AEAFQNS CENTER SILVER SPRING MO F/A 7/4 ELECTROCHEMISTRY … · -, ELECTROCHEMISTRY OF SULFUR DIOXIDE IN NONAQUEOUS SOLUTIONS, PART I BY LINDA S. LAUGHLIN RESEARCH AND TECHNOLOGY

A0-A113 033 NAVAL SURFACE AEAFQNS CENTER SILVER SPRING MO F/A 7/4ELECTROCHEMISTRY OF SULFUR DIOXIDE IN NONAGUEOUS SOLUTIONS. PAR--ETCIUI)MAY Al L S LAUGHLIN

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NATIONAL BUREAU OF STANDARDS 1963-A

NSWC TR 80-533

-, ELECTROCHEMISTRY OF SULFUR DIOXIDEIN NONAQUEOUS SOLUTIONS, PART I

BY LINDA S. LAUGHLIN

RESEARCH AND TECHNOLOGY DEPARTMENT

18 MAY 1981

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4.TTZfand Subtitle) II. TYPEL OF RE9PORT a 191100 CovOV119

ELECTROCHEMISTRY OF SULFUR DIOXIDE IN NONAQUEQUS IteiSOLUTIONS, PART I ____________

S. gPEORMING onG. REPORT NurmaER

7. AUTNORt.) *N 0 A U

Linda S. Laughlin

S. PER1FORMING ORGANIZATION NAMEt AND ADORESS 10. 1RGAM . .P11133C. TASKNaval Surface Weapons Center (Code R33) AE OI N' UU61152N; ZROOOl ; ZROl301;White Oak ROIAA071Silver Spring, MDl 20910

11. CONTROLLING OFFICE NAME AND AOORESS IL RUPORT DATE

18- ILA, 198113. NUWOER OF PAGECS

42I&. MONITORING AGENCY NAME ADORESS(H1 eeiUnt barn CoMUeflb Office) I&. SECURITY "LASS, (ofthlia sp

Unclassified

11114 gLCNk~UrICATION/ DOWNGRADING

IS. OISTRIS111UTION STATEMENT (44 Moo ftepet)

Approved for public release; distribution unlimited

'I. OISTRIUUTIONd STATEMENT (of00 SOM an ft Wham" Se0 if offerut itat Repea)

IS. SUPPLEMENTARY NOTES

19. KEYV woRDS (Gadw e new"e WOd 19 semeeMp 400 N1100114 4W Wek -mb)

Sulfur DioxideCyclic Voltaietry

Konaqueous Solvents

2L. ABSTRACT (CMONSIin Ovine side II aeto"e ANUM ' da

A survey and discussion of published reports on the reduction of sulfurdioxide in nonaqueous solutions is presented as an introduction to theexperimental plans and preliminary results contained in Part II.

DD I C'1I 1473 EDITION OF I Nov soels OUoEtcr UNCLSSIFIED

blNSWC TrR 80-533

" - I

FOREWORD

A survey of the literature on the reduction of sulfur dioxide in nonaqueous

solvents was carried out as part of a program to investigate safety hazards in

nonaqueous ambient temperature lithium batteries. Comparison and discussion of

conclusions and contradictions presented have led to the experimental plans and

results in part II of this report.

This work was funded by the Independent Research Program of the Naval

Surface Weapons Center and by the Electrochemistry Technology Block Program of

NAVSEA.

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CONTENTS

Chapter page

1. INTRODUCTION ........... ... * . ........*..... 7

2 DIMETNYLFORMAHIDE (DM1)................................ 9

-3 DIIEETHYLSULFOZIDE (DMSO) ... ................. 13

4 ACETONITRILE (AN) o.......... .o......... ... .17

5 PROPYLENE CARBONATE (PC) ............ ... 19

6 OTHER SOLVENTS ...... ... oo. .o .... .. .. 21

8 SUMMARY AND PLANS ............ .. . ........... 27

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TABLES

Table Page

1 PUBLISHED STUDIES OF SO2 REDUCTION ......... 292 CYCLIC VOLTAMMETRY STUDIES ........................ * . 303 OTHER VOLTAMMETRIC STUDIES .................... ............... 31

4 ASSIGNMENT OF ANODIC PEAKS IN CYCLIC VOLTAIMETRY INTETRAALKYLANMONIUM SALT SOLUTIONS ........................... 32

5 ASSIGNMENT OF ANODIC PEAKS IN CYCLIC VOLTAMMETRY IN ALKALISALT SOLUTIONS ......... % .... *.................................. 33

6 DIELECTRIC CONSTANTS (e) AND DIPOLE MOMENTS (5) OFNONAQUEOUS SOLVENTS ......................................... 34

7 EQUILIBRIUM CONSTANT FOR COMPLEX FORMATION .................... 358 VARIATION OF El/ 2 WITH [SO 2) IN DMP-TEAP ...................... 369 VARIATION OF El/ 2 WITH TEMPERATURE ............................ 37

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Chapter 1

INTRODUCT ION

Considerable interest in the electrochemical behavior of nonaqueoussolutions of sulfur dioxide has been generated by the use of these systems inhigh energy density lithium batteries. During the past twenty years, SO2 insolvents such as dimethylformamide (DNF) and dimethylsulfoxide (DMSO) has beenextensively studied. Only since 1978 have reports been appearing on actualbattery solvents, propylene carbonate (PC) and acetonitrile (AN).

The reduction of SO2 has been studied by a variety of techniques,including ultraviolet (Uv) and raman spectroscopy, electron spin resonance (ESRor EPR), controlled potential electrolysis, coulometry, chronopotentiometry,voltammetry and chemical reduction. There has been little consistency as tosolvent, supporting electrolyte, working electrode surface or referenceelectrode, resulting in a patchwork of data. Published reports have beensummarized in Table 1.

Whether the electron is transferred chemically or electrochemically, thefirst product is S02-. Controversy begins over subsequent chemicalprocesses, proposed to be either

SO2 + x S02 (SO2)xSO2 (1)

or

2So2- S2o42- (2)

Voltaumetric and spectroscopic data have been interpreted to support one or theother pathway. Only recently have the many contributing factors beeninvestigated. A review of the various studies should prove useful in theplanning of future work. These have been grouped into cyclic voltanmetricstudies, Table 2, and other voltammetric studies, Table 3.

The experimental work performed in this laboratory will be presented inPart II of this report.

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Chapter 2

DIMETHYLFORMAMIDE (DMF)

I. CHEMICAL REDUCTION

Rinker and Lynn reduced SO2 in several solvents, using sodium amalgam asthe reductant, producing several species. 1 A blue radical ion obtained in DMFwas identified as SO2SO2-(x=1), as well as a white precipitate which wasprimarily sodium dithionite, Na2 S204.

II. ELECTROCHEMICAL INVESTIGATION

A. TETRAALKYLAMMONIUM SALTS.

1. Platinum Electrode. Dinse and Mobius electrolytically generatedSO2- in DMF-TPAP (tetrapropylamnonium perchlorate) solution, studying theresulting species by EPR and UV spectroscopy. 2 From analysis of the changesin the EPR spectra with temperature, it was determined that for the complex,(SO2)xSO2-, x-2 and the formation constant is l.3x10

4 M-2 at

0°C. A blue-green color was observed at higher SO2 concentrations.

Martin and Sawyer found not only a blue species in DMF-TEAP* -(tetraethylammonium perchlorate), but a red species as well.3 %hen NaC104

was added to a completely electrolysed solution of SO2, sodium dithioniteprecipitated. Cyclic voltammetry resulted in four peaks, two cathodic and twoanodic. IN spectroscopy revealed two peaks, at 485nm and 580unm. One electronreduction of SO2 occurred at -0.85V (vs. SCE), with the correspondingoxidation of S02- at -0.74V. The UV absorption at 580nm was identifiedwith the blue complex, (SO2)xSO2-, which was oxidized at -O.24V. Thered complex, SO2S2042-, absorbed at 485nm and was reduced at -1.65V.The absorption at 350nm of a completely electrolysed solution of S02 wasattributed to the presence of dithionite ion, $2042. The equilibriumconstants were calculated to be 24 Wl for eqn. 2 and 10 for eqn. 3.

1lRinker, R. G. and Lynn, S., "Formation of Sodium Dithionite from SodiumAmalgam and Sulfur Dioxide in Nonaqueous Media," Industrial and EngineeringChemistry: Product Research and Development, Vol. 8, 1969, pp. 338347.

2Dinse, K. P. and Mobius, K., "EPR - Untersuchungen an elektrolytischerzeugtem SO2-," Zeitschrift fur Naturforach A, Vol. 23, 1968,pp. 695-702.

3 Martin, R. P. and Sawyer, D. T., "Electrochemical Reduction of Sulfur Dioxidein Dimuthylformamide," Inorganic Chemistry, Vol. 11, 1972, pp. 2644-2647.

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S2 + S2042- S02S2042- (3)

Magno, Hazzochin and Bontempelli found that SO2 was reduced at -0.95V(vs. aq. SCE) in DMF-TBAP(tetrabutylammonium perchlorate),4 resulting inanodic peaks at -0.65V and +O.IV. The first peak was attributed to theoxidation of S02- resulting from the dissociation of dithionite and thesecond peak to the direct oxidation of dithionite itself. The blue solutionabsorbed at 580nm and analysis of the spectra indicated that the complex,(SO2)xS02-, was formed. No correlation was observed between the peakat 580nm and any voltammetric peaks. Coulometry showed that S02 and the bluecomplex were both reduced at the SO2 reduction potential. The absorptionspectra were interpreted to suggest that x>l, probably equal to 2. SO2absorbed at 275nm and a saturated solution of Na2S204 at 297nm.Completely reduced solutions of SO2 also absorbed at 297nm. The positiveshift of the anodic peak potentials and the increase in the second anodic peakat lower temperatures were assumed to result from weak adsorption of dithioniteonto the platinum surface.

Kastening and Gostisa-Mihelcic investigated complex formation inDMF-TEABr(tetraethylammonium bromide) by ESR in order to determine the correctvalue of x.5 SO2 was partially electrolysed and the ratio of the tworadical species determined, leading to the conclusion that xl and theequilibrium constant for complex formation (eqn. 1) is 230 -1. The spectraldata reported by Dinse and Mobius were recalculated for xl (rather than 2)2and K was found to be 65-200 M-1 , depending on S02 concentration.

Bowden and Dey found two anodic peaks in the cyclic voltamograme of SO2in DMF-TBAPF6 (tetrabutyliumonium hexafluorophosphate)

6 at -0.13V and +0.63V(vs. AgCl coated Ag wire), which were assigned to the oxidation of S02- and62042-.

Fouchard observed that the electrolysis of S02 in DNF-TEAP graduallyturned the solution from bright blue to red to colorless. 7 Raman spectra of a

4%asno, F., Maszochin, C. A. and Bontempelli, G., 'Voltammetric Behavior ofSulphur Dioxide at a Platinum Electrode in Dimethylformamide,"Blectroanalytical Chemistry and Interfacial Electrochemistry, Vol, 57, 1974,pp. 89-96.

5Kastening, B. and Gostisa-Mihelcic, B., "An USR Investigation of the ComplexFormation Between 802 and Electrogenerated S02-," Journal ofElectroanalytical Chemistry, Vol. 100, 1979, pp. 801-808.

2S.e footnote 2 on page 9.

6 owden, W. L. and Dey, A. N., "Primary Li/SOC12 Cells XI. SOC12 ReductionMechanism in a Supporting Electrolyte," Journal of the Electrochemical°i ciety, Vol. 127, 1980, pp. 1419-1426.

7Fouchard, D. T., Gardner, C. L., Adams, W. A. and Laman, F. C., "Roman and ISR

Spectroscopic Studies of the Ulectroreduction of Sulphur Dioxide," inProceedints of the Symposium on Lithium Batteries, Electrochemical SocietyMeeting, Hollywood (FL), October, 1980, pp. 156-158.

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completely electrolysed solution were characteristic of the dithionite ion. ESRspectra were used to estimate the equilibrium constant for eqn. 1 as 600 M1.In addition to the anions already identified, S02- and S02SO2-, athird was proposed:

2 SO2 -.* SO + S032- (4)

2. Mercury Electrode. Gardner et al initiated a study of the effectsof solvent and electrolyte on the distribution of products8 among S02-,SO2 SO2-, and S2042 - . Most of the reported results were obtainedin DMF. Two anodic cyclic voltametry peaks, two ESR peaks and two pulsepolarographic oxidations were observed in DMF-TEAP, which were attributed toS02 - (-1.17V vs. Ag/AgNO 3 ) and SO2 SO2- (-0.65V). The equilibriumconstant for formation of the complex, S02 SO2-, was calculated to be4360 M-1 .

B. ALKALI SALTS.

I. Mercury Electrode. The reduction of SO2 in DMF-LiCl0 4 wasshown by Gardner et al to be almost completely irreversible 8 with an entirelydifferent product, as evidenced by the potential of the single anodic oxidationpeak in cyclic voltammetry (-0.61V). Several alkali perchlorates were studied,with the lithium salt producing dithionite and the potassium salt producing thecomplex, SO2 SO2-.

8 Gardner, C. L., Fouchard, D. T., Leman, F. C. and Fawcett, W. R., "TheKinetics and Mechanism of the Reduction of Sulphur Dioxide in NonaqueousMedia," in Proceedings of the Symposium on Lithium Batteries, ElectrochemicalSociety Meeting, Los Angeles, October, 1979, pp. 545-559.

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Chapter 3

DIMETHYLSULFOXIDE (DMSO)

I* CHEMICAL REDUCTION

Rinker and Lynn's solution of Na(Hg) and S02 in DMSO changed color fromblue to purple to black to violet to red, finally resulting in a whiteprecipitate.1 The formation of the blue radical ion (S02SO2") occurredmuch faster than in DMF. The nature of the red radical ion (RRI) was not clearuntil experiments had been run in formamide (see Chapter 6). The whiteprecipitate, shown to be chiefly sodium dithionite, also formed much faster thanin DMF.

II. ELECTROCHEMICAL INVESTIGATION

A. TETRAETHYLAMKONIUM SALTS.

1. Platinum Electrode. Bonnaterre and Cauquis observed that thereduction potential of S02 in DHSO-TEAP varied with the concentration ofSO2, with a sharp change in the behavior occurring at about 10-2M.9 Adimerization, as written in equation 2, resulting in a coating of dithionite onthe electrode surface, was used to explain anomalous results. The ratio of the

*anodic to cathodic peaks from cyclic voltammetry decreased as the SO2concentration increased, as well as when the scan rate was increased. Thisdecrease followed the predictions of Olmstead, Hamilton and Nicholson for adimerization following electron transfer.10 The change of reduction potentialwith the concentration of SO2 agreed with Laviron's observation that this wascharacteristic of diffusion controlled electron transfer followed bydimerization.1 1 The weak EPR signal attributed to SO2- indicated thatthe equilibrium was largely shifted toward the dimer.

lSee footnote 1 on page 9.

9 onnaterre, R. and Cauquis, G., "Dimerization Consecutive a un TransfertMonoelectronique II. Reduction Electrochinique du Bioxyde de Soufre dane IsDimethyl Sulfoxyde," E lectroanalytical Chemistry and Interfacial

Electrochemistry, Vol. 32, 1971, pp. 215-223.

10Olmatead, M. L., Hamilton, R. G. and Nicholson, t. S., "Theory of CyclicVoltametry for a Dimerization Reaction Initiated Electrochemically,"Analytical Chemistry, Vol. 41, 1969, pp. 260-267.

IlLaviron, E., "Mechanism of the Polarographic Reduction of Aromatic CarbonylDerivatives: Influence of Dimerization of the Free Radical Formed in theFirst Reduction Stage," Collection of Czechoslovak Chemical Communications,Vol. 30, 1965, pp. 4219-4236.

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2. Mercury Electrode. Dehn et al obtained a polarographic wave forS02 in DMSO-TEAP at -0.77V (vs. aq. SCE), 12 observing that the potential ofthe wave depended on the amount of water in the solution. The extensivesolubility of SO2 in DMSO was explained by the formation of the complex

(CH3)2SO + SO2 1' (CH3)2SOS02 (5)

which could be reduced in the following manner:

2[(CH 3)2SOS0 2] + 2e- . 2(CH 3)2S0 + S2042- (6)

B. LITHIUM SALTS.

1. Mercury Electrode. Differential pulse polarography wasinvestigated by Garber and Wilson as a means of analysis of SO2 in air.13Garber's samples were bubbled through DMSO-LiCI, deaerated and analyzed. Thesensitivity was sufficient, with the resultant peak heights proportional toS02 concentration, up to at least millimolar concentrations. Interferenceswere found from nitrogen oxide and with water concentrations greater than 5%.

Bruno et al sought to evaluate Garber and Wilson's method, finding thatnitrogen oxides did not interfere and could be analyzed simultaneously.14 Thehalf-wave potential was -0.875V (vs. calomel in DMSO-O.1M LiC1).

Bruno, Caselli and Traini also studied the cyclic voltammetric behavior ofSO2 in DMSO-LiNO3.

15 The three anodic peaks were assigned to theoxidation of S02- resulting from the monomerization of S2042-(-0.67V vs. calomel in DMSO-O.lM LiCl), the direct oxidation of dithionite(-0.54V) and the oxidation of the complex, S02S204

2-(-0.22V). Aftercomplete electrolysis, a fourth anodic wave appeared which could have been dueto the oxidation of S022 obtained from the disproportionation reaction:

S2042-* - S02 + S02

2- (7)

A solution of SO2 was electrolysed on a Hg pool at -l.OV (vs. SCE), with theamount of unreduced SO2 monitored by voltammetry. When the electrolysis wasstopped at less than 300s, the original S02 wave could be regenerated by

12Dehn, H., Gutmann, V., Kirch, H. and Schober, G., "Gaspolarographie,"Honatshefte Chemie, Vol. 93, 1962, pp. 1348-1352.

13Garber, R. W. and Wilson, C. E., "Determination of Atmospheric Sulfur Dioxideby Differential Pulse Polarography," Analytical Chemistry, Vol. 44, 1972,pp. 1357-1360.

14Bruno, P., Caselli, M., DellaMonica, M. and DiFano, A., "SimultaneousDetermination of SO2, NO and NO2 in Air by Differential PulsePolarography," Talanta, Vol. 26, 1979, pp. 1011-1014.

15Bruno, P., Caselli, M. and Traini, A., "Study of Sulfur Dioxide ReductionMechanism in Dimethyl Sulfoxide," Journal of Electroanalytical Chemistry,Vol. 113, 1980, pp. 99-111.

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oxidation at -O.15OV, but only after times exceeding 700s. When the

electrolysis was allowed to proceed longer than 300s, it was impossible to

regain all of the original SO2 by reoxidation. The effect of changing the

anion of the supporting electrolyte was also noted, with E11 2 for S02 found

to be -0.711V in O.IM LiNO3 , -0.736V in LiCl and -0.710V in LiC104.

F

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Chapter 4

ACETONITRILE (AN)

I. ELECTROCHEMICAL INVFSTIGATION

A. TETRMALKYhAMONIUM SALTS.

1. Platinum Electrode. Bowden and Dey found no evidence for thereoxidation of SO2- in AN-TBAPF6

6 and assigned the single anodic peakobtained by CV (-0.3V vs. AgCl coated Ag wire) to the oxidation of dithionite,S2042-.

2. Mercury Electrode. The El/ 2 for S02 in AN-TEAP was found byGardner et al to be -l.080V (vs. Ag/AgNO ).8 The equilibrium constant forequation I was calculated to be 18470 M-I, indicating much more extensivecomplex formation than in DMF.

B. LITHIUM SALTS.

1. Platinum Electrode. Geronov, Moshtev and Puresheva reported oneanodic peak appearing during CV in AN-LiBr, but only at SO2 concentrationsgreater than 0.1M.16 This peak (+0.17V vs.SCE) was proposed to be due tooxidation of a product adsorbed onto the surface of the electrode as a result ofthe reactions:

SO2 + e- S02-(ads) (8)

and

S02-(ads) + Li+ "-- LiSO2 (ads) (9)

rather than to oxidation of dissolved species from the bulk of the solution.SO2 was regenerated from the oxidation occurring at +O.17V, as demonstrated byrepeated cycles in the ranges OV to -O.7V and +O.4V to -O.7V. Lithiumdeposition was observed to occur at potentials more negative than -1.5V and

6See footnote 6 on page 10.


16Geronov, Y., Moshtev, R. V. and Puresheva, B., "Electrochemical Reduction ofSulphur Dioxide on Inert Electrodes in Acetonitrile Solutions," Journal of

*i..i Electroanalytical Chemistry, Vol. 108, 1980, pp. 335-346.

17

.. .. . . _ _ _ _ _ _ _ _ __-_ _ _ __.. ._ . ... . . ... . .. .. ... . . . .

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bromide oxidation at potentials more positive than +0.6V. The reductionpotential for SO2 changed linearly with the logarithm of the scan rate, aspredicted by Srinivasan and Gileadi for slow electrochemical adsorptionprocesses.1 7 Calculation of the associated electric charge showed that theelectrode surface was being covered by a monolayer of product.

2. Glassy Carbon Electrode. The electric charge calculated by Geronovet al for glassy carbon was approximately twice that for platinum, presumablydue to differences in the electrode surfaces. 16 The ratio of the anodic peakto the cathodic peak was larger for GC than for Pt, possibly due to differentbonding energies between the adsorbed coating and the substrates.

17Srinivasan, S. and Gileadi, E., "The Potential Sweep Method: A TheoreticalfAnalysis," Electrochimica Acta, Vol. 11, 1966, pp. 321-335.

16 See footnote 16 on page 17.

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Chapter 5

PROPYLENE CARBONATE (PC)

I. ELECTROCHEMICAL INVESTIGATION

A. LITHIUM SALTS.

1. Platinum Electrode. Tikhonova et al observed one anodic peakduring CV in PC-LiC104 (+0.525V vs. TIC1 in PC) 1 8 which was assigned to theoxidation of anion radicals which have not yet reacted according to the scheme:

S02 + Li 4 LiSO 2 (ads) (10)

2LiSO2 4_ Li 2 S204(ads) (11)

The insoluble lithium dithionite coated the electrode surface, blocking furtherreaction. On repeated cycles, the peaks were greatly diminished.

Shembel' et al found two anodic peaks during CV of a similar PC-LiCl04solution containing about 20 times as much S02 (2.2M vs. 0.12M).

19 Thesepeaks were attributed to the oxidation of S02 (3.42V vs. Li) andS2 042-(4.08V). Blockage of the electrode surface by product buildup wasnoted, but with the additional observation that at low sweep rates (10wlVs),during repeated scans, the cathodic peak increased from cycle to cycle. Whenthe reduction product was completely reoxidized (i.e., both anodic peaks weretraversed in the scan), the system could be operated in reverse. The anodicpeak became more positive as the switching potential became more negative andthe potential difference between the cathodic and anodic peaks increased(indicating increased overpotential). An SO2 - solvent complex was alsopostulated to explain lower than expected cathodic currents.

2. Glassy Carbon Electrode. Shembel' et al observed only one anodicpeak on glassy carbon, as opposed to two on platinum.19 The anodic peak grew,the more negative the switching potential. Lithium was deposited at negativepotentials, as indicated by an anodic dissolution peak.

* t 8 Tikhonova, L. S., Kurchashova, N. A., Raikhel'son, L. B., Nikol'skii, V. A.and Sokolov, L. A., "Investigation of the Electrochemical Reduction of SulfurDioxide in Propylene Carbonate," Soviet Electrochemistry, Vol. 14, 1978,pp. 1563-1566.

19Shembel', E. M., Litv6nova, V. I., Maksyuta, I. M., Moskovskii, V. Z.,Sokolov, L. A.. and Ksenzhek, 0. S., "Cathodic Reduction of Sulfur Dioxide inNonaqueous Electrolytes: Lithium Perchlorate Solution in Propylene Carbonate,"Soviet Electrochemistry, Vol. 15, 1979, pp. 1671-1677.

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Chapter 6

OTHER SOLVENTS

1. FORMWEIDE

A. CHEMICAL REDUCTION.

Rinker and Lynn observed that when sodium amalgam was added to asolution Of SO2 in formamide, a red intermediate (RRI) quickly developed, justas quickly becoming colorless.1 When DUIF was added to RI, blue radical ion(ElI) appeared. Apparently both had the formula, S02S02-, with colordifferences due to changes in solvation. The colorless ion (CI) appeared to beS02S 2

2-. A solution of CI could be changed to RRI by adding SO2.Addition of SO2 to sodium dithionite dissolved in formamide resulted in RI.When RRI was further reduced by Na(Rg), CI was formed. This series ofreversible reactions demonstrated thrs chemical relationships among the species.

It. UEXLANETEYLPHOSPHORAMIDE (EWPA)

A. ELECTROCHEMI CAL INVESTIGATION.

1. Tetraethylamnium Salts.

a. MERCURY ELECTRODE. The El/ 2 for SO2 in IDIPA-TEAP wasfound by Gardner et al to be -1.167V (vs.- Ag/AgNO3), with the equilibriumconstant for equation 1 equal to 1456K 1. This indicated considerably lessstability of the complex in HMPA.8

111. METHYLENE CHLORIDE

A. ELECTROCHEMICAL INVESTIGATION.

1. Tetraalkylamonium Salts.

a. PLATINUM ELECTRODE. Bowden and Dey found in MeC12-TBAPF 6that CV Of SO2 resulted in two anodic peaks, assigned to the reductionproducts, S02 -(-0.4V vs. ASCl coated Ag wire) and S20(4O6V.



6Seea footnote 6 on page 10.

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Chapter 7

DISCUSSION

The reduction of SO2 in nonaqueous solvents results in two majorproducts: (S02)xSO2 and S2042. Some disagreement exists, asshown in Tables 4 and 5, but the predominant species depends on numerousfactors, including the nature of the solvent, the supporting electrolyte, 802concentration, temperature, electrode surface and the presence of impuritiessuch as water.

1. SOLVENT. All of the solvents studied were polar and aprotic.Differences in the observed mechanism for S02 reduction seem to be directlyrelated to individual capacities for solvation. Baranski and Fawcett found thatwhen the solvent was changed, keeping a fixed supporting electrolyte, the log ofthe standard rate constant for the reduction of Na+ ions at a Hg electrodevaried linearly with free energies of salvation.20 Gardner et al demonstratedthat the products of reduction depend on the degree of ionic association betweenthe free radical and supporting electrolyte cations.8 Polar solvents of highdielectric constant discourage ion pairing, because of stronger solute-solventinteractions.2 1 The solvents in these studies are all highly polar, but thedielectric constants vary considerably, as shown by Covington's2 1 values inTable 6.

In poor solvators, the free radical ion is more available for complexationby unreduced S02 to form the species, (S02)xSO2-, demonstrated byGardner et al with their values for the complex formation constant of 1456M1in HMPA, an excellent solvator, 4363N-1 in DMF and 18470M- 1 in AN8,presented with other reported values in Table 7. Also observed by Gardner et alwas a shift in half wave potential, from -1.257 in DHF to -1.167 in UNPA to-1.080 in AN. This transition does not follow the trend of complex formationattributed to solvation.

2 0Baranski, A. and Fawcett, W. R., '"edium Effects in the Electroreduction ofAlkali Metal Cations," Journal of Electroanalytical Chemistry, Vol. 94, 1978,pp. 237-240.

r 8fee footnote 8 on page 11.

2 1Covington, A. K. and Dickinson, T., Ed., Physical Chemistry of OrzanicSolvent Systems (London: Plenum Press, 1973), pp. 332-393.

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Color differences have been observed in solutions of reduced SO2 . Someof this effect can be attributed to differences in solvation mechanisms. Rinkerand Lynn found that amides with one or more protons on the amide nitrogen gave ared intermediate, where others gave a blue species. 1 Dry DMF is aprotic andnonbasic, but other such solvents, e.g. dioxane, cannot stabilize BRI. Theunexpected stability of BRI in D1F was explained by the absence of H+ andOH-, the solvent "cage" effects resulting from complexes between solvent andradical, and resonance structures.

An equilibrium constant was calculated for the dimerization of S02- inDMS015 to be 105, as opposed to 24 in DMF8 .

It would seem that generally AN favors complex formation, whereas DNSOfavors dimerization.

2. SUPPORTING ELECTROLYTE. Studies by Gardner et al showed that theformation constant for the complex varied when the supporting electrolyte waschanged.8 In DMF, decreasing the cation size from K+ to Na+ to Li+

resulted in increasing dimer formation. It was concluded that the much largertetraalkylammonium salts favor complexation, while lithium salts favordimerization.

Interesting comparisons can be made, however, when examining Table 4. BothMartin3 and Gardner et al8 reported two anodic peaks in DHF-O.lM TEAPseparated by 0.5V, assigned to the oxidation of S02- and SO2SO2-.Magno et a l and Bowden and Dey 6 found two anodic peaks in DMF-TBA+

separated by 0.75V which were attributed to the oxidation of S02- and82042-. Magno et al observed the complex spectrophotometrically, butreasoned that it was too unstable to be detected voltammetrically. It wouldseem that different species are responsible for the peaks observed with TEA+

and TBA+, as evidenced by the different voltage separations.

Bruno et al found a shift in the SO2 reduction potential when the anionwas changed in a series of Li+ salts.15 If complexation of SO2 by theanion had been responsible, a cathodic shift would have occurred on increasingthe electrolyte concentration. It was proposed that changing the anion modifiedthe rate of dimerization. The observed half-wave potential was plotted

18ee footnote 1 on page 9.



r3See footnote 3 on page 9.

4 See footnote 4 on page 10.I 680e footnote 6 on page 10.

24

L.

NSWC TR 80-533

vs. 1o, the potential at the outer Helmholtz plane, for differentconcentrations of LiClO4 and LiNO3 , resulting in reasonably linearcorrelations.

3. CONCENTRATION OF SO. Geronov et al reported that anodic peaksdid not appear in cyclic voltammograms until the 802 concentration was atleast 0.1M ,6 but many other studies observed peaks at concentrations in therange of 10-4 to IO-2M.

Gardner et al found a trend toward more anodic potentials for the half-wavepotential of S02 in DMF-TKAP8 , as shown in Table 8. On the other hand,Bruno et al 15 and Martin and Sawyer3 reported that the peak potentialshifted cathodic as (SO2] increased. In a detailed study, Bonnaterre andCauquis observed that the peak potential became more anodic as the concentrationof SO2 increased9 , until the effect reversed at a concentration between10-3 and 10-2M (depending on the sweep rate) and the potential moved towardmore cathodic potentials with continued increase in S02.

As the amount of S02 increased, the ratio of the more negative anodicpeak to the more positive anodic peak decreased, as observed by Magno et a1

4

and Martin and Sawyer3 in DMF and Bonnaterre and Cauquis in DMS0 9 . Thisfact could be taken as evidence for the assignment of the second anodic peak tothe oxidation of the complex, (S02)xSO2-.

4. TEMPERATURE. Magno et al found a strong dependence of thepotentials of the anodic peaks on temperature.4 At 0°C the peaks were at-0.5 and +0.2V (vs. SCE), but at 450 C they shifted cathodically to -0.8 and-0.25V. The cathodic peak potential was practically unchanged. The ratio ofthe anodic peaks decreased as the temperature decreased.

* Gardner et al reported that the shift observed in the reduction potentialof SO2 depended on the nature of the electrolyte

8 , as shown in Table 9. InDMF-TEAP, the potential became more positive at lowered temperature.

According to Bonnaterre and Cauquis, increased temperature (550 C) reducedthe effect of the diner costing on further reactions at the electrodesurface9 , making the results more reproducible and less affected by otherparameters.

5. ELECTRODE SURFACE. Martin and Sawyer reported that cyclicvoltammograms of S02 in DMF-TAP obtained on Pt and Au electrodes are similar,

168ee footnote 16 on page 17.




9S* footnote 9 on page 13.


25

i*

NSWC TR 80-533

but the ratio of the anodic peak to the cathodic peak is larger at Pt3. It

was also pointed out that the Au electrode is less susceptible to hysteresis.Gardner et al ran cyclic voltametry in the same DMF-TEAP solution at the Hgelectrode (HlNDE) and observed analogous results 8, as can be seen in Table 4.Geronov et al studied the anodic electric charge for Pt and GC electrodes,finding it to be much lower for Pt, presumably due to the relative smoothness ofthe surface. 16 The GC surface was more rapidly passivated, as demonstrated byrepeated scans. Only one anodic peak was observed on GC by Shembel' et al,compared to two on Pt. 19

6. EFFECT OF WATER. Tikhonova et al added water to the PC-LiCl04electrolyte, increasing the reduction peak and producing a second reduction peakat more negative potentials.18 Geronov et al concluded that the effect ofwater was not evident immediately, but appeared gradually over a period of hoursas the cathodic peak increased. 1

Dehn et al had found that small amounts of water increased the half-wavepotential of SO2 in DMSO, whereas the presence of greater than 5 vol Z splitthe wave into two adjacent waves.

12



-' 16 See footnote 16 on page 17.



* 16See footnote 16 on page 17.

12 See footnote 12 on page 14..

26

J _

NSWC TR 80-533

Chapter 8

SUMMARY AND PLANS

As the previous discussion shoved, there has been no general agreement onthe mechanism of the reduction of S02 in nonaqueous solvents. The choice ofsolvent affects the outcome of the reduction, with DMF and AN favoring thecomplex and DMSO the dimer. Electrolytes with large cations favor the complex,but Li+ ions produce the dimer. This generalization is questionable in viewof the discrepancies reported between studies in TEA* and TBAt solutions(see Chapter 7, section 2). Electrode surface is also important because of thepossibilities of catal)r' effects and/or side reactions.

Cyclic Voltm , -,erates information about reaction reversibility,intermediate sp.v *r,; reduction mechanisms. The effect of changes inelectrolyte and , surface is observed in the position and shape ofvoltammetric peak_,

Coulomet:.t detesiaiines the number of electrons involved in the electrontransfer, possible revealing mechanistic information in the shape of thecurrent-time cur',e.

Controlled potential electrolysis produces reactive intermediates which can* . be detected by other voltammetric techniques or by UV-vis spectroscopy, which

monitors the appearance and disappearance of charge-transfer type complexes duringthe course of the reduction. Those species which have an unpaired electron arealso detectable with ESR spectroscopy.

Rotating ring-disc electrodes provide information about the half-life ofsoluble intermediates, possibly detecting adsorption at the electrode surface.

More detailed experimental plans will be presented in Part II of thisreport, along with preliminary results from cyclic voltanmetry and other studies.

27/28

II

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TABLE 1 PUBLISHED STUDIES OF S02 REDUCTION*

PUBLICATION WORKING

AUTHOR DATE SOLVENT ELECTROLYTE ELECTRODE

DEHN 1962 DMSO TEAP Hg

DINSE 1968 DMF TPAP, TBAI Pt

RINKER** 1968 DMF, DMSO, F

BONNATERRE 1971 DMSO TEAP Pt

GARBER 1972 DMSO LiC1 Hg

MARTIN 1972 DMF TEAP Pt

MARTIN 1973 DMF TRAP Pt

MAGNO 1974 DMF TBAP Pt

TIKHONOVA 1978 PC LiC10 4 Pt

BRUNO 1979 DMSO LiCI Hg

KASTENING 1979 DMF TEABr Pt

SHEMBEL' 1979 PC LiC104 Pt, GC

GARDNER 1979 DMF, AN, HMPA TRAP, LiC104 Hg

GERONOV 1980 AN LiBr Pt, GC

BOWDEN 1980 DMF, AN, MeC12 TBAPF6 Pt

BRUNO 1980 DMSO LiC1, LiNO3 , LiCI0 4 HS

FOUCHARD 1980 DMF TEAP Pt

* ABBREVIATIONS EXPLAINED IN TEXT

• CHEMICAL REDUCTION WITH Na(Hg)

29

29

L..i ', ..

NSWC TR 80-533

0 m ~ c4 Go el-- I * * I - • U% 40 • 40

+ 0 0°

0 CD0 0 -00I +I

Lft

'0n

04 - U.cu 11.

-~ 0 0 * ,0.4- 0 0 C 0 .-

00- 4 I < II I

!.t 00

0G

041

404 41 00."- '4--4

a 0100 .. _. " 'U3 2 w4~ z. I- .4 0."

x04 E. o ' 3po &^ - %0

I04 N 4

x c4

Go CD o

c4 c4. C- e0 0 C

9d0

0 0 m 0 a X 0 0 0. 0 Q 0u

00-C(N ( N0 d (

a 0

Al r

w3

N$WC TR 80-533

-3x

40 I * s 4 i S I 00.

-CU C4 46 4 1

o 0 0 67 - L4 -.

I3 O 0 0 0

cc

4, -4 A.

x.) X X X X 0. -

d~pQ 0'a'.

o ;31

NSWC TR 80-533

TABLE 4 ASSIGNMENT OF ANODIC PEAKS IN CYCLIC VOLTAMMETRY INTETRAALKYLAMMONIUM SALT SOLUTIONS

ANODICAUTHOR SOLVENT ELECTROLYTE PEAK POTENTIAL SPECIES ELECTRODE

MARTIN DMF TEAP -0.74 SO2- Pt-0.24 S02 S02 - Pt

MAGNO DMF TBAP -0.65 so 2 - Pt+0.1 S2042 P

BOWDEN DMF TBAPF 6 -0.13 SO2 Pt+0.63 S2 04

2 - Pt* GARDNER DMF TEAP -1.17 SO2 - Hg

-0.65 S02S02- g

BONNATERRE DMSO TEAP S2042 P

BOWDEN AN TBAPF6 -0.3 S2042- Pt

* BOWDEN MeC1 2 TBAPF 6 -0.4 S02 - Pt+0.06 S2 04

2 - Pt

3

32

___________________ _________________________

NSWC TR 80-533

TABLE 5 ASSIGNMENT OF ANODIC PEAKS IN CYCLIC VOLTAMETRY INALKALI SALT SOLUTIONS

ANODICAUTHOR SOLVENT ELECTROLYTE PEAK POTENTIAL SPECIES ELECTRODE

GARDNER DMF LiC1O4 -0.62 S2 042 - Hg

GARDNER DMF KC10 4 ?S0 2 S0 2 - Hg

BRUNO DMSO LiNO3 -0.67 S0 2 - Hg

-0.54 S2042 - Hg

-0.22 S02S2 042- Hg

GERONOV AN Li~r +0.17 LiS02 Pt

GERONOV AN LiBr ?LiS0 2 GC

TIKHONOVA PC LiCl04 +0.525 S02-P

SHEMBEL' PC LiC1O4 +3.42 S02- Pt

+4.08 S2 042 - Pt

33

NSWC TR 80-533

TABLE 6 DIELECTRIC CONSTANTS E)AND DIPOLE

MOM4ENTS ()OF NONAQUEOUS SOLVENTS

SOLVENT If a

F 109.5 3.73

PC 64.4 4.98

DMSO 46.7 3.96

DMF 36.7 3.86

AN 36.0 3.92

UNPA 29.6 5.39

34

NSWC TR 80-533

TABLE 7 EQUILIBRIUM CONSTANT FOR COMPLEX FORMATION

SOLVENT ELECTROLYTE AUTHOR K

DMF TPAP DINSE 1.3x104 M72

DMF TEABr KASTENING 230 M71

DMF TPAP DINSE RECALCULATED 65-200 M-1BY KASTENING

DMF TEAP FOUCHARD 600 M- I

DMF TEAP GARDNER 4360 M-1

AN TEAP GARDNER 18470 M'-1

HMPA TEAP GARDNER 1456 WI

3

1 35, I

NSWC TR 80-533

TABLE 8 VARIATION OF El/2 WITH [SO2 ] IN DIIF-TEAP

IS021E 11 2

5X10-4 -1.279

10-3 -1.257

36

NSWC TR 80-533

TAKEI 9 VARIATION OF El2WITHi TEMPERATURE

A. DMF-TEAP

-10 -1.224

O -1.234

20 -1.254

B. DMF-LiC 104

OC E/

-10 -1.260

O -1.250

20 -1.235

37/38

i

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BIBLIOGRAPHY

Baranski, A. and Favcett, W. R., "Medium Effects in the Electroreduction ofAlkali Metal Cations," Journal of Electroanalycical Chemistry, Vol. 94, 1978.

Bonnaterre, R. and Cauquis, G., "Dimerization Consecutive a un TransfertMonoelectronique I. Equation des Courbes Intensite - Potentiel,"Electroanalytical Chemistriy and Interfacial Electrochemistry, Vol. 32, 1971.

Bonnaterre, R. and Cauquis, G., "Dimerization Consecutive a un TransfertMonoelectronique II. Reduction Electrochemique du Bioxyde de Soufre dang leDimethyl Sulfoxyde," Electroanalytical Chemistry and InterfacialElectrochemistry, Vol. 32, 1971.

Bowden, W. L. and Dey, A. N., "Primary Li/SOCl2 Cells XI. SOC12 ReductionMechanism in a Supporting Electrolyte," Journal of the Electrochemical Society,Vol. 127, 1980.

Bruno, P., Caselli, M., DellaMonica, Mi. and DiFano, A., "SimltaneousDetermination of SO2 , NO and NO2 in Air by Differential Pulse Polarography,"Talenta, Vol. 26, 1979.

Bruno, P., Caselli, H. and Traini, A., "Study of Sulfur Dioxide ReductionMechanism in Dimethyl Sulfoxide," Journal of Electroanalytical Chemistry,Vol. 113, 1980.

Covington, A. K. and Dickinson, T., Ed., Physical Chemistry of Organic SolventSystems (London: Plenum Press, 1973).

Dehn, H., Gutmann, V., Kirch, H. and Schober, G., "Gaspolarographie,"l

Monatshdfte Chemie, Vol. 93, 1962.

Dinse, K. P. and Mobius, K., "EPR - Untersuchungen an elektrolytisch erzeugtemSO2-," Zeitschrift fur Naturforsch A, Vol. 23, 1968.

j Fouchard, D. T., Gardner, C. L., Adams, W. A. and Laman, F. C., "Raman and ESRSpectroscopic Studies of the Electroreduction of Sulphur Dioxide," inProceedings of the Symnosium on Lithium Batteries, Electrochemical SocietyMeeting, Hollywood (FL), October, 1980.

Garber, R. W. and Wilson, C. E., "Determination of Atmospheric Sulfur Dioxide byDifferential Pulse PolaroSraphy," Analytical Chemistry, Vol. 44, 1972.

39

NSWC TI 80-533

BIBLIOGRAPHY (Cont.)

Gardner, C. L., Fouchard, D. T., Leman, F. C. and Fawcett, W. R., "The Kineticsand Mechanism of the Reduction of Sulphur Dioxide in Nonaqueous Media," inProceedings of the Symposium on Lithium Batteries, Electrochemical SocietyMeeting, Los Angeles, October, 1979.

Geronov, Y., Moshtev, R. V. and Puresheva, B., "Electrochemical Reduction ofSulphur Dioxide on Inert Electrodes in Acetonitrile Solutions," Journal ofElectroanalytical Chemistry, Vol. 108, 1980.

tastening, B. and Gostisa-Mihelcic, B., "An ESR Investigation of the ComplexFormation Between SO2 and Electrogenerated S02-," Journal ofElectroanalytical Chemistry, Vol. 100, 1979.

Laviron, E., "Mechanism of the Polarographic Reduction of Aromatic CarbonylDerivatives: Influence of Dimerization of the Free Radical Formed in the FirstReduction Stage," Collection of Czechoslovak Chemical Communications, Vol. 30,1965.

Magno, F., Mazzochin, G. A. and Bontempelli, G., "Voltammetric Behavior ofSulphur Dioxide at a Platinum Electrode in Dimethylformamide," ElectroanalyticalChemistry and Interfacial Electrochemistry, Vol. 57, 1974.

Martin, R. P., "The Electrochemical Reduction of Sulfur Dioxide and of ElementalSulfur in Nonaqueous Solvents," Ph.D. Dissertation, University of California,

OF eRiverside, 1973.

Martin, R. P. and Sawyer, D. T., "Electrochemical Reduction of Sulfur Dioxide inDimethylformamide," Inorganic Chemistry, Vol. 11, 1972.

Matsuda, H. and Ayabe, Y., "Zur Theorie der Randles-SeveikschenKathodenstrahl-Polarographie," Zeitschrift fur Elektrochemie, Vol. 59, 1955.

,* Olmstead, M. L., Hamilton, R. G. and Nicholson, R. S., "Theory of CyclicVoltammetry for a Dimerization Reaction Initiated Electrochemically," AnalyticalChemistry, Vol. 41, 1969.

Rinker, R. G. and Lynn, S., "Formation of Sodium Dithionite from Sodium Amalgamand Sulfur Dioxide in Nonaqueous Media," Industrial and Engineering Chemistry:Product Research and Development, Vol. 8, 1969.

Shembel', E. N., Litvinova, V. I., Maksyuta, I. M., Moskovskii, V. Z.,Sokolov, L. A. and KIenshek, 0. S., "Cathodic Reduction of Sulfur Dioxide inNonaqueous Electrolytes: Lithium Perchlorate Solution in Propylene Carbonate,"Soviet Electrochemistry, Vol. 15, 1979.

Srinivasan, S. and Gileadi, K., "The Potential Sweep Method: A TheoreticalAnalysis," Electrochimica Acta, Vol. 11, 1966.

Tikhonova, L. S., Kurchashova, N. A., Raik'helson, L% B., Nikol'skii, V. A. andSokolov, L. A., "Investigation of the Electrochemical Reduction of Sulfur

I Dioxide in Propylene Carbonate," Soviet Electrochemistry, Vol. 14, 1978.

40

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, 1',

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BIBLIOGRAPHY (Cont.)

Wopechall, R. H. and Shain, I., "Effects of Adsorption of Electroactive Speciesin Stationary Electrode Polarography," Analytical Chemistry, Vol. 39, 1967.

Wopschall, R. H. and Shain, I., "Adsorption Effects in Stationary ElectrodePolarography with a Chemical Reaction Following Charge Transfer," AnalyticalChemistry, Vol. 39, 1967.

. /

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p-~

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DISTRIDUTION (Cont.)

Copies Copies

RAY-O-VAC University of MarylandAttn: R. Foster Udell I Attn: J. M. Bellama101 East Washington Avenue Dept. of ChemistryMadison, WI 53703 College Park, MD 20742

CAPT. A. S. Alanis Queens CollegeBMO/ENBE Attn: G. D. DuncanNorton AFB Dept of ChemistryCA 92409 1 Charlotte, NC 28274

Norton AFB Spring Arbor CollegeCode AFISC/SES Chemistry DepartmentCA 92409 1 Attn; Dr, David A, Johnson

Spring Arbor, MI 49283

GTE SylvaniaAttn: Dr. Robert McDonald189B StreetNeedham Heights, MA 02194

Martin Marietta AerospaceAttn: John W. LearP. 0. Box 179Denver, CO 80201

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