Chronopotentiometry in Acetic Anhydride Oxidation and Reduction of the Solvent WILLIAM
B.
MATHER, Jr., and FRED C. ANSON
Gates and Crellin laboratories of Chemistry, California lnstitute o f Technology, Pasadena, Calif.
,The anode and cathode reactions that accompany the passage of current through solutions of sodium perchlorate in acetic anhydride-acetic acid mixtures have been elucidated. Hydrogen gas and acetate ion are formed at the cathode. Hydrogen and acetylium ions are formed at the anode. The acetylium ion condenses with the acetic anhydride to form basic products that effectively remove hydrogen ion from the solution. The occurrence of the condensation reaction causes coulometric generation of hydrogen ions at platinum anodes to be less than 100% efficient.
A
acidimetric titration of bases in which the supporting electrolyte consists of sodium perchlorate in an acetic anhydride-acetic acid solution has been described (11), and the results of titrations of a number of inorganic weak bases have been reported (10). These experiments showed that perchloric acid could be generated electrolytically a t a mercury anode with 100% titration efficiency. Attempts to match this generation efficiency with a platinum anode in the same supporting electrolyte resulted in maximum achievable current efficiencies of 95%. Since all of the electrode reactions that could reasonably be expected to occur at a platinum electrode in this solvent correspond to 100% current efficiency for hydrogen ion generation, the present study was undertaken to elucidate the anodic and cathodic reactions that occur when current is passed through acetic anhydride-acetic acid solutions. COULOMETRIC
EXPERIMENTAL
Except for sodium perchlorate, all chemicals were commercial products of the highest purity and were used without further purification. All batches of sodium perchlorate tested contained significant impurities. One recrystallization from boiling acetic acid reduced these impurities to a n acceptably low level. The recrystallized sodium perchlorate contained acetic acid of crystallization that could be removed by heating to 150' C. The supporting electrolyte solution was prepared by dissolving 0.1 mole Reagents.
1634
ANALYTICAL CHEMISTRY
of sodium perchlorate in 1 liter of acetic anhydride that contained 0.05 to 0.2 mole of acetic acid per liter. Apparatus. The chronopotentiometric apparatus was conventional (6). The electrolysis current was supplied by a bank of four 45-volt batteries connected in series with a large variable resistance. The voltage between the working electrode and the reference electrode was supplied to the input of a high impedance, unitgain follower amplifier employing plugin analog computer amplifiers (4). The output of the amplifier was supplied to the input of a fast response recorder. The working electrode was a 0.2-sq. cm. piece of platinum foil that was positioned directly under the sinteredglass disk a t the end of a tube containing the auxiliary electrode. The mer-
+
1
I 5 SEC
-11
2 v,
-
Figure 1. Cathodic chronopotentiograms for reduction of acetic acid in supporling electrolyte Current density 20 ma./rq. cm. Glacial acetic acid added to 90 ml. of rupporting electrolyte (initially 0.05M in acetic acid] in. dicated for each chronopotentiogram
cury-mercurow acetate reference electrode (ref.) has been described (11). Before electrolysis the platinum electrode was oxidized a t about +2 volts vs. ref. with the formation of a thin film that was reduced a t about 0 volt us. ref. This film was observed to affect electrode reactions in ways similar to those observed with platinum oxide films in aqueous media (1). I n
particular, oxidations appeared to proceed more reversibly on reduced electrodes. Procedure. Standard chronopotentiometry (6, 9) and chronopotentiometry with current reversal (6, 8, 12) were employed to identify electrode reactions by the characteristic potentials a t which they occurred and to follow relative concentrations by changes in transition times. All solutions were deaerated with nitrogen before study. All experiments were performed a t room teniperature (25" f 2" C.). RESULTS AND DISCUSSION
Cathode Reaction. With a supporting electrolyte containing both acetic anhydride and acetic acid, the acetic acid is reduced first. This can be seen in Figure 1, which contains chronopotentiograms for the reduction of several different concentrations of acetic acid in the supporting electrolyte. As the concentration of acetic acid is increased, the transition time for the first wave increases and a plot of irl'* us. the concentration of acetic acid is linear. Beyond the transition time for the reduction of acetic acid very reducing potentials are attained, especially when very little acetic acid is present. During the reduction of acetic acid, hydrogen gas is formed a t the electrode; after the transition time for the reduction of acetic acid, the electrode potential becomes very cathodic, and a dark material appears on the electrode surface. If the current is interrupted, this dark material rises to the surface of the solution and reacts with the solvent to evolve a gas. This substance is probably sodium metal resulting from the reduction of the sodium perchlorate supporting electrolyte. Ohmic contributions to the unusually cathodic potentials in Figure 1 were estimated by measuring the current density-dependence of the final potential reached by the electrode. The experiments showed that the ohmic contribution a t a current density of 20 ma. per sq. cm. was no larger than 2 volts and was probably closer to 1 volt.
The strong dependence of the final cathodic potentials assumed by the electrode in Figure 1 on the amount of acetic acid present is unexpected, and no explanation for this dependence has been discovered. Nevertheless the extremely cathodic potentials that are reached before acetic anhydride is reduced indicate that this solvent could be used advantageously to perform difficult electroreductions, especially if it is necessary to employ electrodes having low hydrogen overvoltages. To determine the products of the reduction of acetic acid, chronopotentiometric experiments were performed in which the acetic acid in the supporting electrolyte was reduced for times much shorter than the transition time corresponding to the concentration of acetic acid, and then the direction of the current was reversed. The resulting oxidation chronopotentiogram exhibits two waves (Figure 2): the first, a t -0.8 volt us. ref.; the second, a t f1.5 volts us. ref. These waves were identilied by comparing the chronopotentiograms in Figure 2 with those for the oxidation of separate solutions of hydrogen gas and acetate ion. The oxidation of a solution of sodium acetate in the supporting electrolyte gives a one-step chronopotentiogram with a potential pause a t $1.5 volts us. ref. Thus the second wave in the oxidation of the reduction products of acetic acid is due to the oxidation of acetate ion. The oxidation of a saturated solution of hydrogen gas in the supporting electrolyte under chronopotentiometric conditions produces a wave a t 0.0 volt us. ref. if no acetate is present, but the wave shifts to -0.8 volt us. ref. in the presence of acetate ion. Thus acetic acid is reduced to acetate ion and hydrogen gas at a platinum cathode in the supporting electrolyte. Since hydrogen gas has limited solubility in acetic anhydride, some of the hydrogen formed in the reduction escapes from the solution before the current reversal. The remaining hydrogen is oxidized in the presence of a n excess of acetate a t the electrode. Were it not for this loss of hydrogen gas, the only electrode reaction expected for the oxidation of the reduction products of acetic acid would be: 1/2H2
+ CHsC02-
=
CH3C02H
+ e-
(1)
Reaction 1 occurs a t -0.8 volt us. ref. Anode Reaction. Oxidation of the supporting electrolyte a t a platinum electrode is not so simple as the reduction. It was not possible to distinguish between acetic acid and acetic anhydride oxidation; but, as shown below, the final products are the same in both cases so long as the acetic
authors reported that a solution of acetylium perchlorate in acetic anhydride darkens rapidly, eventually depositing a crystalline perchlorate and some tar (21, The condensation of acetylium ion with acetic anhydride to produce these colored products in the generating solution probably occurs in the following manner: The hydrogen and acetylium ions produced a t the anode are in equilibrium with the solvent,
-t/ T I ME
H+
Figure 2. Reverse-current anodic chronopotentiograms following reduction of supporting electrolyte Time before current was reversed given with each chronopotentiogram. Current density 4 ma./rq. cm. throughout
acid concentration is small. Acetic acid and acetic anhydride appear to be oxidized according to the following reactions: CH,CO%H= CHaCOz.
+ H+ + e-
+
(CHxC0)zO = CHICOI. CH&O +
+ e-
(2)
(3)
The acetoxy radicals decompose very rapidly into methyl radicals and carbon dioxide (7). According to these electrode reactions a hydrogen ion or a n acetylium ion is produced for each electron entering the electrode. Since hydrogen ion and acetylium ion are in equilibrium in acetic anhydride-acetic acid solutions, CHJO
+ CHaCOsH = (CH&O)zO
+ H+
H+
+ CHaCOnH = CH,CO,H,+
+ (CH8CO)tO
CH,CO+
-
(CBCO)nOH+ ( 6 )
+ CHSCOzH = (CHaCO)IOH+ ( 7 )
The charged species in this solvent doubtless exist to a large extent in the form of ion pairs, but for the sake of clarity this complication will not be indicated in the reactions. One effect of oxidation of the solvent according to Reactions 2 and 3 is to decrease the acetic acid concentration near the electrode, both because it is oxidized and as a result of Equilibrium 5. From Equilibrium 7 it follows that the acetylium ion concentration increases as the concentration of acetic acid decreases. Acetylium perchlorate is a very strong electrophile and would be expected to condense with the enol form of acetic anhydride as follows: 0 0
A&
CHs 0 C H a e CH-
(4)
Reactions 2 and 3 correspond to 100% current efficiency for the generation of acid. However, when titrations of the hydrogen ions formed in the anodic electrolysis of the supporting electrolyte were performed by coulometrically generating acetate ion, not all of the acid that should have been present according to Reactions 2 and 3 was found. The reason for this loss of acid appears to be a condensation of acetylium ions with the solvent to form basic products. On passage of anodic current through either a platinum or mercury electrode in the supporting electrolyte the solution turns yellow. As the solution stands, the yellow color changes to orange and finally to red. The same behavior is observed when a few drops of 9F perchloric acid are added to acetic anhydride. These color changes are very similar to those described by Burton and Praill (2, 9 ) as arising from a reaction between acetylium perchlorate and acetic anhydride. These
(5)
iH 8 -0-
CHs (8)
+
. H+ . .. 0
0 0
CHaACHJObCHa (9)
I
..H+.. 0
0 0
li, A 6
CH, CHz 0 CHI= I 0
0 0
/Ill
CH* CHzCOCCHI + H + (10) The protonated condensation product, I, is a weaker acid than perchloric acid, but experiments described below indicate that it is a sufficiently strong acid to be titrated in the acetic anhydride solvent. However, further condensations can occur: VOL. 33,
NO. 12, NOVEMBER 1961
1635
'i
,
I
SEC
, T
,
I
TIME
Figure 3. Cathodic chronopotentiograms for reduction of perchloric acid in supporting electrolyte at 4 ma. per sq. cm. Concentration of perchloric acid given with each chronopotentiogram
acetyl chloride did not react with acetic anhydride and was without influence on cathodic chronopotentiograms.) Two waves are observed: the first a t -0.6 volt us. ref. and the second at - 0.8 volt us. ref. The wave at - 0.6 volt us. ref. is the reduction of perchloric acid, formed in the reaction of acetylium ion with the acetic acid in the supporting electrolyte according to Reaction 7 . The potential pause is at -0.6 volt us. ref. instead of -0.5 volt us. ref., as observed in the reduction of perchloric acid in acetic anhydride, because of the lower concentration of perchloric acid in the acetylium perchlorate-acetic anhydride mixture. The second wave, a t -0.8 volt us. ref., is most likely due to the reduction of the protonated form of the first condensation product of acetylium ion with acetic anhydride:
.H . .+. ti
'0
o
Ai!
+
CH&" C H ~o C H ~ e - = I
0
0 0
CH$!!CH~~!O&CHI
The higher condensation products are much weaker acids that cannot be titrated with acetate. Consequently, the formation of products such as I11 has the effect of removing hydrogen ions from the solution. Chronopotentiometric Experiments.
A very convenient method for identifying the products of anodic electrolysis of the supporting electrolyte is to compare cathodic chronopotentiograms obtained by reveising the direction of the current after the supporting electrolyte has been oxidized with cathodic chronopotentiograms of perchloric acid in acetic anhydride. Figure 3 shows cathodic chronopotentiograms of a perchloric acid solution in the supporting electrolyte. Only one wave is observed a t -0.5 volt us. ref.; the wave corresponds to the reduction of perchloric acid as follows: HClOd
+ e-
=
1/2 H,
+ Clod-
(12)
The chronopotentiograms in Figure 3 were taken after all the water in the 9F HClO, had reacted with the acetic anhydride, but before appreciable condensation had occurred. Figure 4 shows the reduction of a solution of acetylium perchlorate in the supporting electrolyte. The acetylium perchlorate was prepared by reacting excess acetyl chloride with silver perchlorate in acetic anhydride and filtering off the silver chloride precipitate. (Separate experiments showed that 1636
ANALYTICAL CHEMISTRY
+ 1/2 HI
(13)
That this second wave appears immediately-Le., as soon as the solution could be prepared for study-with solutions of acetylium perchlorate in acetic anhydride is good evidence for a rapid reaction between acetylium ion
-
'h
I SLC.
I
TIME
Figure 4. Cathodic chronopotentiograms for reduction of acetylium perchlorate a t 4 ma. per sq. cm. Concentration of acetylium perchlorate given with each chronopotentiogram. Experiments performed within 3 minutes after preparation of solution of acety lium perchlorate
and acetic anhydride. Reaction 13 proceeds a t a more cathodic potential than Reaction 12 bccause I is a weaker acid than HClO,. The cathodic chronopotentiograms that result when the supporting electrolyte is oxidized and then the direction of the current reversed are shown in Figure 5, Two waves are obscrved:
I
scc
c _ (
T
TIME
Figure 5. Reverse-current cathodic chronopotentiograms following oxidation of supporting electrolyte. Time before current was reversed given with each chronopotentiogram. Current density 4 ma./sq. cm. throughout
a t -0.5 and -0.8 volt is. ref. The wave a t -0.5 volt us. ref. is the reduction of perchloric acid; the wave at -0.8 volt us. ref. is probably the reduction of the protonated form of the first condensation product (I). This second wave is observed only after oxidation times long enough so that the acetylium ions generated in the electrode process have time to react with acetic anhydride before the current is reversed. It is noteworthy that significant condensation occurs in 40 seconds. The fact that no separate waves corresponding to the reduction of acetylium ion are observed in Figures 4 and 5 is doubtless due to the rapidity of its reaction with the solvent to form either perchloric acid or the protonated condensation products. The absence of any cathodic waves except those due to reduction of the acids to hydrogen gas in the reverse-current chronopotentiograms indicates that the other products of the electrolyte oxidation, CHiCOO . , CHs., COz, and C2He,are not reducible or are too short-lived to be detected. Chronopotentiometric Studies of the Condensation Reaction. Perchloric acid disappears from mixtures of perchloric acid and acetic anhydride as the higher, slightly dissociated condensation products form. Since the transition time for the diffusioncontrolled reduction of perchloric acid measures its concentration, the disappearance of perchloric acid from an acetic anhydride solution can be followed chronopotentiometrically. Figure 6 shows the results of a chronopotentiometric study of the reduction of perchloric acid in acetic anhydride at various times after mixing. Little change is observed in the first few minutes, while the water in the 9F HClOl reacts with the acetic anhydride. The shift in the background reduction potential (reduction of acetic acid) is due to the increased acetic acid concentration. The reduction of the perchloric acid
1‘s. ref. .uter 30 minutes a second wave appears at -0.8 volt vs. ref.; this wave is due to the reduction of the protonated form of the first condensation product. The transition time of this second wave increases a t the expense of the wave for the reduction of perchloric acid, and there is also a decrease in the total transition time. The esplanation of this behavior is probably as follows: Initially the only reducible species is thc perchloric acid which gives rise to the wave a t -0.5 volt. As time passes and the condensation of acetylium ion with acetic anhydride proceeds, a second wave, due to the first condensation product, appears a t -0.8 volt. -4s condensation cont,inues with the formation of higher, less dissociated products, there is a n over-all decrease in .the transition time because these products are not reduced before the background. -4nalogous behavior is observed when acetylium perchlorate rather than perchloric acid is added to the supporting electrolyte. A more rapid decrease in the transition time for thc reduction of perchloric acid is observed because the condensation reactions proceed faster at higher acetylium concentrations. Although changes associated with the occurrence of the condensation reaction could be observed in reversecurrent chronopotentiograms after only 40 seconds of oxidation, several hours irere required for the same amount of change in perchloric acid-acetic anhydride mixtures. This difference in the rates of condensation is probably due to the differences in the concentrations of acetylium ion produced at the electrode and present in homogeneous perchloric acid solutions. The concentration of perchloric acid and acetylium ion in the immediate vicinity of the electrode at the end of a 40-second oxidation period was about 0.4F, while the homogeneous perchloric acid solutions were 0.025F in HCIO, and 0.2F in acetic acid. Since the rate of condensation depends on the hydrogen ion concentration, it is not surprising that it is much faster with electrogenerated acid. The rapid condensation near the electrode is greatly accelerated a t higher acid concentrations. When perchloric acid was added to the supporting electrolyte before carrying out oxidations and the direction of the current reversed, the resulting cathodic chronopotentiogram showed no increase in perchloric acid concentration for osida-
-
occu1’s a t -0.5 volt
ing oxidation of acetic anhydride solutions at a mercury anode (10, 11) is somewhat surprising. Apparently the generation of hydrogen ions a t a mercury anode in coulometric titrations of bases does not result in local excesses of perchloric acid near the electrode. The electrode reaction in this case is:
0 5 SEC
1
O
n!
0 5 v.
2Hg I
TIME
Figure
=
Hgz+
+ 2e-
(14)
The mercurous ions diffuse away from the electrode and react with acetic acid to produce hydrogen ions:
6. Cathodic chronopotentio-
grams showing reduction of 0.015F perchloric acid Time since solution was prepared on each chronopotentiogram. Current density 4 ma./sq. cm. throughout
tion times from 5 to 20 seconds. 4 t these oxidation times only a n increase in the transition time for the reduction of the protonated form of the first condensation product was observed. Acetic acid suppresses the condensation reaction by decreasing the concentration of acetylium ion. When acetic acid is added to the supporting electrolyte before recording reversecurrent cathodic chronopotentiograms, only a single wave corresponding to the reduction of perchloric acid results, even at long oxidation times. In this case the acetylium ion formed in the electrode process reacts with the acetic acid to form perchloric acid rather than with acetic anhydride to form condensation products. The potentials at which reversible chronopotentionietric waves for the reduction of acids to hydrogen gas occur in acetic anhydride solutions reflect the strengths of the acids in this solvent. Although the acids studied are not reduced reversibly, the potential a t which each is reduced a t a given current density probably involves a n approximately constant overvoltage due to irreversibility. Accordingly, the following acids and the potentials (us. reference electrode) at which each was reduced at a current density of 4 ma. per sq. cm. are ranked in the order of decreasing acid strength : HClO4( -0.5 volt) > HBR( -0.9 volt) > HSSO,( - 1.0 volt) > HC1( -1.4 volts) > CH&O,H( - 1.7 volte) Coulometric Titrations. I n the light of material presented in this paper the success achieved in coulometric-acidimetric titrations involv-
The hydrogen ions fornied in Reaction 15 then rapidly react with the base present in the solution. I n solutions containing no base, oxidation of the mercury electrode showed the same qualitative behavior as the oxidation of the supporting electrolyte at a platinum electrode. I n this case the hydrogen ions formed in Reaction 15 are not removed by reaction with base, and the condensation reaction proceeds. ACKNOWLEDGMENT
WBM gratefully acknowledges the support of the National Science Foundation. LITERATURE CITED
(1) Anson, F. C., ANAL.CHEM.33, 934 (1961). (2) Burton, H., Praill, P. F. G., J . C h e m SOC.1950.2034. (3) Ibid., 1953,827. (4) DeFord, D . D., AbstractR, 133rd
Meeting, ACS, San Francisco, Calif.,
1958, p. 16B. (5) Delahay, P..
“New Instrumental Methods in Electrochemistrv.” Cham 8, Interscience, New York, 1654.
(6) Ibid., pp. 377-80. (7) Gould, E. S., “Mechanism and Structure in. Organic Chemistry,” p. 717, Henry Holt, New York, 1959. (8) Icing, R. M., Reilley, C. N., J . EleclroanuE. ChRm. 1 , 434 (1960). (9) Lingane, J. J., “Electroanalvtical Chemistry,” 2nd ed., Chap. 22, fnterscience, New York, 1958. (10) Mather, W. B., Anson, F. C., ANAL. CHEM. 33. 132 (1961). ~~~(11) Mathe;, B., Anson, F. C., Anal. Chzm. Acta 21,468 (1959). (12) Testa, A. C., Reinmuth, W. H., ANAL.CHEM. 32, 1512 (1960). ~
iv.
\ - - - - I
RECEIVEDfor review June 5, 1961. Accepted August 28, 1961. Contribution No. 2707 from the Gates and Crellin Laboratories of Chemistry. Work s u p ported in part by the U. S Army Research Office under Grant DA-ORD-31-124-61-
G91.
VOL 33, NO. 1 2, NOVEMBER 1961
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