On the Anodic Oxidation of Hydroxide Ion in Acetonitrile and Dimethyl Sulfoxide L. A. Simonson'
and
Royce W. Murray
Kenan Laboratories of Chemistry, University of North Carolina, Chapel Hill, N.C. 275 14
The electrochemical Oxidation of hydroxide Ion at Pt in dimethyl sulfoxide and acetonitrile solvents is described. Two anodic waves and a connected cathodic one in acetonitrile correspond to surface oxide formation and stripping. Two additional more positive anodic waves are diffusional and have collectively n = 1. The surface waves are missing in dimethyl sulfoxide. Identification of electrolysis products eluded a detailed search, including a search for O2 and hydroxyl radical trapping experiment.
This research was inspired by a report by Goolsby and Sawyer (I ) that the oxidation of hydroxide ion a t Pt electrodes in dimethyl sulfoxide solvent is a one-electron process cleanly producing molecular oxygen. The proposed mechanism was:
OH-
-
OP
-
+
e-
20H* + HZOz 02 + HzO
2H202
(1) (2)
(3)
The purported intermediacy of the reactive hydroxyl radical was of interest to us inasmuch as, if this radical could be captured by added substrates in a manner predictable by known hydroxyl radical chemistry, reaction 1 would possess a substantial electrosynthetic potential. Additionally, the production of oxygen, if monitored, could provide the basis for measuring relative rates of hydroxyl radical scavenging by substrates for comparison to the large body of pulse radiolysis data available (2) on such scavenging. Finally, such studies could shed further light on the extent to which the nonaqueous-produced hydroxyl radical is surface-bound, as opposed to a diffusing, reactive species. In the Hofer-Moest reaction in aqueous alkali, for example, it is thought to be surface bound ( 3 , 4 ) . Our investigation has included cyclic voltammetry, coulometry, and attempts a t product isolation following exhaustive electrolysis. Experiments were carried out principally with Pt electrodes, and in dimethyl sulfoxide and acetonitrile solvents, with emphasis on acetonitrile. Our experimental results show that oxygen is not a product of hydroxide ion oxidation in these solvents.
EXPERIMENTAL Cells and Apparatus. The cell employed for cyclic voltammetry has been described ( 5 ) . The working electrode was a fused Pt bead of area 4 X cm2; the auxiliary electrode, a Pt wire spiral; and the reference electrode, an isolated Ag/O.lM AgNO&H&N electrode. In reporting potentials, data were converted from this reference electrode to the aqueous SCE, which lies --0.30 V negative of the silver reference. The cyclic voltammetric cell was operated at ambient temperature (24 f 1 "C). In DMSO, an aqueous SCE reference electrode was used directly. Deaeration was with either dried N2 or He (in bulk electrolysis experiments). The bulk electrolysis cell was H-type fitted with stopcocks and ground glass accessories such that it could be completely sealed for Present address, Department of Chemistry, Framingham State College, Framingham, Mass. 01701. 290
headspace gas analysis. The anode working department typically contained 35 ml of reactant solution; the headspace volume was 40 ml. The anode compartment with Pt mesh basket working electrode of large but undetermined area, was isolated from the cathode department by a medium porosity glass frit. Gas evolved from the Pt flag auxiliary electrode in the cathode department was bled off as necessary. Reference electrode connection to the anode department was through a capillary salt bridge filled with the nonaqueous solvent system. The cyclic voltammetric system has been described ( 6 ) ;that for exhaustive electrolysis of BudNOH solutions employed a conventional operational amplifier controlled module in combination with a booster amplifier ( 7 ) .Coulometry was carried out by integration of current-time curves or with a gas coulometer (81. Gas chromatography was with a Gow/Mac Model 69-500 thermal conductivity detection. Chemicals and Supplies. Acetonitrile (MCB-spectrograde),dimethyl sulfoxide (Baker-reagent), and BudNOH (Eastman-10% aqueous) were used as received. Et4NC104 (Eastman) was recrystallized twice from water, dried under vacuum at 60 "C, and stored in a desiccator. The Bu4NOH was standardized (0.398M) by titration with aqueous HCIO4. Total initial water concentration in electrochemical experiments was approximately 1.3% (-O.'72M), owing to the aqueous nature of the BudNOH reagent. Attempts to reduce the water concentration of the BudNOH by addition of acetonitrile to the 10% aqueous solution followed by evaporation under reduced pressure to a small volume were unsuccessful. This procedure produced a solution with a characteristic amine odor, indicating Hoffman degradatign. Product Analysis. Analysis for 0 2 in the headspace of the anodic electrolysis compartment with both DMSO and acetonitrile solvents was with gas-solid chromatography using a 3-f00t, 5-8, Molesieve 60-80 mesh column activated a t 250 "C overnight and operated a t 50 "C with H e as carrier gas. Adequate sensitivity of the analysis was proved by the following. A 1O-pl air sample (2 PI 0 2 ) produced a symmetrical 02 peak of height 24 units, with associated noise of 0.5 unit. Taking one unit as detection limit, or 0.1 pl 02,this is equivalent to a 5O-pl sample containing 0.2Oh 0 2 . The volume of 0 2 formed by a typical experiment, if reactions 1-3 held, would be 2.2 ml; if 50% of this escapes into the 40-ml headspace, the headspace 02 concentration (2.8%) exceeds the detection limit by 14-fold. The assumption of adequate escape to the headspace was proved valid by decomposing an equivalent quantity of H ~ 0 2 (with KI) in the same cell, with observation of a very strong 02 chromatographic peak. The electrolysis solutions were covered with He, which, since the chromatographic method also detects N2. made detection of air leaks simple. No difficulties were experienced with such leaks during typical experiments. Analysis for H202 was conducted by addition of aqueous saiurated KI to the electrolysis product solution and measuring evolved 0 2 . No 02 was thus detected, nor did the brown discoloration of liberated I2 appear, in any of the described electrochemical experiments. Gas chromatographic inspection for succinonitrile, acetamide, carbon dioxide, ethylene, and butylene, was on Poropak-Q six-foot columns operated a t 225, 225, 65, 170, and 170 "C, respectively. Samples were taken from both headspace and anode compartment solution. Acetamide was found but was traced to base-assisted hydrolysis of acetonitrile and is not a faradaic product.
RESULTS AND DISCUSSION Oxidation of Hydroxide Ion in Dimethyl Sulfoxide. Results of cyclic voltammetry of tetrabutylammonium hydroxide (Bu4NOH) in DMSO with 0.1M Et4NC104 are shown in Figure 1. The hydroxide cyclic voltammogram is characterized by 2 irreversible anodic waves at f0.25 and
ANALYTICAL CHEMISTRY, VOL. 47, NO. 2 , FEBRUARY 1975
I
1
d --
-n
I
Y
v+===7-5Fa
1 1 IO
0 5
-0.5
0 Volts
v5
I
-IO
-I5
S C E
Figure 1. Cyclic voltammograms in DMSO (0.1MEt4NC104) at Pt and
0.3 Vlsec Curve A, background: Curve 6,air-saturated: Curves C and D, 2.5mM Eu~NOH
+0.65 V, preceded by an incompletely resolved smaller wave a t approximately 0 V. At +0.85 V, the current rises continuously into background. A cyclic voltammogram of dissolved 0 2 is also shown in Figure 1, the pseudo-reversible superoxide-producing cathodic wave appearing a t -0.75 V. This wave (or any other) clearly does not appear on the cathodic-going cycle of the hydroxide oxidation voltammogram (Curve C ). These data thus indicate that oxygen is not formed during the oxidation of hydroxide ion in DMSO. Further substantive evidence concerning O2 as a nonproduct was obtained by exhaustive electrolysis experiments a t controlled potential; 5.lOmM Bu4NOH in 0.1M Et4NC104 was electrolyzed at +0.9 V a t a Pt electrode in a closed, deareated cell. Oxygen was not detected during or after completion of the electrolysis using a sensitive gas chromatographic technique (See Experimental Section). Independent recent re-investigation of the oxidation by Sawyer (9) has confirmed that the product is something other than oxygen. The current-time behavior of the bulk electrolysis showed that log(i) was not linear with time, indicating the electrolysis was not simple first order. Integration of the current-time curve gave a coulometric n of 0.85, which agrees with the value of Goolsby and Sawyer. It is ostensibly somewhat low due to a mild instability of hydroxide ion in DMSO. The bulk electrolysis solution undergoes a definite (brown) darkening during the electrolysis. Because of the complications posed by the hydroxide decomposition, we did not attempt further isolation of reaction products. A brief mass transport characterization was carried out for the two anodic hydroxide waves shown in Figure 1. Both appear to be diffusion-controlled, yielding constant iplu1’2 over a range of sweep rates, u = 0.05-1.0 Vlsecond. In one series of scans, iP/u1l2 ac = 192 f 4 and 332 f 10 A cm mol-1V-1’2. Both values are much smaller than expected for a diffusional one-electron reaction, which is consis1 for the s u m of the tent with the coulometric result of n two waves. The day-to-day distribution of current between the two anodic waves was not especially reproducible; see for instance Curves C and D of Figure 1 obtained on different days. Goolsby and Sawyer ( I ) observed a single (+0.75 V) oxidation wave for hydroxide. Recent results (9) suggest that solution aging effects may attenuate the first wave.
-
L
20
1
IO
-I 0
0
V o l t s vs
SC E
Figure 2. Cyclic voltammograms in 0.1 M Et4NC104/acetonitrileat Pt Current scales for curves are 40 (A), 40 (E), 40 (C). 100 (D) PA. Curve A, oxygen reduction (air-saturated) at 0.3 V/sec. Curve 8, 6mM (--) and 12mM (- - - - -) Eu4NOH at 0.2 V/sec. Curve C, 6 m M 6u4NOH at 0.4. V/sec with reversal for second cycle (- - -) at -0.4V. Curve D, 6 m M Bu4NOH at 0.3 Vlsec with stop scan (2 minutes), then reverse scan at +0.05 V ( a ) .4-0.50 V. ( b ) .and f 1 . 1 0 v (c).Relative areas of resutting peak V are 1.00 ( a ) ,1.6 ( b ) , 1.6 (c),1.64 ( d )
-
Clearly, hydroxide behavior in this solvent is more complex than originally thought. Oxidation of Hydroxide in Acetonitrile. Tetrabutylammonium hydroxide is quantitatively stable, as assessed titrimetrically in acetonitrile for a t least 2 hours; longer periods of use were not needed and were not stability-tested. Solutions displayed no odor of tributylamine (Hoffman degradation product). Given this stability, detailed characterization of hydroxide ion oxidation in acetonitrile could be reasonably attempted. Cyclic voltammograms of Bu4NOH a t Pt in acetonitrile are shown in Figure 2. The general reproducibility of the relative peak heights of the anodic pattern was only fair from day to day; the results shown are typical of a large number of experiments. The anodic electrochemistry of hydroxide in acetonitrile is clearly richer than was the case in DMSO. Approximate positions of the four anodic current peaks shown are peak I (0 V), peak I1 (+0.6 V), peak I11 (+1.0 V), and peak IV (+1.6V). The single cathodic peak V lies at -1.0 V us. SCE. This richness does not include a product wave for molecular 0 2 (compare Curve A with others). It was quickly apparent from the behavior of the anodic waves that surface limitations are involved in hydroxide ion oxidation in acetonitrile. The literature on the oxidation of Pt electrodes in aqueous media is quite extensive. Aside from scans into background in acetonitrile ( I O ) , there appears to have been relatively little attention given to Pt electrode (or hydroxide) oxidation in nonaqueous media. In aqueous solution, scrupulous conditions are required to observe fine structure in the anodic scan at a Pt electrode (11, 12). The fine structure in the anodic patterns
A N A L Y T I C A L C H E M I S T R Y , VOL. 47,
NO. 2,
FEBRUARY
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C u r r e n t Peak 11. Our characterization of current peak I1 includes the following observations: (i) The peak current for current peak I1 is the most variable of the four anodic
Table I. Potential Sweep Peak Currents for 6mM Bu4NOH in 0.1M Et4NC104/CH3CN a t P t Current peak,
I
Sweep rate
v, se;
0.10 0.20 0.30 0.40 0.50 0.60
0.70 0.80 0.90 1.00
ui?
iD
8 16 23 30 35 40 50 55 60 67
i
I?
80 80 77 75 70 67 71 67 67 67
I p f k
17 24 33 38 45 55 d
170 120 110 95 90 92 d
54 54 60 60 64 71 d
... ... ... ... ... ... ... ... .. .
28 45 57 64 73 80 87 95 103 108
88 101 104 101 103 103 104 106 108 108
18 33 47 61 76 85 105 117 130 150
180 167 157 152 152 142 150 146 144 150
a Start at +0.1 V (after stirring 60 sec). Start a t +0,6V (after stirring 120 sec). Potential swept through Peaks I-IV during anodic cycle. Severe interference from Peak 111.
of Figure 2 is obtained casually. This suggests that basic nonaqueous media might have some usefulness in electrode surface studies. Accordingly, we have attempted an overall characterization of the complex anodic hydroxide oxidation pattern into surface and nonsurface components. During these studies, the working electrode (Pt bead) was subjected only to mechanical and then electrochemical cleaning. Relatively reproducible results were obtained by sweeping the electrode negative (by approximately 0.4 V) of cathodic peak V before initiating the anodic sweep from -1.4 V us. SCE. That peak V did not appear in this “cleaning” operation indicates it results from anodic processes during the main anodic sweep. C u r r e n t Peaks I ‘and V. These current peaks are considered together since they are both surface-controlled. The following comprises the evidence for surface control: (i) At constant scan rate, current peaks I and V are independent of the concentration of Bu4NOH (See Curve B, Figure 2). The potential of current peak V does depend, extra-thermodynamically, on the hydroxide concentration. (ii) Scanning through peaks I-IV (or only through peak I), reverse scanning, halting at - 0.5 V for a period of stirring, and then continuing the cathodic sweep does not affect current peak V. (iii) The height of current peaks of I and V are linearly related to potential scan rate, as shown in Table I. The small increase in the ratio i,/u at low sweep rate u is due to an apparent inadequacy in the background current correction; a plot of i us. u for current peaks I and V is nicely linear with a small current intercept a t u = 0. These data are interpreted as corresponding to formation of a surface film of surface area or site-limited quantity in anodic current peak I, and to cathodic stripping in current peak V of this film plus any further film formed in anodic current peaks 11-IV. That the anodically-formed film must be stripped via peak V before further peak I film can be formed is clearly shown by Curve C of Figure 2, where peak I is absent on a second anodic-going potential scan. That the anodically formed film in peak I is a Pt oxide film is the most obvious further interpretation. We have, however, no evidence to distinguish the type of oxide film present and will refrain from speculating on this point. We do find, from experiments in which the anodic potential sweep through peak I is stopped and then reversed at 0.05 V, the charge under the resulting cathodic current peak V (Curve D,a, Figure 2) from the film formed in peak I is reproducibly about 610 pc/cm2, more than enough charge for a monolayer of oxide even assuming a roughness factor of two.
,
292
hydroxide oxidation peaks, on occasion being larger than current peaks I11 and IV. (ii) Current peak I1 is somewhat enhanced by increased hydroxide concentration, as shown by Curve B, Figure 2, suggesting a mass transport-limited component. (iii) When the electrode potential sweep is stopped at +0.6 V, the current for peak I1 decays rather quickly to a small value which is increased only slightly by stirring the solution, indicating surface control. (iv) The potential sweep rate dependency of current peak I1 was investigated by potentiostatting the electrode a t +0.05 V (allowing the current peak I to decay) and then scanning anodically at a series of sweep rates. The resulting current peak data, shown in Table I, show that neither i ,/u nor i p/ u1I2 is constant. (v) Additional surface film is formed by scanning the potential through peak 11. This is demonstrated by Curve D,b, Figure 2, which shows that the cathodic stripping peak V grows by approximately 60% (and shifts cathodically somewhat) as a result of scanning through current peak I1 as compared to scanning through current peak I only. Collectively, these results show that current peak I1 definitely contains a surface-limited film-forming component, probably as its major component. Some diffusional contribution to current peak I1 is also indicated, however, through observation (ii) and (iv). We believe that the current peak 11 is actually exclusively surface-limited, and that the apparent mass transport control of it arises from its overlap with current peak 111, which is rather broad and, as will be shown below, is mass transport-limited. That the surface film formed in cathodic peak I1 is again a “Pt oxide” is the most reasonable interpretation but, in the absence of further analysis, we will refrain from speculating on the type of oxide layer or on the origin of multiple, oxide-forming current peaks. C u r r e n t Peaks I11 a n d IV. Current peak I11 appears to be largely mass-transport as opposed to surface-limited. Thus, reversing the potential scan following current peak I11 does not increase the magnitude of the film-stripping peak V as compared to that produced by anodic peaks I and 11. Additionally, current peak I11 is increased substantially by stirring the solution and by increasing the concentration of hydroxide ion, and in experiments where peaks I and I1 are pre-electrolyzed exhibits a fairly constant value of i , / ~ ’ / (Table ~ I). Current peak I11 is diminished in a second anodic scan (Curve C, Figure 2), but this is to be expected simply from hydroxide ion depletion around the Pt electrode surface. As noted above, an overlap of this peak with current peak I1 would lend an apparent mass transport-limited component to the latter. We have only qualitative observations for current peak IV, for which the background current correction is appreciable and which is overlapped by the tail of current peak 111. Current peak IV is increased significantly by stirring the solution, and is increased at higher concentration of hydroxide ion in the solution. These qualitative observations indicate a mass transport-limited character for this current peak. Anodic Coulometry. Controlled potential coulometry was carried out on Bu4NOH solutions ranging in concentration from 6-12mM at +0.8, +0.9, +1.0, $1.1, and +1.7 V us. SCE. Results for n are 0.67, 0.75, 0.72, 0.70, and 1.03, respectively. Thus, about three quarters of the hydroxide ion is consumed by electrolysis at any potential impinging on current peak 111, whereas a complete one-electron consumption of hydroxide is effected by electrolysis at the potential of current peak IV. The electrolytic decay curves
ANALYTICAL CHEMISTRY, VOL. 47, NO. 2, FEBRUARY 1975
were first order a t +1.7 V [log ( i ) us. t linear], but showed a definite plateau a t lower potentials. An example is shown in Figure 3. Search for Products of Hydroxide Ion Oxidation. The one-electron coulometric result indicates that the primeval step in the anodic reaction of hydroxide ion is reaction 1, producing a diffusing or surface-bound hydroxyl radical. In either case, the hydroxyl radical is expected t o be exceedingly reactive and the irreversibility of the cyclic voltammetric pattern is not surprising. We have attempted to indirectly detect the hydroxyl radical by (i) isolation from exhaustively electrolyzed solutions of products which have hydroxyl radical as a logical precursor, and (ii) trapping the hydroxyl radical with substrates known to be kinetically more reactive toward t,he radical than toward acetonitrile. Results of both approaches are described below and are negative. Following reactions 1-3 suggested by Goolsby and Sawyer ( I ) for hydroxide oxidation, a scrutinous inspection for oxygen and hydrogen peroxide product was carried out. Preliminary evidence that 0 2 was not produced has already been shown in the cyclic voltammogram of Figure 2. T h e headspace of the electrolysis cell was examined for 0 2 by a gas chromatographic technique (see Experimental) during and following exhaustive electrolysis a t potentials ranging from H . 8 t o +1.7 V us. SCE, emphasizing the potentials +.l.l and $1.7 V. In no instance was 0 2 detected, even in traces. The gas chromatographic method had sensitivity for 0 2 detection of 0.1 ~1 absolute. This sensitivity exceeds by a factor of 30-fold t.hat required to detect 0 2 evolved in the electrolysis cell headspace via reactions 1-3 for a typical Bu.lNOH sample. This sensitivity was verified in a practical sense by decomposing an equivalent quantity of HZOZ, by added HI, in the electrolysis cell, whereby sufficient oxygen was evolved t o yield a clearly measurable chromatographic peak with signalhoke > 10. At the same time, a discoloration of the solution due to iodine production was observed, as expected. T h e above experiment eliminates reaction 3 in acetonitrile but does not preclude the eventuality that reactions 1 and ?: occur to yield hydrogen peroxide as a stable product or as a product which decomposes by a non-On-producing pathw;iy. Accordingly, an inspection was conducted for hydrogen peroxide in each of the above electrolyses by the addition of KI a t the completion of the electrolysis. In no case was the distinctive darkening of iodine production observed. Hydrogen peroxide in neutral acetonitrile is stable for long periods o f time (many hours). In hydroxide-containing solutions, however, hydrogen peroxide (or HOz-) decays slowly without 0 2 production. This decomposition was detected by KI-induced decomposition and 0 2 analysis, as a function of age of the solution. For example, the quantity of 0 2 evolved from a lOmM H202 solution containing 5mM BudNOH by the addition of KI decreases by approximately 40% after one hour, as compared to 0 2 evolved from a fresh solution. Inasmuch as the non-02-producing decay of hydrogen peroxide in basic acetonitrile might constitute a reaction pathway following reaction 2, a simulated electrolysis experiment was coniluct.ed in which hydroxide ion was destroyed and hydrogen peroxide concurrently added a t the same rate as in an actual electrolysis as based on known charge-time data. HClOd and H202 were added a t 5-minUte intervals to a 5mM Bu,iNOH in a stirred electrolysis cell which contained the platinum electrode surface, exactly simulating electrolysis conditions except that no current was allowed to flow. At the end of the simulation, the solution was examined for H 2 0 2 by addition of KI followed by gas Chromatography of t h e headspace for 0 2 . On addition
\ I
> L - -
-- - LO
ne
\A
-
L '0
P
Figure 3. Current-time decay during exhaustive electrolysis of 13mMBu4NOH at +O 8 V on Pt in 0.1MEt4NC104/CH&N
of the KI, the solution darkened, indicating iodine, and a clearly detectable quantity of 0 2 was found in the headspace. Thus, if diffusing hydrogen peroxide were formed in hydroxide oxidation, it would be stable enough to be observed as a final product. The experimental results eliminate reaction 2 as a pathway for hydroxyl radical decay. A possible criticism of the above conclusion is that it does not take into account a possible anodic Faradaic consumption of H2O2 or H02-. We have observed the anodic cycle voltammetry of these species a t Pt in acetonitrile. Their complex, distinctive anodic patterns are absent from the cyclic voltammograms of hydroxide ion in acetonitrile. A final experiment related to the production of 0 2 involved an electrolysis of Bu4NOH a t +2.3 V us. SCE, well into the rising background reaction. The electrolysis current decreased rapidly during a period in which 0.9-1.0 Faraday/mole of hydroxide was passed. No 0 2 production was detected during this portion of the electrolysis. The current then leveled off and decreased quite slowly thereafter. Increasing quantities of 0 2 were found in the cell headspace during this period. We interpret this as 0 2 production by oxidation of water in the acetonitrile, an oxidation which does not proceed rapidly until the solution becomes neutral by exhaustion of the hydroxide ion. Oxygen production in the early stages of the anodic oxidation of' water has been reported in moist acetonitrile by Portis, Roberson, and Mann ( 1 3 ) . An alternative reaction for the decay of hydroxyl radical is attack of solvent or Bu4N+ supporting electrolyte cation. We have unsuccessfully inspected electrolyzed solutions of hydroxide ion for succinonitrile, ethylene, butylene, carbon dioxide, carbon monoxide, and acetamide. Acetamide was detected but in quantities which would result naturally from base hydrolysis. All of the above compounds are conceivable products of various decay schemes ( 1 3 ) which we have considered. Neither did we detect by gas or thin layer chromatography any unidentified signals. Another approach to the detection of hydroxyl radical is use of a trapping substrate as is done in pulse radiolysis. Hydroxyl radicals will abstract protons from aromatic hydrocarbons and aromatic hydrocarbons bearing alkyl substituents to yield neutral hydrocarbon radicals which dimerize: HAr
+
OH.
-
Ar-
+
H,O
(41
2Ar- --+ A r , (5) Successful competition of substrate HAr with solvent, acetonitrile, is assured if the rate constant for reaction 4 is very large. From pulse radiolysis data ( Z ) , rate constants for the reactions of benzene, toluene, and acetonitrile with
ANALYTICAL CHEMISTRY, VOL. 4 7 , NO. 2, FEBRUARY 1975
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hydroxyl radical are 3.3 X 109, 3.0 X lo9, and 2.1 X lo6, respectively. Thus a diffusing OH. should be consumed entirely by benzene or toluene when generated in mixtures of those hydrocarbons with acetonitrile. We have accordingly searched for the formation of the coupling products biphenyl and bibenzyl in exhaustive electrolysis of Bu4NOH in mixed solvents of benzene/acetonitrile (40/60)(v/v) and toluene/acetonitrile (40/60)(v/v). In electrolyses conducted a t +1.1and f1.7 V, no coupling products were observed by thin layer chromatography. In these electrolyses, as well as those preceding (no trapping substrate), we observed a darkening (brown coloration) of the reaction mixture during the electrolysis. While the negative product findings in this investigation are somewhat unsatisfying, a partial interpretation of the anodic hydroxide oxidation in acetonitrile is possible, and follows. We regard anodic current peaks I and I1 and cathodic current peak V as an ensemble of Pt “oxide-forming” and “oxide-reducing” waves. The two anodic waves yield only surface-limited quantities of anodic electrolysis products and so, in a sense, are incidental to anodic current peaks I11 and IV, which appear to constitute the steps in which bulk quantities of hydroxide ion are consumed. We unfortunately do not know the chemical identity of the product(s) of the bulk electrolysis. They appear to be species of higher molecular weight than would be readily detected by our extensive chromatographic examination, and may be polymeric. As for the nature of the electrode reaction which is initiated in waves I11 (and IV?), the coulometric n = 1 result indicates the initial product must be (Pt-OH)” or (OH-)soln.The hydroxyl radical trapping experiment demonstrates that a freely diffusing hydroxyl radical containing the aromatic hydrocarbon in its solvation shell is not generated. The hydroxide ion reactant on the other hand probably does not contain the aromatic hydrocarbon in its solva-
tion shell as it arrives a t the electrode surface. Then, if the decay of the hydroxyl radical oxidation product exceeds the rate of acetonitrile solvation shell relaxation, no trapping of the radical by the aromatic hydrocarbon would be observable. Such a fast decay reaction is unexpected from pulse radiolysis kinetic data for hydroxyl radicals produced in solution ( 2 ) , which in turn suggests that the hydroxyl radical decay reaction is catalyzed by the electrode surface and, in fact, occurs a t the electrode surface. Whether the surface-catalyzed decay of the hydroxyl radicals is occurring by modes of chemical reactivity not available to the solution radical species is an interesting question. Exploring this question requires knowledge of the ultimate electrolysis products, frustratingly unknown.
LITERATURE CITED (1)A. D. Goolsby and D. T. Sawyer, Anal. Chem., 40, 83 (1968). (2)M. Anbar and P. Neta, lnt. J. Appl. Radiat. Isotopes, 18, 493 (1967). (3)M. Fleishmann and F. Goodridge, Discuss. Faraday SOC., 45, 254 (1968). (4)G. Atherton, M. Fleishmann, and F. GoodridQe, Trans. Faraday SOC., 63. 1468 (1967). (5)R. P. Van Duyne and C. N. Reilley, Anal. Chem., 44, 142 (1972). (6)L. Fox, Ph.D. Dissertation, University of North Carolina, Chapel Hill, 1972 . _ (7)W. S. Woodward, T. H. Ridgway. and C. N. Reilley. Anal. Chem. 45, 435 (1973) (8)J. J. Lingane, “Electroanalytical Chemistry,” 2nd ed., Interscience, New York. N.Y.. 1958. (9)D. T.’Sawyer, Personal Communication, 1974. (IO)A. I. Popov and D. H. Geske, J. Amer. Chem. SOC., 80, 1340 (1958). (11) H. Angerstein-Kozlowska, B. E. Conway, and W. B. A. Sharp, J. Electroanal. Cbern. hterfacial Electrochem., 43, 9 (1973). (12)B. E. Conway and S. Gottesfeld, J. Chem. Soc., Faraday Trans. 1, 69, 1090 (19733. (13)L. C. Portis, J. C. Roberson, and C. K. Mann, Anal. Chem., 44, 294 (1972). I
- I
RECEIVEDfor review August 16, 1974. Accepted October 16, 1974. We acknowledge support of this research in part by National Science Foundation Grant GP-38633X and Materials Research Center, U.N.C., under National Science Foundation Grant GH-33632.
Voltammetric Deposition and Stripping of Selenium(1V) at a Rotating Gold-Disk Electrode in 0.1M Perchloric Acid Richard W. Andrews and Dennis C. Johnson’ Department of Chemistry, Iowa State University, Ames, Iowa 500 10
The electrodeposition of Se(lV) at a Au electrode in 0.1 M HC104 is concluded to produce Se in three distinct states of activity, and three anodic stripping peaks are observed for large quantities of deposited Se. Approximately a monolayer is initially deposited which is apparently stabilized by one-dimensional interaction with the Au electrode surface. The formation of a bulk deposit of Se produces a large activity gradient which is the driving force for irreversible diffusional transport of Se into the Au electrode forming a AuSe alloy of unknown stoichiometry. Application of stripping voltammetry for determination of trace Se(lV) in 0 . 1 M HC104 is possible if the total deposit does not exceed the equivalent of one monolayer. The detection limit of the technique is approximately 0.04 ppb Se(lV) in 0.1 M HCIOd. Author to whom correspondence should be addressed. 294
The physiological significance of selenium has been extensively documented; Se is an essential nutrient a t trace levels ( I ) and toxic when ingested in exces:, ( 2 ) .The beneficial role is reported to involve a synergistic relationship with vitamin E ( 3 ) . When Se is present in animal feeds a t concentrations less than 0.1 ppm, deficiency symptoms develop. These include white muscle diseme, stiff lamb disease ( 4 ) , high embryonic mortality in ewes ( 5 ) ,hepatosis dietetica in weaning pigs ( 6 ) , and kwashikor in Guatamalan children ( 7 ) . The consequences of Se deficiency are sufficiently severe that the American Feed Manufacturers Assoication has petitioned the Food and Drug Administration to allow the supplementing of chicken, turkey, and swine feeds to adjust the Se concentration to 0.1-0.2 ppm (8).When Se is present in feeds a t concentrations exceeding 5 ppm, chronic selenosis develops manifesting itself as
ANALYTICAL CHEMISTRY, VOL. 47, NO. 2, FEBRUARY 1975
