Electrochemistry of Carbon Dioxide in Dimethyl Sulfoxide at Gold and Mercury Electrodes Louis V. Haynes and Donald T. Sawyer Department of Chemistry, University of California, Riverside, Calif.
Chronopotentiometry, controlled potential coulometry, gas chromatography, and pH titrations have been used to investigate the electrochemical reduction of dissolved carbon dioxide in dimethylsulfoxide at gold and mercury electrodes. Reduction at gold gives well defined, irreversible chronopotentiograms that are diffusion controlled and involve a one-electron process; the transfer coefficient, CY,is 0.30 and the average forcm ward heterogeneous rate constant, kl,ho, is sec-1. Reduction at mercury is complicated by equilibria involving water, but is diffusion controlled with a value for unoof 0.64 and a heterogeneous rate constant of 10-25 cm sec-1. Controlled potential coulometry at mercury indicates an overall one-electron process. The reduction products at both electrodes a r e carbon monoxide and carbonate ion in anhydrous media; some formate and bicarbonate ions are formed as water is introduced into the solvent system. Reduction mechanisms a r e proposed which a r e consistent with the data.
ALTHOUGHCARBON DIOXIDE and its chemistry have been extensively studied, investigations of its reduction have been limited primarily to aqueous systems. Examples of the latter are the electrochemical reduction of carbon dioxide at platinum electrodes ( I ) and at mercury electrodes ( 2 , 3), reduction by sodium amalgam ( 4 ) , and reduction by ultraviolet radiation (5). Polarographic reduction of carbon dioxide in dimethylsulfoxide (6, 7), the reduction at aluminum in fused salts (8), and a voltammetric study in dimethylsulfoxide ( 9 ) represent the previous electrochemical studies of carbon dioxide in nonaqueous media. The present investigation has been undertaken as a part of a general study of the electrochemistry of dissolved gases in aqueous and nonaqueous systems. The specific goal has been to study the electrochemical kinetics and mechanisms for the reduction of carbon dioxide at gold and mercury electrodes in a nonaqueous solvent. Chronopotentiometry, controlled potential coulometry, and the galvanostatic method have been used to determine the kinetic parameters and the overall stoichiometry of the reduction process as well as for the identification of the reduction products.
(1) J. Giner, Electrochim. Acta, 8, 857 (1963). (2) J. Jordon and P. T. Smith, Proc. Cliem. SOC.,1960, p. 246. (3) T. E. Teeter and P. Van Rysselberghe, J. Ckem. Phys., 22, 759 (1954); Proc. 6th Meeting C.I.T.C.E., Butterworths, 1955, pp, 538-42. (4) A. A. Smirnov and V. D. Semchenko, Tr. Novockerk. Polytekhn, Inst., 133, 113 (1962). ( 5 ) N. Getoff, G. Scholes, and J. Weiss, Tetrahedron Letters, No. 18, 17 (1960); Nature, 186, 751 (1960). (6) H. Dehn, V. Gutmann, H. Kirch, and G. Schober, Monntsh. Cliem., 93, 1348 (1962). (7) G. Schober, Abliandl. Deut. Akad. Weiss Berlin, KI. Chem., Geol., Biol., (1) 496-7 (1964). (8) J. Thonstad, J . ElecfroclzemicalSOC.,111, 955 (1964). (9) J. L. Roberts, Jr., and D. T. Sawyer, J. Electroanal. Chem. 9, 1 (1965). 332
ANALYTICAL CHEMISTRY
92502
EXPERIMENTA L
The electrochemical measurements were made with an instrument constructed from Philbrick operational amplifiers following a design by De Ford (IO); circuits of the three electrode type were used to minimize IR errors. Potentialtime curves were recorded with a Sargent Recorder, Model SR-2. For galvanostatic measurements, the potential at zero time was determined by extrapolating the essentially linear portion of the potential-time curve to zero time; this method gives identical intercepts to those obtained by ] pH measurements extrapolating E cs. log [l - ( f / 7 ) 1 i 2plots. were made with a line operated Leeds & Northrup instrument (No. 7401) employing a Leeds & Northrup Miniature pH Electrode Assembly (No. 124138). The cell employed for the chronopotentiometric measurements was a Leeds & Northrup coulometric cell (No. 7961) consisting of a 100-ml tall-form beaker which was connected to a fitted top by a polyethylene snap ring. The fitted top contained holes for the working electrode, the reference electrode, a fritted glass tube for isolation of the auxiliary electrode, and two gas inlet tubes (one for passing gas through the solution; the other for passing gas over the solution surface). The cell solution temperature was controlled with a water bath at 25.0 + 0.1" C. Stirring was accomplished with a water-submersible, air-driven magnetic stirrer and a Tefloncovered stirring bar. Solution volumes were generally about 50 ml. The cell used for coulometric measurements was fabricated from a 40/35 standard taper borosilicate joint. The inner portion of the joint was sealed offin the form of a beaker and contained a tungsten wire in the bottom to act as a contact when employing a mercury pool working electrode. The top of the cell was made from the outer portion of the joint and contained a fine-porosity fritted glass tube for isolation of the auxiliary electrode, an inlet for insertion of a reference electrode employing a 7/15 standard taper joint, and an inlet covered with a syringe cap for solution introduction and sample extraction. When a gold foil electrode was employed, the gold wire lead was inserted through the syringe cap. Stirring was accomplished by means of a magnetic stirrer and a Teflon-covered stirring bar. The total cell volume was approximately 43 ml. The reference electrode consisted of a 5-inch length of 3-mm i.d. borosilicate glass tubing with a small ball of soft glass sealed in one end to give a cracked tip. The resistance across the cracked tip was approximately 6000 ohms and the leak rate through it was about 0.1 ml per week. A short piece of capillary tubing (3-mm i d . ) was used as a top. A silver wire was threaded through the capillary with a loop at the top to provide a contact, and with a coating of silver chloride formed on the lower end. A piece of shrinkable polyolefin was used to hold the capillary to the borosilicate glass tubing and allowed easy filling of the electrode. An aqueous 1F tetramethylammonium cL!oride solution was used in the reference electrode. The potential of this electrode was f0.240 volt cs. N.H.E., and was 0.000 volt when compared
(10) D. D. DeFord, private communication, presented at the 133rd National Meeting, ACS, San Francisco, Calif. 1958.
to a Leeds & Northrup No. 1199-31 saturated calomel electrode. A Luggin capillary, made from 5-mm i.d. borosilicate tubing, was employed to minimize ZR errors (the uncompensated resistance was estimat1:d to be less then 60 ohms based on a.c. measurements and geometric considerations). The reference electrode was inserted into the Luggin capillary and the working electrode placed within 0.5 mm of the Luggin capillary tip. The gold working ele8:trodes used for chronopotentiometric studies were of two tyr'es; foil and inlay. The foil electrodes were made by fusing the foil to a gold wire and sealing the wire in polyethylene. The gold inlay electrode was made by attaching a gold wire to a 3/8 in. piece of 1/4-in,diameter gold rod. The rod was sealed in polyethylene with the polyethylene trimmed to expose the end surface of the rod. The inlay electrode was then polished with 400 and 600 mesh jewelers rouge until a smooth surface was obtained. Final pretreatment of the gsld foil and inlay electrodes consisted of anodizing each for 15 minutes in a solution consisting of 3F HC1O4,3F HOAc, iind 0.03F HC1 followed by rinsing with distilled water and drying. A mercury electrode for chronopotentiometric studies was made by abrading a platinum inlay electrode (Beckman No. 39273) with Carborundum paper (400 grit) in a mercury pool (11) until a smooth, uniform film of mercury covered the surface. The electrode was stored in mercury when not in use; excess mercury was removed by a gentle tap before inserting in the sample solution. Pretreatment was accomplished by cathodically scanning two or three times prior to recording chronopo,:entiograms. Controlled potential coulometry with a mercury pool was carried out by placing reagent grade mercury in the coulometric cell. The area of this pool was approximately 7 cm2. A large gold foil electrode, approximately 1 by 3 inches, was used for controlled pot8:ntial studies at gold. Pretreatment of this electrode was accomplished by heating the foil to dull redness over a Fischer burner and cooling to room temperature in a desiccator. The electrode areas were measured by employing the Sand Equation ( 1 2 ) with known concentrations of K3Fe(CN)~in 0.5F KC1; a diffusion coeffusion coefficient of 7.67 X 10cm2/sec. was used for the calculations (13). The calculated areas were: gold foil electrode, 0.515 cmS2;gold inlay electrode, 0.225 cm2; arid amalgamated platinum electrode, 0.222 cmz. A 0.1F solution of tetraethylammonium perchlorate was used as the supporting electrolyte. The salt, obtained from Eastman Organic Ch1:mical Co., was recrystallized from doubly distilled water, dried and stored in a vacuum desiccator over calcium chlc'ride. Dimethyl sulfoxide (J. T. Baker Co., reagent grade) was obtained in 250-gram bottles with a lot analysis of less than 0.001 water. Because the solvent gave clean cathodic and anodic backgrounds, no further purification was attempted. Mixtures of carbon dioxide and nitrogen containing 1.01%, 3.1%, l O . l % , 29.9%, and 100% CO? (volume per cent) were obtained from the Matheson Co. Carbon dioxide solutions were prepared by saturating the solvent system with the appropriate carbon dioxide-nitrogen mixture. The solvent system was satural.ed until a consistency in the transition time for the chronopotentiogram of the solution was obtained. Chronopotentiometric data for tetraethylammonium formate, carbonate, and bicarbonate were obtained using 5 x lO-3F solutions in dimethyl sulfoxide containing 0.W tetra-
(11) S. A. Moros, ANAL.CHEM., 35, 1088 (1963). (12) P. Delahay, "New Instrumental Methods in Electrochemistry," Interscience, New York, 1954, pp. 184, 187. (13) M. von Stackelberg, M. Pilgram, and V. Toorne, Z . Elekrrochem., 57, 342 (1953).
1
5 sec.
100% (3.0mAJ ,
I
l
0
0
0
l
,
0 0 Time, s e c .
Figure 1. Chronopotentiogramsfor reduction of carbon dioxide in dimethyl sulfoxide at mercury electrode
Per cent, by v01, of COZin N P to saturate solution noted on curves
ethylammonium perchlorate. Tetraethylammonium formate was prepared by adding equivalent amounts of 10 % aqueous tetraethylammonium hydroxide and formic acid, evaporating the water by vacuum evaporation, and drying in a vacuum desiccator over calcium chloride. Tetraethylammonium carbonate and bicarbonate were prepared by bubbling 100% C02 through 10 aqueous tetraethylammonium hydroxide until the appropriate pH was obtained. The solutions were concentrated by vacuum evaporation over a steam bath and cooled. Whitish crystalline solids were collected and stored in a vacuum desiccator. A saturated solution of ammonium oxalate in dimethyl sulfoxide (containing 0.1F tetraethylammonium perchlorate) was used to obtain the chronopotentiometric data for oxalate ion oxidation. The solubility of COz was determined by adding an aliquot of the carbon dioxide-nitrogen saturated solution to 20 ml of redistilled water containing approximately 0.8 equivalent of carbonate-free sodium hydroxide in an atmosphere of prepurified nitrogen. The resulting solution was titrated, using a miniature pH electrode assembly, with carbonatefree sodium hydroxide solution from a Gilmont microburet. The concentration of dissolved carbon dioxide was found to be linear with the volume per cent carbon dioxide in the gas mixture, with concentration equal to 0.816 X I O d 3 times the volume per cent of carbon dioxide in the gas mixture. pH titrations in dimethyl sulfoxide were carried out by adding an aliquot of the solution to be titrated to 30 ml of 0.1F tetraethylammonium perchlorate in dimethyl sulfoxide. Standardized solutions (1.50 x lO-3F) of carbonate, bicarbonate, formate, oxalate, and hydroxide ions were used to prepare reference titration curves; the total volume of water added was adjusted to 1.0 ml before titrating. (The addition of small amounts of water had no effect on the titration curves but improved the electrode response markedly.) The solutions were titrated with perchloric acid, prepared by diluting 1.0 ml of 0.4F aqueous HC10, acid to 250 ml with dimethyl sulfoxide. A Gilmont microburet and a miniature pH electrode assembly were employed in the titration; an atmosphere of prepurified nitrogen was maintained over the titration solution. A molecular sieve column (Type 5A), maintained at 25" C , was used for the gas chromatographic detection of carbon monoxide. The carrier gas was helium and the injection volume was 1 to 5 pl. An alumina column was used for the determination of ethylene and a 2 Carbowax-on-silica-gel column was used for determining dimethyl sulfide. Both columns were maintained at 40" C, the carrier gas was nitrogen, and the injection volume was 30 pl. Injection of a lO-3F VOL. 39, NO. 3, MARCH 1967
333
,
0.0
1
Table I. Chronopotentiometric Data for Reduction of COz at Mercury and Gold Electrode areas: Hg, 0.222 cm2; Au, 0.225 cm2 -
Hg
COz Concn, molesiliter 2.39 X 8.05 X
2.53 x 10-3 8.31 x 10-4
i ~ l l / ~ / A C , i72liZ/AC, Ampcm Ampcm sec 1:2/mole sec’‘z/mole 528 593 525 598 558 595 617 617
iTl’*/AC,
Ampcm sec”*/rnole 219 290 288 283
solution of dimethyl sulfide in dimethyl sulfoxide containing 0.1F tetraethylammonium perchlorate indicated a detectability level of 5 x 10-4F dimethyl sulfide; the same detectability was assumed for ethylene. RESULTS
Typical chronopotentiograms for the reduction of dissolved carbon dioxide in dimethyl sulfoxide at mercury are shown in Figure 1. The appearance of a double wave at intermediate carbon dioxide concentrations is similar to the effect observed when studying the influence of water concentration on the chronopotentiograms. The chronopotentiograms shown in Figure 2 are for a solution saturated with 100 carbon dioxide, to which increments of 0.1F tetraethylammonium perchlorate in redistilled water have been added. The occurrence of a double wave appears to be due to two different solvated carbon dioxide species; carbon dioxide solvated by dimethyl sulfoxide and water. The reduction of carbon dioxide at gold electrodes gives well defined, one-step chronopotentiograms at all carbon dioxide concentrations and at all levels of water concentration between 0 and 30% by weight. The reduction wave for carbon dioxide is obscured by the solvent background at water concentrations above 35 %. This occurs for both electrode materials (see Figure 2) and may have a direct relation to the eutectic point for the dimethyl sulfoxide-water system which occurs at the same water concentration (14). The variation of the reduction potential (extrapolated to zero time) with respect to water concentration, for dimethyl sulfoxide saturated with 100% carbon dioxide, is small for water concentrations below 15 % at mercury electrodes (-2.1 1 volts us. S.C.E. a t 0 % ; -2.03 volts at 15%); however, a significant effect is observed at higher concentrations (the reduction potential decreases to - 1.64 volts at 25 % water). From Figure 2 it is apparent that the double wave coalesces t o a single wave at approximately 15 % water. The variation in reduction potential with water concentration is more pronounced at gold (-1.91 volts us. S.C.E. at 0 % water; -1.62 volts at 10% water) and may be related to the absence of the double wave observed at mercury. When high current densities ( 5 to 20 ma/cm2) are used for the reduction of carbon dioxide at mercury and gold electrodes, extensive visible gas evolution occurs on the electrode surface; the effect is more pronounced at the gold electrode. Small amounts of nitric acid (0.005F) do not appear to influence the carbon dioxide wave, although hydrogen ion
(14) J. M. G. Cowie and P. M. Toporowski, Can. J. Chem., 39,
2240 (1961). 334
ANALYTICAL CHEMISTRY
v
2
s
I
i
-10-
i
W
>
(0
2’--2.0 1 W
‘
-3.0
I
0
5 10 .Time, s e c .
15
I
20
Figure 2. Effect of water on chronopotentiograms for carbon dioxide reduction at mercury electrode in dimethyl sulfoxide Per cent of water, by wt, indicated on curves; solution saturated with pure CO? gas at 1 atm is reduced at a much more positive potential (-0.85 volt US. S.C.E. for a current density of 0.04 ma/cm2) than is carbon dioxide. At higher acid concentrations (0.05F) the formation of hydrogen gas in large quantities at the electrode surface causes the carbon dioxide wave to be poorly defined. Ammonium ion gives similar results to those observed with nitric acid. Values for i71’2 have been determined for four different carbon dioxide concentrations at both mercury and gold electrodes; five current densities have been used a t each concentration to give transition times between 3 and 30 seconds. The results are summarized in Table I with two values of i71’2/AClisted for mercury. These correspond to the transition times 71 and r2 where 71 is measured from zero time to the inflection of the first wave and 7 2 is measured from zero time to the inflection of the second wave. Comparison of the data for the two electrode materials indicates that the number of electrodes involved in the reduction process at mercury is twice the number at gold. The data support the conclusion that‘the reduction process is limited by semi-infinite linear diffusion at both electrode materials. The large overpotential for the reduction process, illustrated by Figure 1, indicates an irreversible electrode reaction. For such systems a potential--time relation (at constant current) has been derived for electrode reactions with mass transfer controlled by linear diffusion (12); at 25” C the relation is
I+-
n FAC k f , h o
E = - l 0.059 og[ ana
0.059 log ana
[-
(5)1:2]
where E is the potential of the working electrode us. the normal hydrogen electrode, a is the transfer coefficient, n, is the number of electrons in the rate determining charge-transfer step, A is the electrode area, k f , h ois the heterogeneous rate constant for the reduction reaction, i is the current, C is the concentration of the electroactive species in the bulk of the solution, F is the faraday, r is the transition time, and t is the time of electrolysis. This expression indicates that the quantity ano may be evaluated from the slopes of plots for
[
(a) EL6s. log 1 -
(f)’”],
(b) El=ous. log i (at constant C ) ,
- -2.4. . rn
2 -23Y
u -2.2 . u,
2 -21-
n ’
0
=. -2 0 .
-I 0
-2 0
00
log [l-(t/ij’+]
Figure 3. Analysis of chronopotentiograms for reduction of carbon dioxide in dimethyl sulfoxide at mercury and at gold electrodes Solution saturated with :I% COZ(by vol in N2); current, 17 pamp (Au) and 35 pamp (Hg); electrode areas, 0.225 cm2 (Au) and 0.222 cm2(Hg)
Figure 4. Galvanostatic studies of reduction of carbon dioxide in dimethyl sulfoxide at mercury electrode for three different concentrations (vol % C o nin N2) Potentials extrapolated to zero time
and (c) E t = cs. log (7(at constant i). Figures 3, 4, and 5 indicate typical plots for these expressions for the chronopotentiometric reduction of carbon dioxide. Table I1 summarizes the values for the kinetic parameters obtained from such plots for a variety of conditions. The average values of ana for mercury and gold are 0.64 and 0.30, respectively. Reverse chronopoterdograms have been recorded at several current densities for both mercury and gold electrodes; three of these are shown in Figure 6 for a gold electrode. Similar results are obtained with a mercury electrode. At low current densities t qe ratio of the forward to the reverse transition times is approximately 3 :1 at both electrode materials. As the current density is increased the ratio becomes larger and in the limiting case no reverse wave is observed. Formation of gas at the electrode surface increases in volume as the ratio of the forward to the reverse transition times increases. Addition of water to the solvent system
reduces the amount of gas formed and increases the reverse transition time. The Eo.zzvalues for the reverse waves at mercury and 0.40 volt, respectively. gold are -0.20 volt (us. S.C.E.) and A dimethyl sulfoxide solution saturated with 100% carbon monoxide does not exhibit an oxidation wave. However, formate ion gives an oxidation wave at mercury and gold electrodes with Ell4 values of -0.18 volt and +0.40 volt, respectively. Oxalate ion also exhibits an oxidation wave at both electrode materials with E114 values of -0.20 volt at mercury and +0.42 volt at gold. These data indicate that formate ion and/or oxalate ion are possible reduction products, but the reverse chronopotentiometric data do not permit a conclusive identification of the reduction product. The values of n for the overall reduction process at gold and mercury have been determined under various experimental conditions by controlled potential coulometry ; the data are
+
Table TI. Kinetic Parameters for Carbon Dioxide Reduction in Dimethyl Sulfoxide at Gold and Mercury Electrodes Method used to [con] x 103, i, Et-0, volts evaluate a ~ , moles/liter PamP an. cs. N.H.E. log k r , h a A. Mercury Et cs. log (1 E , cs. log (1 Et CS. log (1 Et,o cs. log i Et,o cs. log i Et,o cs. log i ElmoUS. log i
(t/T)112) (t/T)I’Z)
(th)’”)
Et-0 CS.
log (COz)
Et-o
log (COP)
US.
El-o us. log (COP)
0.831 2.53 8.05 0.831 2.53 8.05 23.9 8.05 8.05 8.05
35 100 300 30 30 30 30 10 50 100
0.60 0.57 0.53 0.76 0.72 0.74 0.76 0.59 0.57 0.62 0.64
-1.95 -1.96 -1.94 -1.98 -1.93 -1.88 -1.86 -1.84 -1.90 -1.92
-22.8 -21.9 -20.5 -28.5 -26.1 -27.7 -28.4 -22.9 -22.2 -23.7 Av. = -24.5
0.27 0.27 0.27 0.30 0.38 Av. = 0.30
-1.60 -1.62 -1.73 -1.62 -1.58
-10.4 -10.5 -10.9 -12.0 - 14.4 Av. = -11.4
Av. B.
=
Gold
Et,o US. log (1 - ( t / ~ ) ” ’ ) Elmo1;s. log (1 - ( f / T ) l ’ ? Et-0 CS. log (1 - ( t / ~ ) ” * ) Etmous. log i Elmocs. log i
0.831 2.53 8.05 8.05 23.9
17 45 170 30 30
VOL. 39, NO. 3, MARCH 1967
8
335
-2 2
-
-2 2
V
0
W
0
m > ln
2 - -21 > 0
w^
-2 0
I
I
I
log ( % C o p ) , by volume
Figure 5. Galvanostatic studies of reduction of carbon dioxide in dimethyl sulfoxide at mercury electrode as function of percentage composition of CO, in saturating gas at three current densities Potentials extrapolated to zero time
summarized in Table 111. The results indicate that the overall rgduction process consumes one electron per mole of carbon dioxide at a gold electrode and is independent of applied potential. However, at a mercury pool electrode the overall reduction stoichiometry is potential dependent; one electron per carbon dioxide is consumed at the less negative potentials ( - 2 . 3 volts 6s. S.C.E.) and two electrons are consumed per carbon dioxide at more negative potentials ( - 2 . 5 volts). Gas formation is observed not only on the electrode surfaces when using high current densities but also during the controlled potential coulometry experiments. Usually 1 to 3 ml of gas are formed in the coulometric cell during exhaustive reduction of the carbon dioxide solutions. In general, larger volumes of gas are formed when using a gold electrode than when using a mercury pool electrode, but gas evolution is observed for all coulometric experiments (see Table 111). Gas chromatography has established this gaseous product to be carbon monoxide.
Table 111. Coulometric and pH Titration Data for COz Reduction at Gold and Mercury Electrodes Formate Applied formed, potential, Water theo[cod x volts retical, 103, concn., OS. % n molesiliter S.C.E. Wt. % Gold
Mercury
336
-
-2.6 -2.6 -2.3 -2.3 -2.5 -2.5 -2.5 -2,3 -2.3 -2.3
0.005 0.005 0.001 1 .00 0,001 0.001 0.001 1.00 0,001 0.001
ANALYTICAL CHEMISTRY
28
1.08 0.96 0.94 0.90
64 70 60 62 30 40
2.06 1.92 2.11 1.16 0.94 1.14
16 20
trace
2.5 2.5 2.5 2.5 2.5 2.5 2.5 8.0 8.0 2.5
The coulometric results at gold indicate an overall one electron reduction process, but carbon monoxide represents a two electron reduction. One way to account for this is to postulate the formation of an equal amount of carbonate andlor bicarbonate ion with the formation of carbon monoxide. This requires that carbonate ion and bicarbonate ion not be electroactive at the potentials where carbon dioxide is reduced; attempts to reduce these ions in dimethylsulfoxide by chronopotentiometry have been unsuccessful. The coulometric solutions have been titrated with standard acid to give titration curves for the reduction products. For comparison, titration curves for standard solutions of carbonate ion, bicarbonate ion, formate ion, oxalate ion, and hydroxide ion, as well as various combinations of these ions have been prepared. Inspection of the curves indicates that formate and oxalate ions give distinctive curves suitable for their detection; however, carbonate, bicarbonate, and hydroxide ions give curves so similar to each other that qualitative identification of these species is not possible. Analysis of the titration curves for the coulometric solutions establishes that formate ion is a reduction product; no evidence for oxalate ion is observed. The presence of carbonate ion and/or bicarbonate ion, and possibly hydroxide ion, also is apparent. pH titration curves of the solutions resulting from carbon dioxide reduction at mercury indicate the presence of a weak base when the reduction is made at -2.5 volts (n = 2 ) ; this base is not present when the reduction is made at - 2 . 3 volts (n = 1). The total amount of acid required to titrate the solutions reduced at - 2 . 5 volts is twice that required to titrate the solutions reduced at -2.3 volts. The total amount of initial carbon dioxide present, assuming a one electron reduction, is exactly equal to the amount of acid required to titrate the coulometric solution reduced at the less negative potential. The first half of the pH titration curve for the coulometric solutions reduced at - 2 . 5 volts and the pH titration curves for the coulometric solutions reduced at - 2 . 3 volts are similar differing only in the relative amounts of prod-
ucts. Therefore, reduction of the solvent system apparently occurs at the more negative potential, probably by a catalytic process involving the carbon dioxide reduction intermediate. The solvent system csn be reduced either by reduction of dimethyl sulfoxide to dimethyl sulfide or by reduction of tetraethylammonium cation to ethylene and triethylamine. The gas chromatographic data are inconclusive, but the latter process appears more probable. The coulometric data, the gas chromatography data, and the pH titration datz. can be explained for both electrode materials if it is assumed that no hydroxide is formed, that carbon monoxide and carbonate ion are formed in equal amounts, and that forrnate ion and bicarbonate ion are formed in equal amounts by a process involving one molecule of water. By making these assumptions, which agree with the qualitative nature of the titration curves, the total amount of acid required to titrate the coulometric solutions corresponds exactly to the total initial amount of carbon dioxide present in the coulometric cell. Further evidence ir support of these assumptions is obtained by varying the water content in the dimethyl sulfoxide solutions and observii~gthe effect in terms of the products formed on exhaustive ::eduction. Table I11 indicates the per cent of the theoretical amount of formate ion that is actually formed for various water and carbon dioxide concentrations. The theoretical amount of formate ion is based on reduction of one mole of carbon dioxide to one-half mole of formate ion and one-half mole of bicarbonate ion. An increase in the per cent of formate-bicarbonate formed occurs as the ratio of water to carbon dioxide is increased. This result is in agreement with the 0bserva;ions made when studying the effect of water on the reverse chronopotentiograms. The quantitative correspondence between the total initial carbon dioxide concentration and the totai acid used in titrating each coulometric solution has been observed for all systems. Two other interesting effects are apparent from the data in Table 111. First, under identical solution conditions exhaustive reduction of carbon dioxide at gold produces significantly more carbon monoxidl? and carbonate ion than at mercury. Second, the more negative applied potentials produce larger amounts of formate, and bicarbonate ions. CONCLUSIONS
The reduction of cartion dioxide at gold occurs by a straightforward process that gives well defined chronopotentiograms characteristic of an irreversible, diffusion controlled process. The transition times sre reproducible, but somewhat dependent on water cc'ncentration. The overall reduction process occurs by a one-electron mechanism. Product analysis has established that formate ion, bicarbonate ion, carbonate ion, and carbon monoxide are reduction products. Carbon monoxide and carbonate ion are the major products in anhydrous media; the competing products, formate ion and bicarbonate ion, appear when water is present in the solvent system. The value of c q a , 0.30, also implies that the rate determining process is a one-electron step. These conclusions can be summarized in a proposed set of mechanisms which are consistent with the kinetic data: (a) anhydrous system.
:5
.o
c
0.0 w v
v; wl
> m t:
> 0
u- -1.0
-2.0
20
10
0
T i n e , sec.
Figure 6 . Reverse chronopotentiograrns for reduction of carbon dioxide in dimethyl sulfoxide at a gold electrode for three different currents Solution saturated with 30
COn in N2
(b) competing process when water is present, COS
+ e-
-+
. C02-
+ H20 .OH + COzCOz-
+
HCOZ-
+
+ .OH
HCO3-
Two complicating features in the reduction of carbon dioxide at mercury make the interpretation difficult. First, the chronopotentiograms exhibit one or two waves depending on the water-to-carbon dioxide ratio in the solvent. The appearance of the two waves is thought to be due to two different solvated carbon dioxide species, the equilibrium involved being dependent on water content. Second, the coulometric results are dependent on applied potential, giving a one-electron over-all process at - 2.3 volts and a two-electron overall process at - 2 . 5 volts. The two-electron value appears to result from catalytic solvent reduction by an intermediate species at the more negative applied potential, The reduction products formed at mercury are the same as those formed at gold; however, formate-bicarbonate formation is more favorable at mercury under similar solution conditions. Although the mechanisms for the reduction of carbon dioxide at mercury may be the same as at gold, the chronopotentiometric data and the ana values imply that a two-electron rate-determining step is possible. Mechanisms consistent with the data are: (a) anhydrous system.
+ 2e-
COS
-
cos-' 3- co,
ConF2
co + cos-'
VOL. 39, NO. 3, MARCH 1967
337
(b) competing process when water is present (favored under most conditions). COz -/- 2e-
-
+ H 2 02!% HC02- + OHCOn + OH- -+ HC03-
COz-2
The reactions involving the COz-2 species must be slow to account for the two-electron stoichiometry observed with chronopotentiometry. This appears reasonable, especially if the competing process involving water is predominate.
Additional work on the electrochemical oxidation of carbon monoxide, formate ion, and oxalate ion in dimethyl sulfoxide is currently in progress. When completed this should provide additional insight concerning the intermediate species for the reduction process. RECEIVED for review September 12, 1966. Accepted January 7, 1967. Work supported by the National Science Foundation under Grant No. G P 4303. Division of Analytical Chemistry, 150th National Meeting, ACS, Atlantic City, N. J., September 1965.
Rapid Condensation Procedure for Determination of Hydroxyl in Silicone Materials R. C , Smith and G . E . Kellum Dow Corning Corp., Midland, Mich. 48640 A new silanol condensation procedure was developed which allowed rapid and reproducible analysis of total hydroxyl (silanol and water) in a wide variety of silicone materials. A new catalyst system consisting of boron trifluoride, acetic acid, and pyridine was employed. The water originally present and that formed by silanol condensation is removed by azeotropic distillation and titrated with Karl Fischer reagent.
METHODSPROPOSED for silanol determination in silicone materials include manometric LiAIH4 ( I ) and Zerewitenoff procedures (2, 3), infrared (4-6), NMR (3,phenyl isocyanate reaction ( 4 , 5), Karl Fischer reagent titration (8-IO), and silanol condensation (11-16). Condensation procedures have employed alkali, acid, or iodine as catalysts to complete the reaction. The water produced in the condensation is continuously removed and recovered by azeotropic distillation
(1) G. H. Barnes and N. E. Daughenbaugh, ANAL.CHEM.,35,
1308 (1963). (2) F. 0. Guenther, Ibid.,30, 1118 (1958). (3) J. F. Lees and R. T. Lobeck, Analyst, 88, 782 (1963). (4) K. Damm and W. Noll, Ko//oidZ., 158,97 (1958). ( 5 ) W . Noll, K. Damm, and W. Krauss, Farbe Lack, 65, 17 (1959). (6) E. R . Shull, ANAL.CHEM., 32,1627 (1960). ( 7 ) J . F. Hampton, C. W. Lacefield, and J. F. Hyde, Inorg. Chem., 4, 1659 (1965). (8) K . Damm, D. Bolitz, and W. Noll, Angew. Chem., 76, 273 ( 1964). (9) H. Gilman and L. S. Miller, J. Am. Chem. SOC., 73,2367 (1951). (10) W. T. Grubb, Ibid.,76, 3408 (1954). (11) J. Haslam and H. A . Willis, “Identification and Analysis of Plastics,” pp, 256-7, Van Nostrand, New York, 1965. (12) G. M. Kline. “Analytical Chemistry of Polymers, Part I, Analysis of Monomers and Polymeric Materials,” 9. 373, Interscience, New York, 1959. (13) G. R. Lucas and R. W. Martin, J. Am. Chem. SOC.,74, 5225 (1952). (14) L. H. Sommer and G. E. Ansul, Ibid.,77,2482 (1955). (15) L. H. Sommer, R. M. Murch, and F. A. Mitch, Zbid., 76, 1619 (1954). (16) L. H. Sommer and L. J. Tyler, Ibid.,p. 1030. 338
ANALYTICAL CHEMISTRY
with benzene or toluene using Dean and Stark or sirrilar apparatus. The quantity of water collected is then measured either volumetrically or by Karl Fischer reagent titration. Most catalyst systems do not allow rapid condensation of silanol even with monomers or simple structures containing silanol. The condensation is usually incomplete and not reproducible, particularly with silicone resins. A new catalyst system consisting of a mixture of boron trifluoride, acetic acid, and pyridine was employed in our laboratory to give rapid and reproducible total hydroxyl analyses (silanol plus water) in a variety of silicone materials. This method avoids many of the interferences, empirical calibration, and miscellaneous problems as poor solubility, incomplete reaction, and interfering siloxane cleavage associated with many of the above methods. It also has broader application since monomers, fluids, and resins may be analyzed using the same procedure. The apparatus is simple and inexpensive to assemble. EXPERIMENTAL
Apparatus. Barrett type Moisture Test Receivers of 10ml capacity (Ace Glass, Inc., No. 7745) with conventional water condensers were used for azeotropic distillation of large samples. Flasks, receivers, and condensers werc all connected by 24/40 joints. For small samples a miniature distillation unit was employed. It consisted of a Bantamware Distillation Receiver of 2-ml capacity (Kontes No. 28870) with suitable condenser and 50-ml flasks with 14/20 joints. Karl Fischer reagent (KFR) titrations of azeotroped water were performed employing the recording biamperometric apparatus previously described (17). Reagents. Boron trifluoride etherate was Eastman, purified grade. Condensation catalyst was prepared by dissolving 42 ml of boron trifluoride etherate in 208 ml of dry glacial acetic acid.
(17) R. C . Smith and G. E. Kellum, ANAL.CHEM., 38,67 (1966).
