn. absorption maximum in a Beckman Model DB spectro­photometer, against a deionized water blank. A standar­dization curve was prepared uding KIO^ (primary standard), which has the same molar equivalency as XeO^. A 9.997 X 10”^! standard stock solution was prepared to use throughout the study. Dilutions similar to the test solu­tions were run daily to test the procedure and reagents. It was found necessary to store the water and HgSO^ under ni­trogen before dispensing for analysis. This was done by fitting gas dispersion tubes from a Research Grade nitrogen tank to the dispensing flasks.

The purity of the sodium perxenate was tested early in the study by weighing out 35.57 mg., and placing it in a

100 ml, volumetric flask, which was filled with 0„1M Eg30^. One milliliter, aliquots of the above solution were taken to 250 ml„ final volume with the necessary reagents, and read in the spectrophotometer„ The results indicated that the perxenate was 97®3$ pure taken as Ea^XeQg « 2H20, which was a highly probable form. This also verified that the acid solution contained Xe(VI) as expected. The possi­bility of Xe(VIII) being present in some of the solutions was verified by the addition of solid KI to the test ali­quot before acid was added, as recommended by Appleman and Malm (1964:) 0 Thus, by adding Kl before acid was added, all (available Xe (.VIII) was reacted with I" before it could de­compose to Xe(VI); this decomposition occurs instantane­ously in acidic solutions.

EquipmentTest cells for the study were built in the chemis­

try glass shop. The first cell design tried had no isola­tion compartment between the reference electrode and the working electrode compartments. As a result, the test so­lution was contaminated by leakage from the reference electrode. Figure 1 shows the second cell design with the rotating disk electrode in place. The other electrodes fit through the hole in the top plate to the same depth. ,

The electrical equipment was set up around the cell in such a way that the (BDS) rotating disk electrode

14

WORKING- ELECTRODE

I t i LROTOR

COUNTERELECTRODE

REFERENCEELECTRODE

JL

KC1SAT TEST SOLN

Fine porosity sintered glass

LUGGINCAPILLARY

REFERENCE ISOLATION WORKINGCELL CELLS CELL

Figure 1 Cell Design Used During this Study

: 15voltarametry P, stationary electrode cyclic voltaimnetry and the DICE polarography could be done in that order without disturbing the cell for more.than a few seconds to change electrodes. The DME and RBE voltammetry were performed using the Heath EIB'/<-20a recorder; ETJW-19 operational amplifier manifold; BIJV7-19-2 polarographic module; and a chopper stabilizer, similar to the EOT-19-4.

Gyclic voltammetry was performed using a manifold of Philbrick operational amplifiers similar to that used by Ramaley and Enke (1965 ) ,

The rotating disk electrode was made by sealing a short length of 0,025 in, platinum wire into the center of a piece of 6 mm, soft glass and lapping the end flat. This provided a flat disk with the boundary far removed from the edge of the platinum. The rotor for the system was a 600 RPM Sargent Synchronous Rotor, The solid spherical elec­trode was made by melting a platinum wire9 sealed in soft .glass tubing$, into a sphere roughly 2 ram in diameter or

20,15 cm area.

CHAPTER 3

POL ARO GRAPH!

A typical polarogram of Xe(VT) at pH 5.0 is shoi-m in Figure 2» The presence of a maximum was typical of the wave at air pH's and concentratiphs greater than 2 X. Fortunately^; the rapid fall of the maximum allowed the dif­fusion limited current to be measured at a potential of -Oo25 Vo At concentrations of 1 X .10"% or less, little or no wave distortion was recorded „ The maximum made estima­tion Of E _ very difficult when accompanied by the close , proximity of large anodic currents 0 In the: case of waves such as in Figure 3$ where no indication of i^ can be seen on the ascending portion of the wave5 the "apparentn E, . was qualitatively estimated by using one-half the average current value at -Qo2;5 V and extrapolating back to the as­cending portion of the wave 9 as shown by example„ Current measurements in this study were all taken at one-half the maximum recorded current fluctuation per drop because of the high damping necessary to obtain.a useable polarogram. Correctionifor background current was not considered im- • portant because of its very small magnitude, The back­ground scan for the 0,065M HgSO^, pH 1,0 system was run at 2 microamperes full scale sensitivity to yield a. zero

' 16 '■ '

Microamperes

200-

150-

pH 3.0 9.65 X 10_4M Xe(VI)

Scan rate: 0.1 V/min

50-

Volts vs N.H.E+0.55 +0.45 , m +0.35 +0.25 +0.15Figure 2 A Typical Polarogram with Maximum +0.05 -0.05 -0.15 -0.25

18

200 jia

100 yn

1.05 X 10~ M Xe(VI)

_ _ _50_jia.___

5.29 X 10-4M Xe(VI)

10

r 2 i ~ x ~f6 m ~xo~(vi') " T '

2 na.1.06 X 10-4% xe(YI)

-0.200.0+0.40 +0.20+0.60

Volts vs N.H.E.

Figure 3 Concentration Study at pH 1.0.

current potential of.+0.27 V and an apparent electro~capil­lary maximump ECMf of ~0e45V. Zero current potential was the potential where the cathodic and anodic currents were . equal and thus the net indicated current was zero j ie, equivalent to grounding the recorder input. The base of anodic or cathodic waves does not necessarily start on the indicated zero current line, but depends also upon the amount of current due to charging the mercury surface and reducible impurities in the solution. The -ECM was deter­mined as. the potential of-minimum current fluctuation per drop in absence of Xe(iri), The zero current line was used . as the baseline for all later polarograms because the resid­ual current at -0.25 V was only 0.1 microampere and. 0.5 microampere at pH 6.7. Data for the calculation of diffu­sion coefficients and other plots was corrected for the residual current, and were all estimatedusing the above criteria on waves without maxima when possible, No maximum suppressors were used in this study because of the strong.oxidizing capabilities of Xe(VI)0

Solutions of pH 1,0 and pH 3.0 appear to have the same half-wave potential. This is due? in part, to the , shift in Xe(VI) reduction potential: ;and the close prox­imity of the anodic oxidation of mercury, as explained below. In Figure. 2, the reduction wave was separated from the anodic currents by a sufficient amount such that can be measured; while at pH's 1.0 to 1.2, Figure 3, a definite .

■ ' 20wave cannot be distinguished; between anodic currents and the maximumc Figure 4S "apparent” vs pHp illustratesthe slope change at pH 3=0 very clearly0 In all cases,the apparent” E _ was never more than 0.020 7 negative from the zero current potential and the start of mercury oxi­dation, and thus a plot of E _.q vs pH would look almostexactly the' same. Reduction of Xe(VI) commenced very soon after cessation of anodic currents, and may have actually been present in the oxidation of the mercury surface at pH 1.0. The latter point can only be surmised. For a bet­ter understanding of the complexity of the system. Figure 5, a composite of polarograms at pH9s 3.0, 3.5, 6.7, and 11.0 should be consulted.

Examination of the polarograms; at pH 1.0 and pH 3.0 indicates that there is a difference in potential between the oxidation of Hg and reduction of Xe(Vl). At pH 1, the reaction was apparently already underway, but masked by the anodic current flow. The lack of any sign of the foot of the wave substantiates this position. Diffusion cur­rent measurements at more negative potentials did not appear to be disturbed by obscuring the wave foot or the maximum. - The only waves really;.,suitable -for "wave analysis ” were obtained at pH 3.0. The curve in Figure 2 is, however,% of such low sensitivity, 500 ua full scale, that an ac­curate picture of the wave is impossible. E^y^ - j^ has been estimated to be 0.012 volts. Examination of other

21

+0.65 -

Slope = -40 mv/pH-0.20 Vrr

pH = 5 units

rr

+0.15

All potentials vs IJ.H.B.

Figure 4 "ApparentM vs pH of Xe(Vl) Solutions.

22

! 1pH 6.7 I |2.68 X 1D-4M Xe(TI)

I 1I pH 3.5 I, 6.25 X lb-4M

50 %a

| pH 3.0 I /[ t I I 9.6 X"I0I"4K I . 50 pa II ! J I- 1 >* ^1 1 11

. - 1 /------| /pH 11.(j | O . m NagS041/ 1

^ I .; i jit A Ar 1 f I > y........ i

! ^ i ^ y pH 3.0 I 1

8.57IX 10_4M Xd(VI)I

1 A 3.25 X jl0-4M Xe(vt) 4->-----4------------1-------------- ! I

+0.65

I| 50 pa }| I \ I

+0.55 40.45 +0.35Volts vs N.H.E.

+0.25 +0.15

Figure 5 Composite of Polarograms at Various pH's.

23curves obtained with more sensitivity yielded - E^y^ranging from 0,016 V to 0.020 V. Applying equation 5-48 of Meites (1963), (^3/4 - El/4 ~ ~0»0517/<xna), (Xna values of 2.58 to 4.3 were obtained, which are all rather large.

Two completely independent concentration studies were run using a 0.065M H^SO^ supporting electrolyte. The average diffusion current constant, I, for both studies was 11.7 based on peak currents, corresponding to a diffusion coefficient of 8.0 X 10"^ cm^/sec. The study of 7-21-66 deserves more attention due to the greater care taken in protecting the solutions from oxygen contamination. The greater care was rewarded by a narrower spread in the data. See Appendix A for the complete synopsis of both studies. These values are comparable to Jaselskis* experience and to I05“ in perchloric or sulfuric acid. Meites (1963) quotes a diffusion current constant of 12.06 for I0y~.

Very favorable results are presented (vide infra) when compared against the papers of B. Jaselskis (1964a,b). Our diffusion coefficients and I values are in close prox­imity and the general behavior of the system verifies the re­sults of Jaselskis. Jaselskis used a 0.1M KgSO^ supporting electrolyte and 0.02M acetate, or 0.2 phosphate buffer for pH control. In this study, 0.1M H^SO^ was the supporting electrolyte at pH 1.0, but for the other pH's the buffers, themselves, were the supporting electrolyte. This feature may well account for the difference in E^/pH. Buffers for

24the pH study were; 0 =, 1M phosphoric-mono phosphate for pH 3.0, 0„1M monophosphate-biphosphate for pH 6.7, 0.1 hioarbonate- carhonate for pH's 10 and 11, and 0„1M biphosphate-triphos­phate for pH 13. All solutions were also 0.01M in sulfate from the stock Xe(VI) solution.

Polarograms of this study, Figure 5, definitely show the anodic waves', some of which appear to have maxima. Sim­ilar anodic waves were present with the supporting elec­trolyte only at pH 6.7$, 10, 11, and 13. Unfortunately, at pH 1.0 it was impossible to see any anodic waves because of the strong mercury oxidation, but their presence at higher pH's sheds doubt o® the importance of an unstable xenate.

No attempt was made in this study to prove that the current was diffusion controlled, but since Jaselskis did prove diffusion control, there is no reason to presume that diffusion control was not in effect in this study.

The polarography of Xe(VI) in 0.065M H^SO^ (vide infra) substantiated the experiences of Jaselskis (1964a) and provided a basis for further investigation. The dif­fusion current constant and diffusion coefficient point out the predicted similarities of XeO^ and 10^ and will be useful ip later calculations and comparisons. The differ­ence in slope of Si vs.pH, between this study G.040 V/pH

. • ; . : T \ " / V ; .. 'unit and Jaselskis 0.077 V/pH unit, denotes a change in thereaction controlling the apparent E,'. The .differences in

' ■ ■ ,supporting.electrolyte may have beep partially responsible„ Within the limited ranges of concentration and pH,Jaselskis (1964a,b) made a more thorough investigation than found in this study. The results obtained here are in- general agreement with Jaselskis^ and cover a much wider range of concentration and pH. It is interesting to note that Jaseiskise reported results are all below the range, of concentrations where maxima were found in this study.

CHAPTER 4

STATIONARY ELECTRODE VOLTAHMETRY

. Stationary electrode voltammetry was performed to investigate the possible reactions of Xe(VT) at the platinum electrodeo Figures 6? 7 and 8 describe the current flow to an oxidized stationary platinum electrode for five concen­trations of Xe(VI) swept at 3*0, 9=3 and 45.7 millivolts per second between the potentials >1.23 V and +0.167 VH .H .E . Throughout this studyp 0.065M HgSO^ was used as the supporting electrolyte. Just beyond the limits of the current-potential (l/E) curves $, oxygen and hydrogen evo­lution occurred. Pretreatment of the electrode consisted of oxidation at +1.23 V for 30 seconds with nitrogen bubbling through the cell.

Collective examination of the I/E curves showed five definite trends.

1o All anodic currents at 1.23 V were independent of the Xe(VI) coneentration. This is best seen in "A" in Figure 7.

2. Cathodic current peaks were a non-linear function of Xe(VI) concentration and occurred at the same. potential as the peak current for platinum oxide reduction. Note region "B” in Figures 7 and 8.

' 7 . ■' 26 -

+0.60 V(1) 1.05 X 10~3M Xe(Vl)(2) 5.29 X 10-4M Xe(VI)(5) 2.10 X 10-4M Xe(VI)(4) 1.06 X 10-4M Xe(VI)(5) 0.065M H2S0.Scan rate: 3.0 X 10"3 V/sec.(----) Cathodic sweep(----) Anodic sweep

20 pa—

10 ua—

N.H.E.

+1.050 V

potentials

+0t33 +oT5i +0:60 +ol69 ^ +0:87-r xy k v I i V * V xy i vy # v v / • iFigure 6 Stationary Electrode Voltammogram at 3.0 mv./sec. ro-4

28

r B i— >* |

1

(1) 1 .05 X 10-% Xe(VI)(2) 5.29 X 10"4M Xe(VI)(5) 2.10 X 10_4M Xe(VI)(4) 1.06 X 10-4M Xe(VI)(5) 5.3 X 10-% Xe(VI)(6) 0.065M H2S0.(--- ) Cathodic sweep(--- ), Anodic sweepScan rate: 9.3 X 10“^ V/s

+0.167 0.39 0.67 0.95 +1.23 Vvs N.H.E.

Figure 7 Stationary Electrode Voltammogram at 9.3 mv./sec.

29

160 4-(1) 1.05 X 10 M Xe(VI)(2) 5.29 X 10~4M Xe(VI)(3) 2.10 X 10_4M Xe(VI)(4) 1.06 X 10-4M Xe(VI)(5) 5.3 X 10-5m Xe(VI)(6) 0.065M HoS0.

150 -

60 —Scan rate:

4.57 X 1CT2 V/sec.50 -

40 —

)ia50 -

(3)

20

40.96 V 0.56 4-0.167 0.42 0.85 4-1.25VVolts vs N.H.E.

Figure 8 Stationary Electrode VoItammogram at 45.7 mv./sec.

30

3o Between 4-0.60 V and 4-0.85 V the cathodic current flow is higher on the anodic sweep than on the cathodic sweep.

4o Cathodic current at scan reversal was a non-linear function of Xe(VI) concentration.

5. As scan rate increasedP the cathodic curves became more sharply peaked.The independence of the anodic current with Xe(VI)

concentration was expected due to the very irreversible nature of the system and the instability of the lower oxi­dation states of xenon in solution. Enhancement of the platinum electrode reaction suggests catalytic reduction of Xe(VI) in a manner similar to the surface catalyzed reduc­tion of iodate in dilute sulfuric acid (Anson 1959). Since XeO^ and 10^“ are iso-electronic and so similar in, other respects? it is not unusual that the l/E curves for the two should be similar. Unfortunately? no work was available to compare the concentration dependence of ICy™ and XeO^ re­duction at a platinum electrode. Anson's (1959) study found that I05~ was not reduced at a reduced electrode in a manner such that a distinct wave could be ascertained. However? at a platinum electrode whose surface is being reduced? a large peak is formed at the same, potential as the reduc­tion of the oxide film. Current decay of this peak was so rapid that at 0.40 V cathodic of E„ only 20% of peak cur- . rent flowed. Anson explains the catalysis of the 10^”

. 31reduction "by a mechanism involving electron transfer to iodate from the electrode via a platinum oxide bridge* which is conductive only when the oxide is being reducedo The cathodic current then divides between iodate and platinum oxide reduction. As time passes*.the platinum oxide is removed and the current decays very rapidly. Such a mech­anism is also plausible for the Xe(VI) reduction with minor adaption. The major differences that can be noted are, that here* the current does not decay as rapidly as in the ICy™ reduction and that an actual plateau was formed* denoted as Region C* Figures 6 and 7. The slow current decay and plateau may be due to a separate reduction reaction or to oxidation of the electrode by Xe(VI). If Xe(VI) is electro­active independent of the oxide bridge at more cathodic potentials* the ascending portion of the wave may not be distinguishable from the PtO reduction* but the plateau cur- . rent should be treatable mathematically from a bulk diffu­sion concept. The expected reaction would be a totallyirreversible charge transfer.

■ ■ ■ *The various mechanisms through which it might be

possible to explain the observed l/E curves are listed below.1. Simple irreversible reduction. Ox + ne — R..2. Chemical reaction preceding an irreversible reduc­

tion, Z =2= Ox, Ox + ne —s-R.3. Irreversible reduction on a clean platinum surface,

Ox + ne R.

324» Irreversible reduction involving the platinum.sur­

face film Ox + ne 5c Irreversible reduction preceding through the

platinum oxide reduction,' PtO -f 2e + 2H+-^Pt + H20 n Pt + Ox-e-n PtO 4- R,

Other mechanisms involving reversible reductions or chemical reactions following electrochemical reduction can be dis­missed because of the irreversibility of the Xe(Yl) to Xe° reduction, the instability of the lower oxidation states of xenon in solution and the lack of suitable ions in solution for other reactions.

Figure 9 shows the reduction current, corrected forbackground,, at E„ to be a numerical function of concentrationPat each scan rate, but to vary with scan rate as a non­linear function. The corrected Xe(71) current at 4-0.167 Y was probably also linear with respect to concentration, but it was much more difficult to accurately compile. The ap­parent influence of scan rate was less than at E , butPdefinitely not linear. Eased on Figure 9 alone, one mightthink that the Xe(YI) current at E_ was due to a diffusion- Pcontrolled transfer while that at 4-0.167 Y was kinetically. involved,

Figure 10 shows plots log ip and log i+Q y vs log v to obtain the functional dependence of 1 vs v. The

40 45.7 nnr/sec.

35

3018.7 mv/sec.

25>ia2 0 -

3.0 mv/sec.15"

10 -

10

X 10_4M Xe(Vl)Figure 9 Ip and I+q j ^ v vs C

1+0.167 V Ts C

40-

35-

jia

9.3- mv/sec.20-

3.0 mv/sec.15-

10 -

10

X 10~4M Xe(Vl)(See Appendix B for data.) vi

34

log 1+0.167 7(1) 1.05 X10-5M Xe(VI)(2) 5.29 X 10-4M Xe(VI)(5) 2.10 X 10-4M Xe(Vl)(4) 1 .06 X 10"% Xe(VI)(5) 5.3 X 10 M Xe(VI)(6) 0.065M HoS0.

2.04-(1)(2)

(4)

(6)2.00.8

log v

l°g Ip(1) 5.29 X 10"% Xe(VI)(2) 2.10 X 10"% Xe(VI)(3) 0.065M HgSO.(4) 1.06 X 10"% Xe(VI)

1.54

(4)0.54

log vFigure 10 log 1 and log v vs Log v

35slope of the Xe(VI) curves at are centered about 0.5, indicating i is proportional to v^. This would fit any of the simple irreversible reductions. Log i+Q vs log v, however, is a family of curved lines, indicating a kinetic involvement. Log i (H^SO^) has two slopes, between 5.0 and9.3 mv/sec. s = 0.457 or probably i v *, and between9.3 and 45.7 mv/sec. s = 0.90 or i^ v. At +0.167 V, the HgSO^ curve has a slope of 1.1 to 1.2 possibly indicating a linear function.

In a study of the platinum oxide system, Laitinen and Enke (1960) found that the film formation was so irre­versible that it could not establish the potential of the electrode even under zero current conditions. Inherent hys- tersis found between the oxidation and reduction curves, indicated that the reactions were either very slow or oc­curred via different mechanisms. Using the variation of double layer capacitance as a function of surface coverage, Laitinen and Enke (1960) were also able to conclude that the rate of oxide removal increased with increasing cathodic potential. At +0.85 V the "half-life" of the oxide was 60 seconds and at +0.75 V only 40 seconds.

To investigate the observed reactions from a less empirical point of view, the l/S curves were analyzed using the diagnostic criteria of the Theory of Stationary Elec­trode Polarography as summarized by Nicholson and Shain (1964). Nicholson and Shain characterized eight classes of

36

reactions involving reversible and irreversible charge transfers coupled to, preceding and following chemical re­actions and catalytic reactions. Those of concern to this study are shown below.Class II Irreversible Charge Transfer 0 + ne ->R.

IV Chemical reaction preceding an irreversible charge transfer Z^O, 0 + ne-^R.

In their diagnostic scheme, Nicholson and Shain(1964) used the variation of current function, peak poten­tial shift and anodic to cathodic current ratios as a func­tion of scan rate. Figure 11 presents peak current function, ip/v^ and i+Q y/vT vs v in an effort to separate dif­fusion processes from kinetically controlled processes. This analysis can be misleading in that a very fast coupled re­action may behave like a simple reduction, depending upon the scan rate. In this light, note that here we have cov­ered only a 30-fold range in scan rate, whereas some reac^ tions only change characteristics over a 100-fold range.H^SO^ curves are included in Figure 11 for reference only. They are the product of a non-homogeneous reaction, which does not fit the mathematical model from which the analysis guides were derived. Xe(VI) curves at Bp and +0.167 V resemble only Class II, the simple irreversible reductions. The fit of the experimental data to the sample cases is not, however, particularly close.

Figure 11

38For. the sake of convenience, the curves for the

oxidized electrode in 0o065M shall be knovm as thePtO or curve and labeled on-, the figures as 0 »065Mh 2so4.

If the Xe(VT) reduction is a simple irreversible' reduction as possibly indicated by Figure 10 and 11, then comparison of a theoretical irreversible reduction wave to an experimental curve $ corrected for background, should provide positive proof 0 Figure 12 contains 3 corrected experimental curves and one theoretical curve calculated from the current functions of Nicholson and Shain (1964?page 710) assuming c<n = 0,40 for 5,29 X 10"% Xe(YI) at

.

9.3 mv/sec. scan rate. Ihe diffusion coefficient was8.0 X 10“ sq. cm./sec. derived from the polarography.The most serious differences between the curves are; 1) The theoretical curve has a much higher current peak, 10 times the experimental, curve, 2) a .much narrower:current peak and 5) no indication of a definite current plateau. The experimental curves are definitely some reaction or reac­tions other than a simple irreversible reduction.

Figure 13? the peak potential variation with scan rate? does not provide much data rsince the scan rate range was only thirty-fold. ; The data presented also suffers from a drift as high as +0.015 Y in E_ due to the manual syn- chronization of recorder and sweep generator, and the var­iation in the sweep time. Diffusion controlled processes

39

25 — 245 )ia(1) 5.29 X 10"% Xe(Vl) Theoretical irreversible

reduction curve(2) 5.29 X 10-4M Xe(Vl)(3) 2.10 X 10"% Xe(VI)(4) 1.06 X 10"% Xe(VI)(---) Catholic sweep(-- ) Anodic sweep

+0.24 +0.52 +0.80 Volts vs N.H.E.

+1.08 V

Figure 12 Corrected l/S Curves and Theoretical Irreversible Curve.

40

(1) 5.30 X 10-5M Xe(VI)(2) 5.29 X 10-4M Xe(VI)(3) 2.10 X 10-4Ii Xe(Vl)(4) 0.065M HoS0.AS/ Alog v /(5) 1.06 X 10"% Xe(VI)

+0.2 -

+ 0.1 -

- 0.1 -

- 0.2 -

-0.3 --

20 80Millivolts per second

See Appendix B for data tables

Figure 13 Peak Potential Shift with Scan Rate.

again show no dependence on scan rate and the different effects of scan rate on and Ej_ make this plot useful in kinetic cases« At best? Figure 13 resembles none of the sample cases of Eicholson and Shain. Within the limits'of the experimental error? Class IT? chemical reaction pre­ceding an irreversible charge transfer? might be possible? but highly unlikely? due to the failure of the peak height to increase rapidly with an Increase in scan rate? and the lack of any material capable of a: reversible chemical re­action with Xe(TI) save H «, The ratio of anodic to cathodic currents was not plotted because of the lack of any anodic current contributions from Xe(O)„

Despite various indications of Figures 10 and 11? no logical explanation for the behavior of the system can be found? save dependence on the platinum oxide reduction men- tioned earlier0 To explain the shape of the I/E curves? assume that the peak at 4-0.60 T is due to removal of a limited amount of oxide and that the rate constant has been shown to increase with increasing cathodic potentials. In­creasing scan rate increased i rapidly? due to an accelef-Pated rate constant and more rapid charging of the double layer. The increase in the background current at 0.167 T with scan rate is|:due .mostly to currents associated with hydrogen adsorption and charging of the double layer at higher rates. Current to the electrode on the anodic sweep is zero, until normally at 4-0.85 when redeposition of

platinum oxide commences? and rises until the oxide for­mation rate becomes constant? causing an anodic plateau at +1oO V as Seen in Figure 7 - Region A. The presence of Xe(Vl) greatly alters the cathodic and anodic curves below 4-1 <,0 Vo Above this potential? the effects of Xe(VI) cannot be seeno The importance of the' platinum oxide to the reduc­tion of Xe(TI) at -fd o 60 V can be seen in the rapid increase in, current at the PtO reduction potential, and in curves v": where the electrode was not anodically.pretreated. .In the case of curves where the electrode was not anodically pre­treated or conditioned by sweeping the potential range rapidly two or three times the peak current was. less than at an oxidized electrode. The enhancement of the peak cur­rent and the gradual disappearance of the peak with in­creasing Xe( VI) concentration can be explained as increased * current due to Xe(VI) reduction through formation of plat­inum oxide which is subsequently reduced,.and to electro- reduction of Xe(VI) via PtO electron bridging as suggested by Anson (1959). At +0.60 V the rapid rise in current was due mostly to electro-reduction of XeO^ at activated PtO sites. Below 2 X i0“ !€ Xe(VI) the current decayed after the peak in a manner similar to the PtO curves because of active site removal in the PtO. The plateau current at +0.167 V was also controlled by Xe(Vi) concentration but through a different reduction mechanism. At this potential PtO reduction was very rapid and the availability of activated ..

• 45PtO reduction sites very low.. Thus the fraction of the. cur­rent put into PtO formation was higher and the high currents of were not possible« Increasing scan rate, increased the peak height in the polarogram; but after cor­recting. for the increase in PtO current, one will note from Figure 9 that the Xe(VI) current at E only increased four-irfold for a fifteen-fold increase in scan rate, while the PtO reduction current increased ten-fold for a fifteen-fold increase in scan rate.. At -KM 67 V the same fifteen-fold increase: in scan rate increased Xe'(VI) currents only 1.7 times, while the PtO currents increased fifteen-fold with a scan rate increase of ten times, The same rate laws obvi­ously do not apply to Xe(VI) reduction at +0,60 Y and +0.167 V as illustrated in Figure TO. Also of interest was the decrease in multiplication factor with increased Xe(VI) concentration, supporting the mixed reaction concept because of the non-linearity of i vs concentrations and i vs scan rate.

Another explanation of the total wave picture could be conceived involving, two irreversible reactions. The first being PtO + 2e~ + 2H+— -Pt + 2H2O on the electrode surface, and the second starting only 0.100 to 0,200 V ca- thodic to the first wave, the reaction being XeO^ + 6e~ + 6H+—^Xe0 + ^hgO. According the sample I/E curves of Nicholson and Shain (1964), the second reaction might be a special case of Class IV, a chemical reaction preceding

44an irreversible charge transfer, where v^/K-k^ was 3.0 or greater or a case of catalytic reaction with irreversible charge transfer. The catalytic cases that appear similar have been ruled out previously for lack of proper species in solution. Understanding the role of the chemical re­action preceding an irreversible charge transfer, also requires that one find species to fit the system. The most likely equilibrium involving XeO^ would be a protonation reaction H+ + XeO^ HXeO^+, and the equilibrium constant for such a reaction would be very small, according to its molecular nature.

The explanation given by Laitinen and Enke (i960) can be used to interpret the high reverse sweep currents of Figures 6, 7 and 8. The platinum surface reduction pro­duces a peak centered at +0.60 V, while the anodic (reverse) sweep produces a long oxide formation wave starting at approximately +0.85 V and limiting at +1.00 V. Due to the proposed catalytic dependency of mechanisms 4 or 5 mentioned earlier on page 32, the Xe(VT) curves follow the PtO curves in producing a large sharp cathodic and a long anodic wave. This point is illustrated by the cor­rected experimental curves of Figure 12. Note that the current returns to the base line in a long wave in harmony with the hysteresis of the oxidation and reduction reactions. As long as the platinum surface remained electro-active, Xe(VI) was reduced. A mechanism consistent with the simple

irreversible charge transfer scheme, would require that the corrected l/E curve decrease in current starting before the peak current of that cathodic sweep.

Calculation of D using equation 6-29 of Delahay (1954), ip = 3.01 X 10^ n(<na)TAD^CvT for the peak current in an irreversible charge transfer reaction, indicated kinetic complications due to the size and range of D from 5.8 X 10~^ to 9.7 X 1(T7 cm^/sec. D was largest for the lowest concentration at the slowest scan rate, and smallest for the fastest scan at the highest concentration. The function V n 11 was estimated to be 0.40 by comparison with Delahay1s Figure 6.6. See Appendix B for calculations and data.

Considering the solution species available and the difficulties involved with the Class IV reaction scheme, the original catalytic concept of the PtO surface still ex­plains the system most satisfactorily. It is not possible to differentiate between a reduction involving the surface oxide or one proceeding through the oxide with the data presented.

CHAPTER 5

ROTATING- DISK VOLTAMMETRY

Rotating disk voltammetry was performed using the Heath polarographic equipment, rotating disk electrode (RDE) and solutions as previously mentioned. Electrode pre treatment was by oxidation as in the previous experiment. Voltammograms, l/E curves, for the system at three poten­tial scan rates are shown in Figures 14, 15 and 16. The character of the dependency between Xe(VI) and the PtO re­duction was again emphasized. At Xe(VI) concentration of2.1 X 1 or less, the RDE l/E curves look very similar to the stationary electrode l/E curves, Figures 6 and 7. Above 2.1 X 10*"% Xe(VI) concentration, the wave-like char­acter of the curves indicates either strong catalysis by the platinum surface or independent irreversible reduction of Xe(VI), corresponding to a combination of 2 or all of the 3 following reactions postulated for the stationary electrode experiments. Figure 14 tends to support a

Reaction 1 PtO + 2e-^Pt.Reaction 2 Xe(VI) + 6e P^°%- Xe(0)Reaction 3 Xe(VI) + 3Pt*-*-Xe(o) + 3PtO

mixed current-catalytic scheme much better than a mixed current-irreversible reduction scheme. At 5 X 10""% Xe(VI)

46

Scan rates 1.67 X v/sec.

(2) 5.29 X 10~4M Xe(VI)(3) 2.10 X 10'4M Xe(VI)(4) 1.06 X 10~4M Xe(VI)(5) 5.30 X 10-5M Xe(VI)(6) 0.065M H,S0.pa.

+0.65+0.95 +0.85 +0.75 +0.45+1.05Volts vs N. H. E.

Figure 14 Rotating Disk Voltammogram at 1.67 mv./sec.

48

Scan rate: 3.34 X 10'*'’ v/sec.

(1) 1.05 X 10_5K Xe(VI)(2) 5.29 X 10~4M Xe(VI)(3) 2.10 X 10-4M Xe(VI)(4) 1.06 X 10-4K Xe(Vl)(5) 5.30 X 10_5M Xe(VI)(6) 0.065M H2S04

(2)_-

+1.05 +0.85 +0.65 +0.45 +0.25 V

Volts vs N.H.E.

Figure 15 Rotating Disk Voltammogram at 3*34 mv./sec.

Xe(VI)Xe(VI)Xe(VI)Xe(VI)Xe(Vl)

(4) 1.06 X 10~4](5) 5.30 X 10~5:(6) 0.065M H„S0

(3)2 -

(4) (6)

+0.75 +0.25 V vs N.H.E.+1.25Scan rate: 8.35 X 10- v/sec.

Figure 16 Rotating Disk Voltammograin at 8.35 mv./sec.

50concentration, the peak current was 0.45)ia, only 0.1 pa higher than the PtO reduction current, and the current quickly decayed to a sustained low of 0.25pa. At 5.3 X 1 Xe(VI), the current continued to rise after the peak at a much slower rate. Increasing the Xe(VI) concentration to 1.05 X 10“ at the same scan rate ie., the same ac­celeration of PtO-VPt rate constant, produced a curve with an even higher slope after the peak. Thus the current after the peak was apparently limited by a platinum-Xe(VI) reaction. The fraction of Xe(Vl) reduced directly by PtO electron bridging and that reduced by formation of PtO can only be speculated on. Small hydrogen adsorption waves near +0.25 in Figure 16 are not to be considered in the wave analysis, as they are a function of the supporting electrolyte.

Increasing the scan rate as in Figure 15 and 16, in­creased ip in a manner that the amount of Xe(VI) reduced at Ep seemed to increase. However, analysis of Figure 17, ip and i ^ ^ vs scan rate as a function of concentration, shows that it is the PtO reduction current that increases rapidly with scan rate, and that for most cases Xe(VI) re­duction currents are not affected by scan rate, as would be expected from a 100# convection controlled reduction.Fried and Elving (1965) discussed the effects of scan rate and rotation rate upon the reduction current. In the two extreme boundary cases that can exist in rotating disk

pa

9*0

8.0 -

7.0 -

6.0 ..

5.0 -

4.0 -

5.0 -

2.0 -

1 .0 -

i-p VS V

(1)

pa

9.0-

8.0

7.0

6.0

5.0

4.0

5.0

2.0 -

1 .0 -

Him vs vtil

(1) 1.05 X 10“5M Xe(VT)(2) 5.29 X 10~^I Xe(VI)(3) 2.10 X 10”4!-! Xe(VI)(4) 1.06 X 10-4M Xe(VI)(5) 5.50 X 10-5h Xe(VI)(6) 0.065M H2SO4

(2)

0.1 0.2 0 .5 0 .4 0.5

Figure 17 Ip and H i m vs Scan Rate.

0.1 0.2 0.5 0.4 0.5

Scan rate in volts per minute(See Appendix C for data.) VJl

52voltammetry: 1) a fast scan rate and slow rotation rateand 2) a high rotational rate and slow scan rate, control of the reaction will be by diffusion and convection respec­tively. In the case of pure convection control, scan rate will not affect the limiting current as indicated by the Levich equation: ^ i m ~ • 620nFAD2/ w 1 /^Cv”1 , where v isthe kinematic viscosity and w, the rotation rate. Under pure diffusion control, the voltammograms will have peaks as in stationary electrode voltammetry and the current willobey the Randles-Sevick equation: i = 3.01 X 10^n(ccn )^ ^P &1 /? 1/?AD ' Cv ' , where v is the rate of polarization. With the very low scan rates and high rotational speed (600RPM) all curves in this study should be controlled purely by con­vection as indicated by the study of Fried and Elving(1965). However, the higher concentration curves in Figure 17 show a definite distortion at the slower scan rates, indicating some form of kinetic complication.

The variation of ilim vs v as a function of Xe(VI) concentration, also in Figure 17, shows that there are some kinetic complications at +0.25 V. The slope of all, but the 5.3 X 10” Xe(VI) limiting current curve, indicates a 10-12% increase in current with scan rate in a region where convection control should be 100%. On these grounds the PtO assisted reduction mechanism is still the most likely mode of Xe(VI) reduction. A simple irreversible reduction may still be possible but the dependence of initial Xe(VI)

55reduction upon the platinum surface reduction, and the rise in current after the peak at high Xe(VI) concentrations all lead to a surface catalyzed reaction. Replotting the data of Figure 17, one can obtain i and i1 iTn vs C as a function of scan rate, Figure 18. is shown to be linear inconcentration at all scan rates except below 1 X 10~^M. 1^,however, shows a 25$ variation in current at 1.05 X 10"" strength, as a function of scan rate. The change in slope of ip/C at 5.5 X 1 0 " " Xe(VI) concentration illustrates the behavior noted in Figures 14, 15 and 16.

Assuming that the three hypothetical reactions, mentioned earlier, control the current throughout the range of the experiment, one can explain the curve shape as for the stationary electrode voltammetry. The reactions are simplifications for the sake of the discussion only. The exact species present on the platinum surface is not known and thus the stoichiometric involvement of Ft and Xe(VI) can only be guessed at. At +0.60 V reaction 1 is current limiting since it must occur before Xe(VI) can be reduced. This reaction is known to be irreversible and increase in rate with increasing cathodic potential. Reaction 2 is related to reaction 1 in that Xe(VI) is reduced only at potentials where PtO is being reduced. Thus the current increases at with Xe(VI) concentrations. Active reduc­tion of Xe(VI) via charge transfer bridging is not a proven reaction but hypothesized by Anson (1959) for the

vs Clim9.0- 9.0"

8.0"

7.0- 7.0"

(1) 8.35 X 10"° v/sec(2) 3.34 X 10-5 v/gec (5) 1.67 X 10”5 v/sec (4) 8.35 X 10-4 v/sec

62 4 8 10 62 4 8 10X 10-4%4 Xe(VI) X 10“^ Xe(VI)

Figure 18 I and Ilim vs C. (See Appendix C for data.) VI4*

55electro-reduction of I0^~ on platinum surfaces. This hypothesis is supported by the following reasoning. As the potential proceeds cathodically, the current of the

curve decreases to a non-Faradic level near zero. In the presence of 5 X Xe(VI) the current decreases toa level where Xe(VI) concentration limits the current flow via reaction 3. Increasing to 2.1 X 1Cf^M, the current did not decay because the rate of reaction 3 was fast enough to maintain i despite the increase in PtO reduction rate. Reaction 3 is assumed not to be potential dependent outside the hydrogen adsorption region. Increasing Xe(VI) concen­tration to 5.29 X 10™^M produces a curve with a plateau at +0.60 V where the current is limited by a combination of reactions 1 and 2, depending on the concentrations. How­ever, as reaction 1 increases, the current increases, meaning that reaction 3 was actually greater than reactions 1 and 2, thus maintaining the electrode in an oxidized state. At 1.03 X 1 0 * " Xe(VI), current continues to rise beyond the plateau region of +0.35 V showing that reaction 1 is current limiting. Doubling the scan rate increases i for PtO reduction, but does not increase the corrected peak current of the Xe(VI) curves, except at 1.03 X 10"* M. A further increase to 5 times the scan rate did not change the peak current due to Xe(VI) except at 1.03 X 10” but did cause a peak to show on all cuves from the increase in PtO current (Figure 17). The limiting currents for

565.29 X 10-4M and 1.05 X 10“5M Xe(VI), however, did not change significantly as seen in Figure 17.

The wave shape can also be simply explained as a separate charge transfer reaction, not related to the PtO reduction. The possibility exists that only reactions 1 and 2 are involved producing the peak at +0.60 V due to PtO reduction and the higher limiting currents at +0.35 V from the Xe(VI) reduction. The effective separation of the E^'s would have to be only 0.050 V to produce this type of a wave. The validity of this assumption can be checked by replotting Figure 15 to show l/E curves without surface related currents, Figure 19. Corrected curves appear to be simple irreversible charge transfers; and since Figure 18 showed i^im vs ^ was linear for all scan rates, use of the Levich equation should be accurate to within 10 to 15# at +0.25 V.

Diffusion coefficients calculated for the rotating disk via the Levich equation were from 2.0 to 9.5 X 10"^

ocm /sec. at +0.25 V, comparable to the values for the DME polarography. Variation of D with scan rate and concentra­tion indicates a kinetic involvement. The rotating disk electrode is ideally suited for a kinetic study. However, to investigate a kinetically coupled reaction as postulated vide infra, variation of the rate of rotation is also necessary. Unfortunately, this type of equipment was not

57

(1) 1.05 X 10"% Xe(VI)(2) 5.29 X 10~4M Xe(Vl)(3) 2.10 X 10“4M Xe(Vl)(4) 1.06 X 10-4M Xe(VI)(5) 5.30 X 10"% Xe(VI)

10..

8..

(2)uncorrected

peak currents4..

2 -

i4)

+0.65+0.85+1.05 +0.45+1.25 +0.25 V

Figure 19 RDS Voltammogram Corrected forBackground.

58available at the time of the study. See Appendix C for data and sample calculations.

In preliminary experiments it was noted that a platinum electrode which had been exposed to iodine pro­duced a differently shaped curve for the Xe(VI) reduction.To investigate this anomaly and provide more understanding of the Xe(VI)-platinum reduction, a new set of experiments were performed using the RDE. The upper curve of Figure 20 is a cathodic sweep from +1.25 V in 1 X 10“ Xe(VI) with an oxidized RDB that had been dipped in a solution con- taining 2.5 X 10"% I2, and 0.05M KI in 0.065M HgSO^. Electrode pretreatment consisted of rinsing with doubly distilled water, immersion in the test solution, and being held at +1.25 V for one minute with ^ bubbling through the cell. The effects of exposure to iodine were so lasting that to return the electrode to a clean state, harsh meas­ures were required. Standing over night in concentrated HNO^, electrolyzing in the cell and evolving oxygen for 10 to 15 minutes or regrinding the surface would refresh the electrode. Repeated cycling for +1.25 V to +0.25 V in Xe(VI) solution would also gradually remove the wave depression.

Curve 2 represents the addition of 1.0 ml of 1 X 10 KI in 0.065M I^SO^. Note the dramatic depression of the wave. Xe(Vl) reduction via PtO bridging was greatly reduced. The wave beginning at +0.25 V was dependent upon

- 8.0

-4.0-

jiaKI

-2.0-

0.0.+0.05 V+0.25+0.45+0.85

Volts vs N.H.E.

Scan rate: 3.54 X 10""5 v/sec+2.0

Figure 20 I/B Curves Showing the Effects of I~ and I2.

60Xe(Vl) concentration as seen by addition of another milli­liter of 1 X 10*" M KI. Xe(VI) concentration in Curve 2 was8.1 X 10"" M, with a limiting current of 5• 2p.a. Assuming the wave depression to be due to lowering the Xe(Vl) con­centration alone is not valid because Figure 15 shows that at 5.29 X 10~^M Xe(Vl) concentration the current continues to rise after +0.60 V. The iodide-iodine redox couple changed the base line for the Xe(VI) reduction and sup­pressed the PtO reduction peak to a nominal level of 0.92pa as per Figure 15. Adsorption of electro-active and electro­inactive iodide and iodine have been shown to exist on a clean platinum surface by Hubbard, Osteryoung and Anson (1966). Will and Knorr (1960) also noted that oxide for­mation was hindered by iodide adsorption. Thus suppression of the enhancement of the PtO reduction via a coupled Xe(VI) reduction can be explained via previously noted experimental results.

The wave beginning at +0.35 V in Figure 20, appears to be Xe(Vl) reduction per se. Treating it as a simple irreversible reduction, D was calculated at all concentra­tions to yield the values 7.5, 6.2 and 5.8 X 10** cm2/sec.It was assumed that enough suppression occurred after +0.60 V to separate any coupled reactions between PtO and Xe(VI) at +0.25 V even though there was peak enhancement at 1 X 10~^M. The enhancement of the PtO peak was totally sup­pressed by 2pl-l of I*". The diffusion coefficient was

61calculated for other curves not shown and the average value of D was 7.5 X 10"* cm^/sec. To all appearances, the dif- fusion coefficient was near 8.0 X 10"* cm /sec. for the wave at +0.25 V without iodide and at +0.05 V in the pres­ence of iodide. The reaction responsible in both cases was probably the irreversible reduction of Xe(Vl). Kinetic involvement of the PtO surface was blocked by and I"" on the electrode following the peak at +0.60 V. The presence of adsorbed was also signified by the more cathodic po­tential of the wave and the decrease in hydrogen adsorption potential from +0.25 V to +0.05 V. Clarification of the amount of PtO surface involvement in the unshielded wave at +0.25 V and the character of the iodine-iodide shielding can only be done by further work in varying rotational rates and by chronopotentiometric techniques.

Other systems were also briefly investigated with the rotating disk electrode for the purpose of observing catalytic waves. Among them were; Tl+ - Tl+ , Br~^ - Fe+ - Fe+ , and I0^~ - with Xe(VI) and I0 ~". In all cases the reactions either occurred cathodic of the PtO re­duction peak or no current enhancement occurred. Periodate behaved very similarly to Xe(Vl) at an electrode partially covered by I^. This is not difficult to understand con­sidering the oxidation potential and reduction products of periodate. The effects of scan rate at one concentration

62of lO^” can be seen in Figure 21 which relates l/E and I va v for 1 X 10""% KIO. in 0.065M H2S0..

The results of this study also added insight to the experimental results of Anson (1959). It is doubtful that an oxide bridging mechanism is necessary to explain the results. Since iodine and iodide are products of the io- date reduction, it is probable that the adsorption of iodine and iodide covered the platinum as soon as the plat­inum oxide was removed. lodate was reducible at a platinum oxide film that was being reduced and at a clean platinum surface. Adsorption of iodide and iodine covered the electrode causing the current to decrease and even­tually forcing the iodate reaction to shift to more cath- odic potentials. The behavior of Xe(VT) was very similar to iodate, except that Xe(0) does not adsorb on the elec­trode and cause the current to decrease rapidly. When iodide was introduced to the cell, the Xe(VI) curves be­came very similar to the curves. For both andXe(VI), the platinum oxide reduction initiated their reduc­tion, and in the absence of I” and Ig both are probably electro-active on a clean platinum electrode.

Figure 21

CHAPTER 6

SUMMARY

The dropping mercury electrode polarography of Xe(VI) in 0.065M has been shown to be similar to thepolarography in 0.1M K2SO^ except in the rate of half-wave potential shift with pH. In this study changing the sup­porting electrolyte to vary the pH resulted in an apparent E^/pH of 0.040 V/pH unit instead of a 0.077 V/pH unit change, as noted by Jaselskis (1964a) in varying buffers at a constant 0.1M K2S0^ supporting electrolyte. "I" for Xe(VI) as determined for this study and Jaselskis1 (1964a) was 11.9 average. Jaselskis (1964b) reported a lower value (7.7) which cannot be explained.

The Xe(Vl) reduction at a platinum electrode has an entirely different characteristic. Immediately obvious in both the stationary and rotating electode work was that Xe(VI) was being reduced at platinum oxide reduction poten­tials. Peak currents were a function of Xe(VI) concen­tration and scan rate. Continuation of the reduction current after the peak was more of a function of Xe(VI) concentration than of scan rate in both experiments. The stationary electrode results indicated that a three reactionmechanism was likely. Xe(VI) reduction via the activated

64

65platinum oxide surface takes place at and about E^. PtO re­moval continues at a faster rate as the potential decreases; and since the current at +0.25 V forms a plateau with in­creasing concentration, re-oxidation of the surface must be occurring. The possibility of a simple irreversible Xe(VI) reduction on a clean platinum electrode seems un­likely from the data and the analysis of Nicholson and Shain (1964).

The rotating disk electrode provided a different view of the system. The platinum surface film was defi­nitely responsible for Xe(Vl) reduction at +0.60 V. I for Xe(Vl) was not completely insensitive to scan rate, re­flecting the uncertainty of the reaction parameters. The small variation of i ^ m with respect to scan rate, indi­cates that the concomitant influence of the platinum surface and related possible catalysis in the plateau re­gion is secondary to the simple irreversible reduction.The apparent peak enhancement talked about in Chapters 4 and 5 was due to adding the PtO reduction current on to the Xe(VI) reduction current in the rising portion of the wave. Thus causing the PtO peak to grow with Xe(VI) concentration. The reduction of the platinum oxide undoubtedly initiated the Xe(Vl) reduction. Exposure of the electrode to iodine eventually suppressed all catalysis at +0.60 V resulting in a Xe(VI) reduction starting at +0.35 V. Because iodine

66adsorption shifted effective potentials at the electrode $, kinetic analysis of the wave does not seem profitable. Analyzing the resulting waves for D yielded a value close to that of the wave at +0*25 V.

A more thorough extension of this studyp involving varying the mass transport rate along with scan rate and concentration,' should elucidate the relationship of Xe(VI) reduction and the platinum surface. Analysis of iodate and periodate in presence of iodine.should also be carried out for comparison purposes.

The stated goals of the study, to look for a cat­alytic usage of Xe(VI) were not met in the manner expected. Electrode involvement was not of interest, but to this end new insights might be provided by the. use of tin oxide or graphite electrodes. In striving for the useful, this study may have opened the possibility of iodine trace analysis via adsorption blocking of a catalytic reduction. For practical purposes, however, a more inexpensive oxidizer than Xe(VI) will have to be found.

APPENDIX A

POLAROGRAPHY

The Ilkovic Equation: id = 607nD1^ C m 2'3*1/6

Where: i, = average diffusion limited current inmicroamperes.

n = number of electrons involved in the reaction.D ss diffusion coefficient of the reducible ion

in cm /sec.C ss concentration of reducible ion in millimoles

per liter.m = mass of mercury flow in mg/sec.t = drop time in seconds.

Sample calculations for an average current, determination:C = 5.3 X 10~4M Xe(VT) = 5.3 Xm = 0.1485g for 20 drops or 1.516 mg/sec.t = 4.9 sec.i, = 10.0 microamperes.

Derivation of and V log m = 0.18070 log t = 0.69020

2(log m) = 0.36140 log t/6 = 0.115052/3(log m) = 0.12046 antilog 0.11503 = 1.303

antilog 0.12046 = 1.32

10.0 = 607(6)D1/,2(5.3 X 10_1) (1.32) (1 .303)67

68

"I" the diffusion current constant = i./cm^^t^^

10'°__________ = 10,0 = 10.96(1.30) 5.3 X 10~1 (1.32) 0.912

I 2 10.96 2 , ,D — (----) = (---------?) = 9.05 X 10 cm /sec.

607n 3.64 X 10^

Table 2. Results of Polarofcraphy.

69

July 14, 1966Cone. Xe(VT) 10i6 5.

(10-4%)

mass (s) 0.1557 0.20 drops

m2/ 3 1 . 3 6 1.

id(p.a) ave. 15.8 8 .I 8.5 10.

D X 10"6 c C Qsq. cm./sec.

July 21, 1966Cone. Xe(VT) 5,

(10-4%)mass 0.1473 0.20 dropsm2/3 1.31 1.

id()ia) ave. 19.0 10.I 10.6 11.

D X 10-6 sq. cm./sec.

2.5 1 .01 0.51

0.1549 0.1538 0.1577

1.36 1.35 1.375.4 1.65 1.048.2 9.4 11.4

5.2 6.7 9.8

Cone. Xe(YI) 10e5 5.29 2.10(10-4%)

1.06 0.53

mass 0.1473 0.1485 0.146220 drops 0.1477 0.1433

m2/3 1.31 1.32 1.31 1.31 1.29id()j.a) ave. 19.0 10.0 3.9 2.0 0.94

I 10.6 11.0 10.9 10.9 10.6

8.4 9.1 8.9 9.1 8.5

2

1543

35491

29

1485

3200

1

70

APPENDIX B

STATIONARY ELECTRODE VOLTAMMETRY

Table 3. Data for Figure 9«_________________________Stationary Electrode Parameters at in microamperes

Cone.Xe(VI) 3.0Scan Rate in mv/ 9.3 18.7

sec.45.7

0.065M HgSO. 1.66 2.84 6.8 16.55.3 X 10“5M 0.68 1.92 — 6.51.06 X 10~4M 2.06 3.96 5.2 8.02.10 X 10~4M 5.22 9.61 11.2 21.05.29 X 10-4M 13.7 20.4 29.0 47.51.05 X 10~5M 25.7

microamperes145.0

Stationary Electrode Paramaters at +0.167 V in microamperesCone.Xe(VI) 3.0 Scan Rate in mv/sec.

9.3 18.7 45.7 91 .40.065M H2S0. — 0.2 0.5 1.3 3.05.3 X 10“5M 1.0 1.5 1.9 2.5 3.0

1.06 X 10-4M 2.4 2.9 3.6 4.7 6.82.10 X 10-4M 4.7 5.4 6.5 8.0 11.75.29 X 10-4M 12.8 12.6 14.0 20.0 21 .81.05 X 10-5M 18.0 20.0 25.0 32.0 54.0

microamperes

71

Table 4. Data for Pi^ire 10.

log iCone.

Xe(VI) 0.477log V

0.968 1.271 1.6590.065M KgSO. 0.22 0.45 0.83 1.231.06 X 10"4M 0.51 0.60 0.72 0.902.10 X 10-4M 0.72 0.98 1.05 1.325.29 X 10_4M 1.14 1.31 1.46 1.68

log 1+0.167 V

Xe(VI) 0.477 0.968J L U & V1.271 1.659 1.961

0.065M HgSO. ----- ----- -0.30 0.11 0.485.3 X 10-5M 0.00 0.18 0.28 0.40 0.481.06 X 10~4M 0.38 0.46 0.56 0.67 0.832.10 X 10-4M 0.67 0.73 0.81 0.90 1.075.29 X 10-4M 1.11 1 .10 1.15 1.30 1.341.05 X 10"5M 1 .26 1.30 1.40 1.51 1.73

•' - 72Table 5. Data for Figure 1 1 . ____________________________Peak current function; i^/v^ ^ X 10”^

Cone.Xe(VI) 3.0

Scan Rate in mv/sec.9.3 18.7 45.7 91.4

0.065M H2S0. 3.04 3.0 3.1 7.7 8.25.3 X 10-5M 1.2 2.0 --- 3.0 1.321.06 X 10_4M 3.8 4.2 3.9 3.74 5.32.10 X 10-4M 9.5 9.6 8.4 9.8 12.05.29 X 10-4M 25.0 21.4 20.6 22.0 ---

1.05 X 10_5M 46.8 --- •--- 108.0 ---

v 1/2 0.055 0.095 0.13 0.21 0.30

Limiting current function : ilinA 1/2 X 10"•5

Cone. Scan Rate in mv/sec.Xe(VI) 3.0 9.3 18.7 45.7 91.4

0.065M H2S0. --- 0.21 0.37 0.60 0.985.3 X 10“% 1 .8 1.58 1.4 1.17 0.981.06 X 10"4M 4.4 3.1 2.8 2.24 2.22.10 X 10-4M 8.5 5.2 4.9 3.7 3.85.29 X 10~4M 23.0 12.0 10.5 9.3 7.31.05 X 10“% 34.0 30.6 18.5 14.7 17.0

73Notes for Figure 12

The theoretical curve in Figure 12 was calculated using the following current functions from Nicholson and Shain (1964).i = i(planar) + i(spherical)i = 602n(<%n^)^A(Dv)^Cv5rX(bt) + 0.16O(v /roi/Sfn^)0(bt)C = 5.29 X 10-4M, v = 18.7 X 10“3 v/aec., n = 6, ccna = 0.40, r = 0.1 cm.

Consult Nicholson and Shain for the exact descrip­tion of the potential scale. Zero in this case is 5 mv. cathodic of E . Below is a table relating current and potential.

Table 6. Data for Figure 12________________________________millivolts i (ua) millivolts i (ua)

+160 1 .48 +20 201140 3.95 15 216120 7.91 10 228110 11 .9 5 237100 17.6 0 24290 24.7 - 5 24580 36.1 10 24470 51.4 15 24060 71.6 20 23350 98.4 30 21840 130.5 40 20135 148.0 50 18530 166.5 70 15925 184.0

Table 7. Data for Figure 13.Peak potentials and shifts with scan rate

Cone.Xe(VT)

Scan Rate in mv/sec. 9.3 AE 18.7 AE 45.7 A E 91.4

0.065K HgSO. 0.560 -.040 0.320 -.006 0.314 +.031 0.3455.3 X 10-5M 0.354 -.006 0.348 -.088 0.260 +.085 0.345

1.06 X 10-4M 0.295 -.003 0.292 -.040 0.252 +.024 0.2762.10 X 10-4M 0.327 -.002 0.325 -.047 0.287 +.046 0.3245.29 X 10-4M 0.300 +.010 0.310 -.058 0.252 +.052 0.304

log V 0.3034

A E/Alog v

0.3881 0.3010

Cone. Scan Rate Differentials in mv/sec. Xe(TX) 9.3 - 18.7 18.7-45.7 45.7 - 91.4

0.065M H2S0. -0.132 -0.015 +0.1035.3 X 10~5M -0.020 -0.226 +0.2821.06 X 10~4M -0.010 -0.103 +0.0802.10 X 10-4M -0.007 -0.120 +0.1535.29 X 10-4M +0.033 -0.150 +0.173 -j

75

Table 8,Value of D Related to Scan Rate and Concentration

Cone. Moles Xe(VI)

5.3 X 10-5 1.06 X 10~45.29 X 10"41.05 X 10-3

Scan Rate in mv/sec.3.0 91.4

4.0 X 10 5.8 X 105.4 X 103.3 X 10

-6-6- 6

-6

1.2 X 10 1 .68 X 106.3 X 10

9.75 X 10

-6-6-7-7

Calculation of the Diffusion Coefficient Equation 6-29 (Delahay, 1954):

i = 3.01 X 105n(«na)*/lD*Cv*ip = peak current in microamperes.n = number of electrons involved in total reductionn = transfer coefficient times the number of

electrons in rate limiting step.pA = area of electrode in cm .

oD = diffusion coefficient in cm /sec.C ss concentration in millimoles per liter, v = scan rate in volts per second.

Sample calculation:C = 2.1 X 10-4M Xe(VI), v = 18.7 mv/sec., i = 6.5 )ia

6.5 = 3.01 X 105(6)(.40)*(0.15)D*(.21)(1.82 X 10~2)*6.5D2 =

3.01 X 105(6)(.63)(0.15)(.21)(134)D = 1.8 X 10“6 cm2/sec.

= 1.34 X 10-5

APPENDIX C

ROTATING DISK VOLTAMMETRY

Table 9* Data for Pimires 17 and 18.-—- - ---.-- ~--C. 1I (corrected for PtO currents) :in microamperes

Cone. 0.835 1 .67 3.34 8.35 rav/sec.Xe(Vl) 0.05 0.10 0.20 0.50 volts/min.

0.06514 h2s°4 0.32 0.62 0.95 2.125.3 X 10-5m 0.105 0.15 0.10 0.211.06 X io"4m 0.25 0.45 0.33 0.322.10 X 10~4I-1 0.98 1.0 1.05 1 .15.29 X 10'"4!’! 2.38 2.8 2.95 3.11.05 X 10_3M 6.2 5.4 7.05 7.1

'Cliin at (0.25 V) in microamperesCone. Scan Rate in volts/min.

Xe(VI) 0.05 0.10 0.20 0.500.065H H2S04 ——— 0.08 0.06 0.145.3 X 10-5M 0.27 0.22 0.34 0.381 .06 X 1 o"4h 0.33 0.33 0.38 0.762.10 X 10_4H 1.35 1.53 1.66 1.865.29 X 1 o-4k 4.2 4.10 4.34 4.411 .05 X 10-% 8.65 8.90 9.12 9.16

76

77

Calculation of D usinff the Levich EquationAssuming convective control,

ilim = 0.620n?AX)2/5v-1//6w1'/2C^Vhere: n = number of electrons involved in the reaction

P = Faraday constant, 96,493 coulombs/equiv.2A = area in cm

2D = diffusion coefficient in cm /sec.v = kinematic viscosity in stokesw = angular velocity in radians/sec.

ilim = limiting current in amperesC = concentration of reactant in moles

per milliliter.

A = 0.025" diameter or 3.14 X 10*" cm^ w = 600 RPM = 62.8 R/sec.

Calculation of kinematic viscosity Density of 0.130N H^SO^: 1.0042 (estimated)Vg q = 0.8937 centipoise

Sp. vH^go^(0.130n) = 1.0065

v = x 10-2 = 8.92 x 10-5 stokes

V 1/6 = (8.92 X 10-5)1/6 = 0.456

78

For 5.29 X 10_4M Xe(VI), llim = 4.2pa

2/5 (0.456)(4.2 X 10“6)D

(0.620)(9.65 X 104)(3.14 X 10“5)(5.29 X 10“7)(7.93)3/2

D = (4.06 X 10-4) " 1" = 8.2 X 10“6 cm2/aec.

10. Variation of D with Scan Rate and Concentrat

Cone. Scan Rate in mv/sec.Xe(VI) 0.84 1 .67 3.34 8.35

5.3 X 1 o-5m ---------- ---------- 6.0 7.01.06 X io_4m 2.0 2.0 ---------- 7.22.10 X io“4m 6.0 7.2 8.1 9.65.29 X io-4m 8.2 ---------- 8.6 8.81.05 X io-5m 8.6 9.0 — — 9.5

"D" reported in sq. cm./sec. X 10

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