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Electrochemistry of hydrogen halides in dimethylsulfoxide
ELECTROANALYTICAL CHEMISTRY AND INTERFACIAL ELECTROCHEMISTRY Elsevier Sequoia S.A., Lausanne - Printed in The Netherlands
387
ELECTROCHEMISTRY OF HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
MANFRED MICHLMAYR AND DONALD T. SAWYER
Department of Chemistry, University of California, Riverside, California 92502 (U.S.A.) (Received October 10th, 1968; in revised form May 13th, 1969)
Although the polarographic and voltammetric reduction of the hydrogen halides (HC1, HBr and HI) has been studied extensively, almost no work has been reported for their electrochemical oxidation. In aqueous systems they are dissociated and only the oxidation of the halide ion or the formation of an insoluble mercurous salt by the mercury electrode is observed. However, hydrogen halides are much less dissociated in nonaqueous media than in water 1. Investigations of their electrical conductivity in anhydrous solvents 2- 4 have shown that the systems involve a number of complex equilibria such that the ionization constants obtained have no more than qualitative significance 1. The electrochemical oxidation of halide ions has been studied in several nonaqueous systems 5-10, but little is known about their electrochemical oxidation in dimethylsulfoxide (DMSO) 4'5,11,12. The reduction of HC1 in DMSO has been discussed 13a4. A general mechanism has been proposed for the oxidation of halides in nonaqueous media: 3 X- --0 X3 + 2eXg ~ } X 2 + e -
(1) (2)
based on the data for iodide ion. Two waves corresponding to the two oxidation reactions are observed at +0.5 and +0.7 V vs. SCE for iodide in DMSO 5. The reactivity of C12, Br2, and I2 with a number of organic solvents 8 is relevant to the present study because the reaction forms He1, HBr, and HI, respectively This paper summarizes the results of a detailed investigation of the electrochemistry of HCl, HBr, and HI in DMSO. From the studies, the kinetic parameters for the electrode reactions have been evaluated and the oxidation products have been identified; this has permitted plausible reaction mechanisms to be proposed. EXPERIMENTAL
Apparatus
The electrochemical measurements were made with a versatile electronic instrument constructed from Philbrick operational amplifiers following the design of DeFord is. By appropriate interconnection of the operational modules, the instrument can perform chronopotentiometry, controlled-potential coulometry and cyclic voltammetry. Potential-time curves were recorded with a Sargent model SR strip-chart recorder using a speed of 12 in./min, and cyclic voltammograms with a J. Electroanal. Chem., 23 (1969) 387-397
388
M. MICHLMAYR, D. T. SAWYER
Moseley X-Y recorder model 7030A. A function-generator (Hewlett-Packard model 202 A) was used for the cyclic voltammetry at high scan rates, and a Tektronix model 564 oscilloscope was employed to record these current-potential curves as well as chronopotentiograms with transition times of less than 1 s. The electrochemical cell and the three-electrode assembly used for chronopotentiometric and voltammetric measurements have been described previously 16. The electrochemical area of the platinum working electrode was determined using Fe(CN) 3- in 0.5 F aqueous KC1; its diffusion coefficient 17 was assumed to be 7.67 x 10-6 c m 2 s - 1 . Potentials are reported vs. SCE unless otherwise noted. The pH was monitored with a Leeds and Northrup miniature pH glass electrode assembly and a Corning model 12 pH meter. Reagents
A 0.1 F solution of tetraethylammonium perchlorate (Eastman Organic Chemical Co.) in DMSO (J. T. Baker Co.) was used as supporting electrolyte. Prior to use, the salt was recrystallized from water, dried at 70°C in vacuo and stored in a vacuum desiccator. For a number of experiments 0.1 F lithium perchlorate (G. F. Smith Chemical Co.) also was used as supporting electrolyte. Anhydrous hydrogen chloride, hydrogen bromide and hydrogen iodide were obtained from Matheson Co. ; their solutions were prepared by passing the gas through concentrated H 2 S O 4 prior to bubbling it slowly (with cooling) into DMSO. The concentrations were determined by potentiometric titrations with 0.01 F AgNOa and 0.01 F KOH. Solutions of H2SO4 and HC104 in DMSO were prepared as described previously 13. Acetonitrile (Mallinckrodt, nanograde) was purified by distillation from calcium hydride, and N,N-dimethylformamide (DMF) was obtained from Matheson, Coleman and Bell (spectroquality). All other materials were reagent-grade and were used without further purification. Either argon or prepurified nitrogen was used as the inert gas; the solutions were deaerated for 20 rain before measurements. The gas stream was presaturated by passing through a 0.1 F solution of EtgNC104 or LiC104 in DMSO, DMF, or CHaCN before entering the electrolysis cell. RESULTS
Anhydrous hydrogen halides give well-defined cathodic and anodic chronopotentiograms in D MSO at a platinum electrode (Fig. 1); the quarter-wave potentials for the reduction waves are virtually identical at - 0.50 V. The E~ values for the oxidation waves are +i.00 V, +0.88 V, and +0.60 V vs. SCE for HCI, HBr, and HI, respectively. 1. Reduction
The cathodic waves at -0.50 V correspond to the reduction of hydrogen ion, which is a diffusion-controlled process. This is established by the constant values of iT½/AC°; combination of these data with the Sand equation 18 permits evaluation of the diffusion coefficients, D, for the three acids. The respective values for HC1, HBr, and HI are 4.3 x 10 -6, 4.0 X 10 -6, and 3.7 x 10 - 6 c m 2 s -1 which are in good agreement with previously determined values 13'19 for HC1, HC104, and H2SO4. The diffusion coefficients of H2SO4 and HC104 also have been determined; both have J. Electroanal. Chem., 23 (1969) 387-397
389
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
1.00
i = 500 ~ A
r~
0.00 h
- 1.00
- 2.00
L
i
I
3 4 2 Time (see) Fig. 1. Chronopotentiograms for 3.00 × 10 -3 F HCI in D M S O containing 0.10 F Et4NC10~ at a platinum electrode (area, 0.277 cm2). 0
1
values in D M S O of 4.35 x 10 - 6 c m 2 s - 1 with H 2 S O 4 acting as a monobasic acid. The potential-time relationship for an irreversible system at 25°C is given b y I 8 : E - 0.059 log nFC°k~'hA 0.059 z~-t ~ ~na i + --~na l o g - - z ~
(3)
where E is the potential of the working electrode vs. NHE, ct the cathodic transfer coefficient, na the number of electrons in the rate-determining step of the reduction process, F the faraday, k°,h the heterogeneous rate constant for the reduction reaction, i the current, A the electrode area, C O the concentration of the electroactive species in the bulk of the solution in mol cm- 3, T the transition time and t the time after the electrolysis is started. This equation can be used to evaluate ~na and log k~,h by plotting E vs. log (r ~ - t~/z ~) for the chronopotentiometric waves. The respective values for HCI, HBr, and HI at a platinum electrode are 0.67, -5.14; 0.65, -4.96; and 0.64, -4.63. A more meaningful parameter for comparison of electrochemical kinetic data is the simple heterogeneous rate constant, ks.h. This is the rate constant at the equilibrium potential, E, for the rate-determining couple and can be evaluated from the relation 18 ks,h ---- kfOhexp (-- ~n~FE'/R T)
(4)
Because the formal potential for the reaction H ÷ + e- ~ ½ H 2 is not known in D M S O solutions, evaluation of ks,h requires that the value of E' be estimated. If the thermodynamic value of E ° for the reduction reaction in H 2 0 is used, ks, h is equal to kf0,hin view of the definition of NHE. J. Electroanal. Chem., 23 (1969) 387-397
390
M. MICHLMAYR, D. T. SAWYER
I
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391
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
2. Oxidation
In addition to the major anodic waves for the hydrogen halides in DMSO, HBr and HI exhibit small prewaves at +0.60 and +0.44 V vs. SCE, respectively. Their height is approximately ~ of the total transition time, i.e., much too small for consecutive two- and one-electron electrode reactions. The sum of the waves has been used for evaluations of the effect of current density and concentration on the transition time. Table 1 summarizes such data for the three halogen acids and indicates that the values of iz ~ for the oxidation wave of HC1 and HBr decrease with increasing current density; this normally is indicative of either a preceding chemical reaction or of a catalytic process 18. Consideration of the data in Fig. 1 and Table 1 indicates that the ir~ values for the anodic waves for HC1 and HBr are almost an order of magnitude larger than for the cathodic waves. This implies a catalytic process with regeneration of the hydrogen halide. Such a conclusion is supported by the data of Table 2 which establish TABLE 2 CHRONOPOTENTIOMETRIC DATA FOR THE OXIDATION OF HCI AND HBr AT HIGH CURRENT DENSITIES AT A PLATINUM ELECTRODE (area, 0.227 cm 2) IN DMSO HCl
HBr
C Ox 106/mol cm-3
i/#A
iz½/#A s½
C Ox 106/mol cm-3
i/#A
iz½/pA s~'
0.124 0.124 0.124 0.124 0.240 0.240 0.240 0.240 0.354 0.354 0.354 0.354 0.467 0.467 0.467 0.467 0.580 0.580 0.580 0.580
150 200 300 500 300 400 600 900 400 600 800 900 400 600 800 1000 500 700 800 1000
4.39 4.45 4.51 4.75 8.78 8.93 8.33 8.62 13.0 12.8 12.5 13.3 17.2 t6.6 17.6 16.7 20.6 20.8 21.0 19.8
0.186 0.186 0.186 0.186 0.370 0.370 0.370 0.370 0.554 0.554 0.554 0.554 0.738 0.738 0.738 0.738 0.922 0.922 0.922 0.922
200 300 400 500 300 400 500 550 400 500 600 700 450 600 700 800 500 600 700 900
7.20 7.75 8.30 7.75 15.1 15.4 14.9 14.3 24.2 21.2 21.7 22.0 30.4 28.9 30.0 30.6 39.4 39.7 40.1 39.3
iz ½
Average: A ~ 6
= 160 pA s½em mo1-1
iz ½ A C o - 181 #A s~cm mo1-1
that at high current densities the values of iz~/AC ° become independent of current and concentration and are reduced to values comparable to those for the reduction waves for HC1 and HBr. If the average values in Table 2 are assumed to represent a diffusion-controlled one-electron process that is uncomplicated by the catalytic J. Electroanal. Chem., 23 (1969) 387-397
392
M. MICHLMAYR, D. T. SAWYER
process, they can be used with the Sand equation is to evaluate the anodic diffusion coefficients for HC1 and HBr. The data for HI in Table 1 imply the absence of catalytic complications and can be used to evaluate its diffusion coefficient. The respective values of D for HC1, HBr, and HI are 3.5 x 10 -6, 4.5 x 10 -6, and 4.8 × 10 -6 cm2s -1. By assuming that the high current-density data for HC1 and HBr represent the diffusion-controlled process, the average values for i z ~ / A C ° can be used to calculate the hypothetical diffusional transition times, Zd, for the conditions of Table 1 if there were not catalytic complications. The observed transition times are representative of the catalytic process and can be combined with the calculated Zd values to give the quantity (Z/Zd)+ which is indicative of the extent of the catalytic reaction (Table 1). Delahay 18 has shown that when the value of this ratio is greater than 2.4, a firstorder catalytic process obeys the expression (5)
(Z/Zd) ~ = 2(kcCzz/rc) ~
where kc is the catalytic rate constant and Cz the concentration of the substance that is catalytically oxidized. A plot of (Z/Zd) ~ VS. Z~ should give a linear slope that is proportional to (kc Cz)~. Such a treatment of the data in Table 1 yields curved lines whose slopes decrease with increasing values of z ~ and indicates a more complicated process than a first-order reaction. The values of (k~ C~)½ calculated from the slopes range from 1.1 to 0.5 for HC1 and from 0.3 to 0.1 for HBr. When the anodic current for the hydrogen halides is reversed before reaching the transition time, a reduction wave appears at -0.50 V (Fig. 1) which has the general features of the cathodic hydrogen ion reduction wave (~na=0.67). The ratio of its transition time to the electrolysis time before reversing the current is 1 : 3 for HCI, and between 1:2.5 and 1:3 for HBr. Current reversal prior to the transition time for the anodic HI wave yields a cathodic wave with an E0.22 value of +0.36 V. The ratio of zr/zf is always smaller than the theoretical 1 : 3 and averages about 1 : 8. When the reduction wave at +0.36 is reversed, the anodic wave at + 0.44 (small prewave) is found to be larger than before. For this case, the ratio of the transition times is 1 : 3 (after subtraction of the previous wave). The cathodic wave at - 0.5 V does not show any net increase in ir ~ on a reverse chronopotentiogram. A relation similar to eqn. (3) has been developed for irreversible anodic chronopotentiograms 1s E -
0.059 n F C ° k°,h A (1 - ct)na log i
0.059 z~ - t ~ (1 - c~)n~log - - z ~
(6)
where (1 - a ) is the anodic transfer coefficient and kbOhthe heterogeneous rate constant for the oxidation reaction. By extrapolating the potential to the beginning of the chronopotentiogram, plots of Et=o vs. log i and Et=o vs. log C o should give straight lines with slopes that allow evaluation of ( 1 - a)n a. Figure 2 and Table 3 indicate that for low current densities such plots give meaningful results; at higher current densities (with a transition time of 30 s, or less) irregular (1 - a)na values are obtained for HCI and HBr. By using the average values of (1 - ~ ) n a (0.52, 0.61 and 0.72 for HC1, HBr, and HI, respectively) and assuming that n is unity, the quantity log k°h has been evaluated from eqn. (6); the mean values are - 13.1, - 13.3 and - 12.0 for the three hydrogen halides. J. Electroanal. Chem., 23 (1969) 387-397
393
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
Again, ks,h can be calculated from the relation: ks,h = k°,h exp [(1 - a ) n a F E ' / R T ]
(7)
by using the thermodynamic values, E °, for the half reactions X- ~ ½X 2 + e- in H 2 0 (+ 1.36, + 1.07, and +0.54 V vs. NHE2°). The respective values for log ks,h are -2.0, -2.9, and -5.5 for HC1, HBr and HI. If the platinum electrode is polarized at approximately - 2.0 V prior to running the anodic chronopotentiogram for the hydrogen halides, the iz ~ values are about 10~o smaller. Addition of Et4NOH to the solutions of all three hydrogen halides does not change the anodic i'r~, but the rest potential is shifted considerably as a function of pH. Controlled-potential coulometry has been used to determine the number of electrons involved in the oxidation of HC1 and HBr. Using a small electrolysis cell with a large platinum gauze electrode and a control potential of + 0.95 V, the number of electrons consumed during electrolysis, n, depends on the electrolysis time and varies between 8 and 20. If the electrolysis is interrupted after 50-70~o of the original oxidation wave has disappeared, n has a value between 6 and 10. The solutions turn yellow during electrolysis and a chronopotentiometric reduction wave appears at -0.5 V with a transition time which corresponds approximately to the number of coulombs consumed, assuming the formation of hydrogen ions. Coulometric reduction of the oxidation products changes the pH from 2.00 to 12.45 and evolves H2 gas (identified by gas chromatography). Because the only conceivable anodic electrode process is X----~½X2+e-, the coulometric experiments confirm that a catalytic process occurs which reproduces oxidizable material. During anodic electrolysis the solution pH changes from pH 3.6 for HC1 and pH 3.4 for HBr (1 x 10 -2 F solutions) to approximately pH 2.0 after 10-12 h. Again, hydrogen iodide is the exception; controlled-potential coulometry at + 0.70 V vs. SCE gives reasonable values for n. If the electrolysis time does not exceed 2 h, n is below 1.6 electrons; if the coulometric experiment is completed or interrupted within 1 h, n has a value between 1.04 and 1.18. The electrode process obviously is a one-electron process with higher n-values resulting from a slow chemical reaction of the products with the solvent. Voltammetric studies of the hydrogen halides in DMSO have been carried out using the same platinum electrode and the same cell as for chronopotentiometry. For HCI and HBr, cyclic voltammetry gives an anodic peak between + 0.90 and + 1.05 V and a reduction peak at -0.61 V at slow scan rates (Fig. 3). The two peaks separate further upon increasing the scan rate, and there is no indication of a reversible peak even at scan rates of 20 V s- 1. Reference to Fig. 3 indicates that the reduction peak is much smaller (and due to the I-IX reduction), if the first scan is cathodic. Hydrogen iodide gives an anodic peak at +0.65 V and a cathodic peak at +0.35 V plus a cathodic peak at -0.6 V which is not increased or decreased by the anodic process. Another anodic peak at + 0.45 V, which is very small for the first scan, becomes larger with each sweep. The peak separation indicates a reversible couple for this peak plus the cathodic peak at + 0.35 V. The appearance and behavior of the voltammograms as well as the chronopotentiograms is the same if 0.1 F LiC104 is used as supporting electrolyte instead of 0.1 F Et4NC104. The hydrogen halides have also been studied by chronopotentiometry in J. Electroanal. Chem., 23 (1969) 387-397
394
M. M I C H L M A Y R , D . T. S A W Y E R
+1.00
..-":....... " ./Z//
.if--:...+0.80
."
•
?
./"
+0.60
.~
....-- .~:.. /
~1~"
.i-"
/~I ./
.J ~:~-
~.~"J i
- 1.00
i
i
0.00
i
i
+1.00
i
+2.00
log i (i in pA)
Fig. 2. Plots ofEt= o vs. log i at constant concn, for HC1 in D M S O . (1-6): C O= 0.416, 0.830, 1.651, 3.00, 3.97, 6.34 m M .
TABLE 3 ANODIC TRANSFER COEFFICIENTS FOR THE OXIDATION OF HYDROGEN
HALIDES IN DMSO AT A PLATINUM
ELECTRODE (area, 0.227 cm 2) C O x 106/mol c m - 3
Et = o vs. Io# i Small i
Large i
Et = o vs. log C
(i/uA)
A. HC1
0.416 0.830 1.238 1.651 3.00 3.97 4.52 6.34 Average
0.56 0.50 0.42 0.56 0.60 0.48 0.52 0.55 0.52
0.35 0.30 0.28 0.27 0.16 0.13 0.12 0.10
0.53 0.47 0.50 0.56 0.51 0.39 0.35 0.33 0.51 (up to i = 30 #A)
(0.5) (1) (5) (10) (30) (50) (80) (100)
B. H B r
0.405 0.808 1.603 2.77 3.16 3.54 Average
0.60 0.62 0.52 0.58 0.64 0.61 0.59
0.51 0.44 0.38 0.30 0.25 0.25
0.65 0.59 0.59 0.67 0.53 0.48 0.63 (up to i = 20 #A)
(0.5) (1) (5) (20) (30) (50)
C. HI
0.415 0.979 1.795 2.48 3.08 3.59 Average
0.67 0.63 0.75 0.80 0.76 0.62 0.70
(1) (10) (20) (30) (40) (50)
J. Electroanal. Chem., 23 (1969) 387-397
0.70 0.76 0.74 0.80 0.72 0.74 0.74
395
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
N,N-dimethylformamide (DMF) and acetonitrile (CH3CN);the results are summarized in Table 4. In addition various halide salts have been investigated in DMSO, DMF and CH3CN. In general, their behavior is similar to that of the hydrogen halides (Table 4). The negative slopes for plots of i ¢ vs. i are indicative of the extent of the catalytic process.
+25
0
~-25
-50 -75
i
i
I
i
1.00
i
0.00
i
-1.00
E (V vs. SCE)
Fig. 3. Cyclic voltammograms of 2.00 x 10 -3 F HBr in DMSO containing 0.10 F Et4NCIO4 at a platinum electrode (area, 0.227 cmZ). (1,2,3) first, second, third potential sweep, respectively. Scan rate, 0.1 V s-1.
TABLE 4 CHRONOPOTENTIOMETRIC OXIDATION OF HYDROGEN HAL1DES AND ALKALI HALIDES IN C H 3 C N AT A PLATINUMELECTRODE
DMSO, D M F AND
(Values in parenthesis are for prewaves) DMSO E~/V vs. SCE
DMF Slope, i ¢ vs. i/s½
E~t/V vs. SCE
CH3CN Slope, ir ½ vs. i/s½
E~/V vs. SCE
Slope, i ¢ vs. i/s ½
+ 1.15 (+ 1.00) +0.98 ( + 0.70) +0.50 (+0.22) + 1.15
- 0.203
HC1
+ 1.00
-0.345
+ 1.03
-0.382
HBr
+0.88 ( + 0.60) +0.60 (+0.44) + 1.00
-0.170
+0.85 ( + 0.65) +0.65 (+0.48) +0.98
-0.234
HI LiC1
0 -0.310
0 -0.344
-0.141 0 -0.190
(+ 1.oo) KCI
+ 1.00
-0.293
+ 1.02
-0.328
KBr
+0.85 (+0.60} +0.60 ( + 0.45)
-0.126
+0.85 (+0.65) +0.62 ( + 0.50)
-0.165
KI
0
0
+ 1.15 (+0.98) +0.95 (+0.70) +0.50 (+ 0.20)
-0.184 -0.103 0
J. Electroanal. Chem., 23 (1969) 387-397
396
M. MICHLMAYR, D. T. SAWYER
DISCUSSION AND CONCLUSIONS
On the basis of the electrochemical data and the known behavior of hydrogen halides in nonaqueous systems, the overall process for the oxidation of HC1, HBr, and HI is concluded to be: HX ~=~ H + + X X - ~=~ X X z + e -
(8) (9)
The primary oxidation products are reactive, and solutions of halogens in CH3CN, CH3NO2 and D M F are known to be unstable 8'9. The reaction of CIE with DMSO has also been reported 2. Thus, catalytic post chemical reactions with the solvent and represented by: X
CH3"~- A CH2X\--2 + C H a / b - I D --*
CHa/b-13+
HX
(10)
or X2+CHa-CN
~
CHzX-CN+HX
(11)
are reasonable for the present systems. Such a cyclic mechanism is consistent with the large (Z/Zd)~ values (Table 1) and the coulometric results for HCI and HBr. The oxidation of hydrogen iodide appears to follow the same g~eral pattern, even though a significant increase of iz÷/AC° or of n is not observed. The fact that n is always larger than 1 supports the assumption that there is some reaction with the solvent. Another indication is the small transition time for the reverse wave (ratio 1 : 8 instead of 1 : 3); a part of the oxidation product has disappeared, probably through reaction with DMSO. However, the absence of a significant catalytic process implies that reaction of the oxidation product with the solvent does not regenerate HI. Although the oxidation of the halides corresponds to the formation of halogen, the mechanism is complicated by prewaves for HBr and HI. The sequential two-step mechanism suggested for a number of solvents 5 does not seem likely because the relative transition times are 1 : 5 instead of 4 : 5 for the prewaves. A possible explanation is that the prewave represents a first oxidation according to eqn. (1) which is followed by the reaction represented by eqn. (2) to give the main wave. Cyclic voltammetry and other data indicate that HI is oxidized, in part, by a reversible reaction (+0.45 and +0.35 V). I- ~:~ ½ I z + e I- +I2 --o I3 13 ~ 3 Ia 4-e-
(prewave)
(12)
(major wave)
(13) (14)
Thus, the 12 produced at the electrode reacts with I- to give I3, and I- and 13 are oxidized at different potentials. The same argument would apply to HBr and its prewave. ACKNOWLEDGEMENT
This work was supported by the National Science Foundation under Grant GP-7201. J. Electroanal. Chem., 23 (1969) 387-397
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
397
SUMMARY
The electrochemical behavior of anhydrous hydrogen chloride, hydrogen bromide and hydrogen iodide in DMSO has been studied by chronopotentiometry, controlled-potential coulometry, and cyclic voltammetry at platinum electrodes. The cathodic as well as the anodic reactions have been investigated, and the heterogeneous kinetic parameters for the rate-controlling electrode reactions have been evaluated. The data establish that oxidation products of the hydrogen halides react with the solvent. In the case of HCI and HBr, a catalytic process occurs which enhances the values of iz ~ by a factor of 5-10. REFERENCES l 2 3 4 5 6
7 8 9 10 11
12 13 14 15 16 17 18 19 20 21
G. J. JANZ AND S. S. DANYLUKin B. PESCE(Ed.), Electrolytes, Pergamon Press, Oxford, 1962, p. 225. A. V. THOMASAND E. G. ROCHOW, J. Am. Chem. Soc., 79 (1957) 1843. G. J. JANZ AND S. S. DANYLUK, J. Am. Chem. Soc., 81 (1959) 3846, 3850, 3854. J. A. OLABE, M. C. GIORDANOAND A. J. ARVIA, Electrochim. Acta, 12 (1967) 907. R. T. IWAMOTO,Anal. Chem., 31 (1959) 955. J. F. COETZEEAND WEI-SAN SIAO, Inorg. Chem., 2 (1963) 14; J. F. COETZEE AND J. L. HEDRICK, J. Phys. Chem., 67 (1963) 221. H. S. SWOFFORD,JR. AND J. H. PROPP, Anal. Chem., 37 (1965) 974; R. B. FULTONAND H. S. SWOFFORD, JR., Anal. Chem., 40 (1968) 1373. J. D. VOORHIESAND E. J. SCHURDAK,Anal. Chem., 34 (1962) 939; I. M. KOLTHOFFAND J. F. COETZEE, J. Am. Chem. Soc., 79 (1957) 1852. C. SINICKI, P. DESPORTES, M. BRYANT AND R. ROSSET, Bull. Soc. Chim. France, (1968) 829; R. L. BENOIT, M. GUAV AND J. DESBARRES, Can. J. Chem., 46 (1968) 1261. A. I. PoPov AND D. H. GESKE, J. Am. Chem. Soc., 80 (1958) 1340. R. A. OSTERYOUNG, m. L. McKISSON, P. H. DUTCH, G. LAUERAND E. B. LUCHSINGER,Development of a Light-weight Secondary Battery System, Final Report, Contract DA-36-039 SC-88925 (1962), AD 290 326 [ref. J. N. BUTLER, J. Electroanal. Chem., 14 (1967) 104]. M. C. GIORDANO, J. C. BAZANAND A. J. ARVXA,Electroehim. Acta, 11 (1966) 741, 1553. I. M. KOLTHOFF AND T. B. REDDY, J. Electroehem. Soc., 108 (1961) 980. H. DEHN, V. GUTMANN, H. KmCH AND G. SCHOEBER, Monatsh. Chem., 93 (1962) 1348. D. D. DEFORD, private communication, presented at the 133rd ACS Meeting, San Francisco, California, 1958. D. T. SAWYERAND J. L. ROBERTS, JR., J. Electroanal. Chem., 12 (1966) 90. M. v. STACKELBERG,M. P1LGRAMAND V. TOOME, Z. Elektrochem., 57 (1953) 342. P. DELAHAY,New Instrumental Methods in Electrochemistry, Interscience Publishers, Inc., New York, 1954. E. L. JOHNSON, K. H. POOL AND R. E. HAMM, Anal. Chem., 38 (1966) 183. W. M. LATIMER, The Oxidation States of the Elements and Their Potentials in Aqueous Solutions, Prentice-Hall, Inc., New York, 2nd ed., 1952, pp. 53, 60, 63. F.A. COTTONAND G. WILKINSON, Advancedlnorganie Chemistry, Interscience Publishers, New York, 2nd ed., 1966.
J. Electroanal. Chem., 23 (1969) 387-397
387
ELECTROCHEMISTRY OF HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
MANFRED MICHLMAYR AND DONALD T. SAWYER
Department of Chemistry, University of California, Riverside, California 92502 (U.S.A.) (Received October 10th, 1968; in revised form May 13th, 1969)
Although the polarographic and voltammetric reduction of the hydrogen halides (HC1, HBr and HI) has been studied extensively, almost no work has been reported for their electrochemical oxidation. In aqueous systems they are dissociated and only the oxidation of the halide ion or the formation of an insoluble mercurous salt by the mercury electrode is observed. However, hydrogen halides are much less dissociated in nonaqueous media than in water 1. Investigations of their electrical conductivity in anhydrous solvents 2- 4 have shown that the systems involve a number of complex equilibria such that the ionization constants obtained have no more than qualitative significance 1. The electrochemical oxidation of halide ions has been studied in several nonaqueous systems 5-10, but little is known about their electrochemical oxidation in dimethylsulfoxide (DMSO) 4'5,11,12. The reduction of HC1 in DMSO has been discussed 13a4. A general mechanism has been proposed for the oxidation of halides in nonaqueous media: 3 X- --0 X3 + 2eXg ~ } X 2 + e -
(1) (2)
based on the data for iodide ion. Two waves corresponding to the two oxidation reactions are observed at +0.5 and +0.7 V vs. SCE for iodide in DMSO 5. The reactivity of C12, Br2, and I2 with a number of organic solvents 8 is relevant to the present study because the reaction forms He1, HBr, and HI, respectively This paper summarizes the results of a detailed investigation of the electrochemistry of HCl, HBr, and HI in DMSO. From the studies, the kinetic parameters for the electrode reactions have been evaluated and the oxidation products have been identified; this has permitted plausible reaction mechanisms to be proposed. EXPERIMENTAL
Apparatus
The electrochemical measurements were made with a versatile electronic instrument constructed from Philbrick operational amplifiers following the design of DeFord is. By appropriate interconnection of the operational modules, the instrument can perform chronopotentiometry, controlled-potential coulometry and cyclic voltammetry. Potential-time curves were recorded with a Sargent model SR strip-chart recorder using a speed of 12 in./min, and cyclic voltammograms with a J. Electroanal. Chem., 23 (1969) 387-397
388
M. MICHLMAYR, D. T. SAWYER
Moseley X-Y recorder model 7030A. A function-generator (Hewlett-Packard model 202 A) was used for the cyclic voltammetry at high scan rates, and a Tektronix model 564 oscilloscope was employed to record these current-potential curves as well as chronopotentiograms with transition times of less than 1 s. The electrochemical cell and the three-electrode assembly used for chronopotentiometric and voltammetric measurements have been described previously 16. The electrochemical area of the platinum working electrode was determined using Fe(CN) 3- in 0.5 F aqueous KC1; its diffusion coefficient 17 was assumed to be 7.67 x 10-6 c m 2 s - 1 . Potentials are reported vs. SCE unless otherwise noted. The pH was monitored with a Leeds and Northrup miniature pH glass electrode assembly and a Corning model 12 pH meter. Reagents
A 0.1 F solution of tetraethylammonium perchlorate (Eastman Organic Chemical Co.) in DMSO (J. T. Baker Co.) was used as supporting electrolyte. Prior to use, the salt was recrystallized from water, dried at 70°C in vacuo and stored in a vacuum desiccator. For a number of experiments 0.1 F lithium perchlorate (G. F. Smith Chemical Co.) also was used as supporting electrolyte. Anhydrous hydrogen chloride, hydrogen bromide and hydrogen iodide were obtained from Matheson Co. ; their solutions were prepared by passing the gas through concentrated H 2 S O 4 prior to bubbling it slowly (with cooling) into DMSO. The concentrations were determined by potentiometric titrations with 0.01 F AgNOa and 0.01 F KOH. Solutions of H2SO4 and HC104 in DMSO were prepared as described previously 13. Acetonitrile (Mallinckrodt, nanograde) was purified by distillation from calcium hydride, and N,N-dimethylformamide (DMF) was obtained from Matheson, Coleman and Bell (spectroquality). All other materials were reagent-grade and were used without further purification. Either argon or prepurified nitrogen was used as the inert gas; the solutions were deaerated for 20 rain before measurements. The gas stream was presaturated by passing through a 0.1 F solution of EtgNC104 or LiC104 in DMSO, DMF, or CHaCN before entering the electrolysis cell. RESULTS
Anhydrous hydrogen halides give well-defined cathodic and anodic chronopotentiograms in D MSO at a platinum electrode (Fig. 1); the quarter-wave potentials for the reduction waves are virtually identical at - 0.50 V. The E~ values for the oxidation waves are +i.00 V, +0.88 V, and +0.60 V vs. SCE for HCI, HBr, and HI, respectively. 1. Reduction
The cathodic waves at -0.50 V correspond to the reduction of hydrogen ion, which is a diffusion-controlled process. This is established by the constant values of iT½/AC°; combination of these data with the Sand equation 18 permits evaluation of the diffusion coefficients, D, for the three acids. The respective values for HC1, HBr, and HI are 4.3 x 10 -6, 4.0 X 10 -6, and 3.7 x 10 - 6 c m 2 s -1 which are in good agreement with previously determined values 13'19 for HC1, HC104, and H2SO4. The diffusion coefficients of H2SO4 and HC104 also have been determined; both have J. Electroanal. Chem., 23 (1969) 387-397
389
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
1.00
i = 500 ~ A
r~
0.00 h
- 1.00
- 2.00
L
i
I
3 4 2 Time (see) Fig. 1. Chronopotentiograms for 3.00 × 10 -3 F HCI in D M S O containing 0.10 F Et4NC10~ at a platinum electrode (area, 0.277 cm2). 0
1
values in D M S O of 4.35 x 10 - 6 c m 2 s - 1 with H 2 S O 4 acting as a monobasic acid. The potential-time relationship for an irreversible system at 25°C is given b y I 8 : E - 0.059 log nFC°k~'hA 0.059 z~-t ~ ~na i + --~na l o g - - z ~
(3)
where E is the potential of the working electrode vs. NHE, ct the cathodic transfer coefficient, na the number of electrons in the rate-determining step of the reduction process, F the faraday, k°,h the heterogeneous rate constant for the reduction reaction, i the current, A the electrode area, C O the concentration of the electroactive species in the bulk of the solution in mol cm- 3, T the transition time and t the time after the electrolysis is started. This equation can be used to evaluate ~na and log k~,h by plotting E vs. log (r ~ - t~/z ~) for the chronopotentiometric waves. The respective values for HCI, HBr, and HI at a platinum electrode are 0.67, -5.14; 0.65, -4.96; and 0.64, -4.63. A more meaningful parameter for comparison of electrochemical kinetic data is the simple heterogeneous rate constant, ks.h. This is the rate constant at the equilibrium potential, E, for the rate-determining couple and can be evaluated from the relation 18 ks,h ---- kfOhexp (-- ~n~FE'/R T)
(4)
Because the formal potential for the reaction H ÷ + e- ~ ½ H 2 is not known in D M S O solutions, evaluation of ks,h requires that the value of E' be estimated. If the thermodynamic value of E ° for the reduction reaction in H 2 0 is used, ks, h is equal to kf0,hin view of the definition of NHE. J. Electroanal. Chem., 23 (1969) 387-397
390
M. MICHLMAYR, D. T. SAWYER
I
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391
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
2. Oxidation
In addition to the major anodic waves for the hydrogen halides in DMSO, HBr and HI exhibit small prewaves at +0.60 and +0.44 V vs. SCE, respectively. Their height is approximately ~ of the total transition time, i.e., much too small for consecutive two- and one-electron electrode reactions. The sum of the waves has been used for evaluations of the effect of current density and concentration on the transition time. Table 1 summarizes such data for the three halogen acids and indicates that the values of iz ~ for the oxidation wave of HC1 and HBr decrease with increasing current density; this normally is indicative of either a preceding chemical reaction or of a catalytic process 18. Consideration of the data in Fig. 1 and Table 1 indicates that the ir~ values for the anodic waves for HC1 and HBr are almost an order of magnitude larger than for the cathodic waves. This implies a catalytic process with regeneration of the hydrogen halide. Such a conclusion is supported by the data of Table 2 which establish TABLE 2 CHRONOPOTENTIOMETRIC DATA FOR THE OXIDATION OF HCI AND HBr AT HIGH CURRENT DENSITIES AT A PLATINUM ELECTRODE (area, 0.227 cm 2) IN DMSO HCl
HBr
C Ox 106/mol cm-3
i/#A
iz½/#A s½
C Ox 106/mol cm-3
i/#A
iz½/pA s~'
0.124 0.124 0.124 0.124 0.240 0.240 0.240 0.240 0.354 0.354 0.354 0.354 0.467 0.467 0.467 0.467 0.580 0.580 0.580 0.580
150 200 300 500 300 400 600 900 400 600 800 900 400 600 800 1000 500 700 800 1000
4.39 4.45 4.51 4.75 8.78 8.93 8.33 8.62 13.0 12.8 12.5 13.3 17.2 t6.6 17.6 16.7 20.6 20.8 21.0 19.8
0.186 0.186 0.186 0.186 0.370 0.370 0.370 0.370 0.554 0.554 0.554 0.554 0.738 0.738 0.738 0.738 0.922 0.922 0.922 0.922
200 300 400 500 300 400 500 550 400 500 600 700 450 600 700 800 500 600 700 900
7.20 7.75 8.30 7.75 15.1 15.4 14.9 14.3 24.2 21.2 21.7 22.0 30.4 28.9 30.0 30.6 39.4 39.7 40.1 39.3
iz ½
Average: A ~ 6
= 160 pA s½em mo1-1
iz ½ A C o - 181 #A s~cm mo1-1
that at high current densities the values of iz~/AC ° become independent of current and concentration and are reduced to values comparable to those for the reduction waves for HC1 and HBr. If the average values in Table 2 are assumed to represent a diffusion-controlled one-electron process that is uncomplicated by the catalytic J. Electroanal. Chem., 23 (1969) 387-397
392
M. MICHLMAYR, D. T. SAWYER
process, they can be used with the Sand equation is to evaluate the anodic diffusion coefficients for HC1 and HBr. The data for HI in Table 1 imply the absence of catalytic complications and can be used to evaluate its diffusion coefficient. The respective values of D for HC1, HBr, and HI are 3.5 x 10 -6, 4.5 x 10 -6, and 4.8 × 10 -6 cm2s -1. By assuming that the high current-density data for HC1 and HBr represent the diffusion-controlled process, the average values for i z ~ / A C ° can be used to calculate the hypothetical diffusional transition times, Zd, for the conditions of Table 1 if there were not catalytic complications. The observed transition times are representative of the catalytic process and can be combined with the calculated Zd values to give the quantity (Z/Zd)+ which is indicative of the extent of the catalytic reaction (Table 1). Delahay 18 has shown that when the value of this ratio is greater than 2.4, a firstorder catalytic process obeys the expression (5)
(Z/Zd) ~ = 2(kcCzz/rc) ~
where kc is the catalytic rate constant and Cz the concentration of the substance that is catalytically oxidized. A plot of (Z/Zd) ~ VS. Z~ should give a linear slope that is proportional to (kc Cz)~. Such a treatment of the data in Table 1 yields curved lines whose slopes decrease with increasing values of z ~ and indicates a more complicated process than a first-order reaction. The values of (k~ C~)½ calculated from the slopes range from 1.1 to 0.5 for HC1 and from 0.3 to 0.1 for HBr. When the anodic current for the hydrogen halides is reversed before reaching the transition time, a reduction wave appears at -0.50 V (Fig. 1) which has the general features of the cathodic hydrogen ion reduction wave (~na=0.67). The ratio of its transition time to the electrolysis time before reversing the current is 1 : 3 for HCI, and between 1:2.5 and 1:3 for HBr. Current reversal prior to the transition time for the anodic HI wave yields a cathodic wave with an E0.22 value of +0.36 V. The ratio of zr/zf is always smaller than the theoretical 1 : 3 and averages about 1 : 8. When the reduction wave at +0.36 is reversed, the anodic wave at + 0.44 (small prewave) is found to be larger than before. For this case, the ratio of the transition times is 1 : 3 (after subtraction of the previous wave). The cathodic wave at - 0.5 V does not show any net increase in ir ~ on a reverse chronopotentiogram. A relation similar to eqn. (3) has been developed for irreversible anodic chronopotentiograms 1s E -
0.059 n F C ° k°,h A (1 - ct)na log i
0.059 z~ - t ~ (1 - c~)n~log - - z ~
(6)
where (1 - a ) is the anodic transfer coefficient and kbOhthe heterogeneous rate constant for the oxidation reaction. By extrapolating the potential to the beginning of the chronopotentiogram, plots of Et=o vs. log i and Et=o vs. log C o should give straight lines with slopes that allow evaluation of ( 1 - a)n a. Figure 2 and Table 3 indicate that for low current densities such plots give meaningful results; at higher current densities (with a transition time of 30 s, or less) irregular (1 - a)na values are obtained for HCI and HBr. By using the average values of (1 - ~ ) n a (0.52, 0.61 and 0.72 for HC1, HBr, and HI, respectively) and assuming that n is unity, the quantity log k°h has been evaluated from eqn. (6); the mean values are - 13.1, - 13.3 and - 12.0 for the three hydrogen halides. J. Electroanal. Chem., 23 (1969) 387-397
393
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
Again, ks,h can be calculated from the relation: ks,h = k°,h exp [(1 - a ) n a F E ' / R T ]
(7)
by using the thermodynamic values, E °, for the half reactions X- ~ ½X 2 + e- in H 2 0 (+ 1.36, + 1.07, and +0.54 V vs. NHE2°). The respective values for log ks,h are -2.0, -2.9, and -5.5 for HC1, HBr and HI. If the platinum electrode is polarized at approximately - 2.0 V prior to running the anodic chronopotentiogram for the hydrogen halides, the iz ~ values are about 10~o smaller. Addition of Et4NOH to the solutions of all three hydrogen halides does not change the anodic i'r~, but the rest potential is shifted considerably as a function of pH. Controlled-potential coulometry has been used to determine the number of electrons involved in the oxidation of HC1 and HBr. Using a small electrolysis cell with a large platinum gauze electrode and a control potential of + 0.95 V, the number of electrons consumed during electrolysis, n, depends on the electrolysis time and varies between 8 and 20. If the electrolysis is interrupted after 50-70~o of the original oxidation wave has disappeared, n has a value between 6 and 10. The solutions turn yellow during electrolysis and a chronopotentiometric reduction wave appears at -0.5 V with a transition time which corresponds approximately to the number of coulombs consumed, assuming the formation of hydrogen ions. Coulometric reduction of the oxidation products changes the pH from 2.00 to 12.45 and evolves H2 gas (identified by gas chromatography). Because the only conceivable anodic electrode process is X----~½X2+e-, the coulometric experiments confirm that a catalytic process occurs which reproduces oxidizable material. During anodic electrolysis the solution pH changes from pH 3.6 for HC1 and pH 3.4 for HBr (1 x 10 -2 F solutions) to approximately pH 2.0 after 10-12 h. Again, hydrogen iodide is the exception; controlled-potential coulometry at + 0.70 V vs. SCE gives reasonable values for n. If the electrolysis time does not exceed 2 h, n is below 1.6 electrons; if the coulometric experiment is completed or interrupted within 1 h, n has a value between 1.04 and 1.18. The electrode process obviously is a one-electron process with higher n-values resulting from a slow chemical reaction of the products with the solvent. Voltammetric studies of the hydrogen halides in DMSO have been carried out using the same platinum electrode and the same cell as for chronopotentiometry. For HCI and HBr, cyclic voltammetry gives an anodic peak between + 0.90 and + 1.05 V and a reduction peak at -0.61 V at slow scan rates (Fig. 3). The two peaks separate further upon increasing the scan rate, and there is no indication of a reversible peak even at scan rates of 20 V s- 1. Reference to Fig. 3 indicates that the reduction peak is much smaller (and due to the I-IX reduction), if the first scan is cathodic. Hydrogen iodide gives an anodic peak at +0.65 V and a cathodic peak at +0.35 V plus a cathodic peak at -0.6 V which is not increased or decreased by the anodic process. Another anodic peak at + 0.45 V, which is very small for the first scan, becomes larger with each sweep. The peak separation indicates a reversible couple for this peak plus the cathodic peak at + 0.35 V. The appearance and behavior of the voltammograms as well as the chronopotentiograms is the same if 0.1 F LiC104 is used as supporting electrolyte instead of 0.1 F Et4NC104. The hydrogen halides have also been studied by chronopotentiometry in J. Electroanal. Chem., 23 (1969) 387-397
394
M. M I C H L M A Y R , D . T. S A W Y E R
+1.00
..-":....... " ./Z//
.if--:...+0.80
."
•
?
./"
+0.60
.~
....-- .~:.. /
~1~"
.i-"
/~I ./
.J ~:~-
~.~"J i
- 1.00
i
i
0.00
i
i
+1.00
i
+2.00
log i (i in pA)
Fig. 2. Plots ofEt= o vs. log i at constant concn, for HC1 in D M S O . (1-6): C O= 0.416, 0.830, 1.651, 3.00, 3.97, 6.34 m M .
TABLE 3 ANODIC TRANSFER COEFFICIENTS FOR THE OXIDATION OF HYDROGEN
HALIDES IN DMSO AT A PLATINUM
ELECTRODE (area, 0.227 cm 2) C O x 106/mol c m - 3
Et = o vs. Io# i Small i
Large i
Et = o vs. log C
(i/uA)
A. HC1
0.416 0.830 1.238 1.651 3.00 3.97 4.52 6.34 Average
0.56 0.50 0.42 0.56 0.60 0.48 0.52 0.55 0.52
0.35 0.30 0.28 0.27 0.16 0.13 0.12 0.10
0.53 0.47 0.50 0.56 0.51 0.39 0.35 0.33 0.51 (up to i = 30 #A)
(0.5) (1) (5) (10) (30) (50) (80) (100)
B. H B r
0.405 0.808 1.603 2.77 3.16 3.54 Average
0.60 0.62 0.52 0.58 0.64 0.61 0.59
0.51 0.44 0.38 0.30 0.25 0.25
0.65 0.59 0.59 0.67 0.53 0.48 0.63 (up to i = 20 #A)
(0.5) (1) (5) (20) (30) (50)
C. HI
0.415 0.979 1.795 2.48 3.08 3.59 Average
0.67 0.63 0.75 0.80 0.76 0.62 0.70
(1) (10) (20) (30) (40) (50)
J. Electroanal. Chem., 23 (1969) 387-397
0.70 0.76 0.74 0.80 0.72 0.74 0.74
395
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
N,N-dimethylformamide (DMF) and acetonitrile (CH3CN);the results are summarized in Table 4. In addition various halide salts have been investigated in DMSO, DMF and CH3CN. In general, their behavior is similar to that of the hydrogen halides (Table 4). The negative slopes for plots of i ¢ vs. i are indicative of the extent of the catalytic process.
+25
0
~-25
-50 -75
i
i
I
i
1.00
i
0.00
i
-1.00
E (V vs. SCE)
Fig. 3. Cyclic voltammograms of 2.00 x 10 -3 F HBr in DMSO containing 0.10 F Et4NCIO4 at a platinum electrode (area, 0.227 cmZ). (1,2,3) first, second, third potential sweep, respectively. Scan rate, 0.1 V s-1.
TABLE 4 CHRONOPOTENTIOMETRIC OXIDATION OF HYDROGEN HAL1DES AND ALKALI HALIDES IN C H 3 C N AT A PLATINUMELECTRODE
DMSO, D M F AND
(Values in parenthesis are for prewaves) DMSO E~/V vs. SCE
DMF Slope, i ¢ vs. i/s½
E~t/V vs. SCE
CH3CN Slope, ir ½ vs. i/s½
E~/V vs. SCE
Slope, i ¢ vs. i/s ½
+ 1.15 (+ 1.00) +0.98 ( + 0.70) +0.50 (+0.22) + 1.15
- 0.203
HC1
+ 1.00
-0.345
+ 1.03
-0.382
HBr
+0.88 ( + 0.60) +0.60 (+0.44) + 1.00
-0.170
+0.85 ( + 0.65) +0.65 (+0.48) +0.98
-0.234
HI LiC1
0 -0.310
0 -0.344
-0.141 0 -0.190
(+ 1.oo) KCI
+ 1.00
-0.293
+ 1.02
-0.328
KBr
+0.85 (+0.60} +0.60 ( + 0.45)
-0.126
+0.85 (+0.65) +0.62 ( + 0.50)
-0.165
KI
0
0
+ 1.15 (+0.98) +0.95 (+0.70) +0.50 (+ 0.20)
-0.184 -0.103 0
J. Electroanal. Chem., 23 (1969) 387-397
396
M. MICHLMAYR, D. T. SAWYER
DISCUSSION AND CONCLUSIONS
On the basis of the electrochemical data and the known behavior of hydrogen halides in nonaqueous systems, the overall process for the oxidation of HC1, HBr, and HI is concluded to be: HX ~=~ H + + X X - ~=~ X X z + e -
(8) (9)
The primary oxidation products are reactive, and solutions of halogens in CH3CN, CH3NO2 and D M F are known to be unstable 8'9. The reaction of CIE with DMSO has also been reported 2. Thus, catalytic post chemical reactions with the solvent and represented by: X
CH3"~- A CH2X\--2 + C H a / b - I D --*
CHa/b-13+
HX
(10)
or X2+CHa-CN
~
CHzX-CN+HX
(11)
are reasonable for the present systems. Such a cyclic mechanism is consistent with the large (Z/Zd)~ values (Table 1) and the coulometric results for HCI and HBr. The oxidation of hydrogen iodide appears to follow the same g~eral pattern, even though a significant increase of iz÷/AC° or of n is not observed. The fact that n is always larger than 1 supports the assumption that there is some reaction with the solvent. Another indication is the small transition time for the reverse wave (ratio 1 : 8 instead of 1 : 3); a part of the oxidation product has disappeared, probably through reaction with DMSO. However, the absence of a significant catalytic process implies that reaction of the oxidation product with the solvent does not regenerate HI. Although the oxidation of the halides corresponds to the formation of halogen, the mechanism is complicated by prewaves for HBr and HI. The sequential two-step mechanism suggested for a number of solvents 5 does not seem likely because the relative transition times are 1 : 5 instead of 4 : 5 for the prewaves. A possible explanation is that the prewave represents a first oxidation according to eqn. (1) which is followed by the reaction represented by eqn. (2) to give the main wave. Cyclic voltammetry and other data indicate that HI is oxidized, in part, by a reversible reaction (+0.45 and +0.35 V). I- ~:~ ½ I z + e I- +I2 --o I3 13 ~ 3 Ia 4-e-
(prewave)
(12)
(major wave)
(13) (14)
Thus, the 12 produced at the electrode reacts with I- to give I3, and I- and 13 are oxidized at different potentials. The same argument would apply to HBr and its prewave. ACKNOWLEDGEMENT
This work was supported by the National Science Foundation under Grant GP-7201. J. Electroanal. Chem., 23 (1969) 387-397
HYDROGEN HALIDES IN DIMETHYLSULFOXIDE
397
SUMMARY
The electrochemical behavior of anhydrous hydrogen chloride, hydrogen bromide and hydrogen iodide in DMSO has been studied by chronopotentiometry, controlled-potential coulometry, and cyclic voltammetry at platinum electrodes. The cathodic as well as the anodic reactions have been investigated, and the heterogeneous kinetic parameters for the rate-controlling electrode reactions have been evaluated. The data establish that oxidation products of the hydrogen halides react with the solvent. In the case of HCI and HBr, a catalytic process occurs which enhances the values of iz ~ by a factor of 5-10. REFERENCES l 2 3 4 5 6
7 8 9 10 11
12 13 14 15 16 17 18 19 20 21
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J. Electroanal. Chem., 23 (1969) 387-397