J. RYU,T . IS.ANDERSEN, AND 13. EYRING
3278 Iieve that we deal with radical AI. The observation that these radicals are not detectable at low temper% tures, but only a t room temperature, may account for a rather complicated mechanism which must be involved in the radical formation.
Acknowledgments. We are grateful to Dr. B. Radak and Dr. I. Dragani6 for their interest during the progress of the work. We are also indebted to Dr. R. Herak and Mr. B. Prelesnik for checking the unit cell dimensions of our crystal.
The Electrode Reduction Kinetics of Carbon Dioxide in Aqueous Solution
y J. Ryu, T. N. Andersen, and H. Eyring” Department of Chemistry, University of Utah, Salt Lake City, Utah 84112
(Received December 8, 1971)
Publication costs assisted by the National Institutes of Health, the National Science Foundation, and the A r m y Research Ofice-Durham
The electrolytic reduction of COZ in neutral, aqueous solution a t a mercury pool cathode has been studied to establish the mechanism and to obtain the kinetic parameters by the steady-state galvanostatic method. The log current us. potential curves show two Tafel regions of different slope which is indicative that different, L e . , consecutive, steps are rate determining in the two regions. The most plausible steps considered are (I) CO, €LO -k e @ HCOZ.(ads) 011- and (2) HCOz. (ads) e + 13COZ-. Reaction orders with respect to the partial pressure of the COZ and the Tafel slopes are considered as the criteria for the proposed mechanism. The Langmuir and Temkin conditions for the adsorption of the reaction intermediate are considered and compared. The experimental results show that the Temkin condition is the more suitable and hence there is an appreciable change in the heat of adsorption with the surface coverage of the reaction intermediate.
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Introduction The cathodic reduction of carbon dioxide in an aqueous solution has been the subject of investigations by several authors.l--13 Through the above works, the principal reaction product in aqueous solution has been found to be the formate ion. Although some of the previously published papers1°-13 have proposed mechanisms for the reaction, certain salient points have been overlooked and it is necessary to reconsider and clarify some aspects of the kinetics. Therefore we have proposed and formulated a mechanism, as was done in some previous papers,’O-labut have more carefully considered the form of the rate equations in order t o apply reaction order diffewcntials as a criteria of our mechanism. I n particular, the influence of the partial pressure of C02 on the pH and the reversible potential, and thus on the form of the rate equations, has been taken into account. Also the adsorption of the reaction intermediate and it6 influence on the rate equations have been more fully and explicitly investigated than had been previously. Finally, the effect of temperature on the rate has been studied.
Experimental Section The details of the experimental setup can be found elsewhere.'* An 11-type cell was used which consisted of an anode and cathode compartment in a thermoThp .Journal of Phy.?ical Chemistry, Vol. 7 6 , No. 62, 1972
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stated jacket. The cathode, which was separated by a glass frit from the anode, consisted of a pool of polarographic grade mercury of 2.98-cm diameter which was connected by a tube of Hg to a Pt wire lead located (1) M.E. Royer, C. R. Acad. Sci., 70, 73 (1870). (2) A. Cohen and S. Jahn, Ber. Deut. Chem. Gcs., 37, 2836 (1904). (3) R. Ehrenfeld, ibid., 38, 4138 (1905). (4) F. Fisher and 0. Prziza, ibid., 47, 256 (1914). (5) M . Rabinowitsch and A. Maschowetz, 2. Elektrochem., 36, 846 (1930). (6) T. E. Teeter and P. Van Rysselberghe, J . Chem. Phys., 22, 759 (1954). ( 7 ) Proceedings of the Sixth Meeting of the International Committee on Electrochemical Thermodynamics and Kinetics, Butterworths, London, 1955, p 538. (8) T. E. Teeter, Ph.D. dissertation, University of Oregon, Eugene, Ore., June 1954. (9) M. Hong, Ph.D. dissertation, University of Utah, Salt Lake City, Aug 1969. (IO) W. Paik, T. N. Andersen, and H. Eyring, Electrochim. Acta, 14, 1217 (1969); W. Paik, Ph.D. dissertation, University of Utah, Salt Lake City, Utah, Aug 1968. (11) L. V. Haynes and D. T. Sawyer, Anal. Chem., 39, 332 (1967). (12) J. L. Roberts, Jr., and D . T . Sawyer, J . Electroanal. Chem., 9, I(1965). (13) J. Jordan and P. T. Smith, Proc. Chem. Soc., 246 (1960); P. T . Smith and J. Jordan, “Polarography 1964, Proceedings of the Third International Congress,” G. J. Hills, Ed., Macmillan, London, 1966, pp 407-418. (14) J. Ryu, Ph.D. dissertation, University of Ctah, Salt Lake City, Utah. June 1971.
E:LECTEODE REDUCTION KINETICS OF COz
3279
above the solution level. The anode consisted of a
Pt wire. The solution mas stirred during the experiment. The potential of the cathode was measured with respecb to a saturated calomel electrode (sce) a t room temperature (25 f 0.5”), which was separated from the cathode by a glass Luggin tube and capillary. The electrolytes used were aqueous solutions of sodium bicarbonate and sodium formate prepared from water doubly distilled from a basic permanganate solution. Reagent grade nitrogen gas and commercial COz were mAered together, after purification, to give the desired compositiori. Purification consisted of passing them througlx a wash bottle containing an acidic solution of 0. I M vanadous sulfate and amalgamated zinc granules a presaturator which contained the invesligation. The partial pressure of 6 0 2 was varied from 95.6 to 618 mm a t a tempera. The upper limit of the COZ prest o only COz in the bubbling gas and differed froin atmoi3pheric pressure (of about 645 mm) by the vapor pressure under the experimental conditions. For temperature variation, experiments %‘ere run at 0.2, 10,20, 30, 40, 50, and 60” under a constant 9 < : o 2of 489 k 4 mm. The solution was pre-electrolyzed between a Hg pool A/cm2 for about 24 cathode and Pt anodes at hr under purified KT2gas a t room temperature. After pre-electrol:ysis, the used mercury was replaced for the kinetic measurements. The range of the current density studied was jrom 5 X lo-’ to about loe3A/cm2. The current-potential relations were studied by the steady-state galvanostatic method with stirring. The stable potentials in the Tafel region did not change significantly (zk2 mV) after waiting several hours. The potentials \\we measured with a recorder hooked t o an electrometer amplifier or with a potentiometer, arid the currents ivcre measured with meters accurate t o within 2%. The pbl of .(,hesolution was recorded following each experiment.
Results and Discussion A . Proposed Mechanism. To postulate the reaction steps me consider previous studies on the subject in light of our own results. The reduction of carbon dioxide in the neutral pH region has been found to give the formate ion as the principal product. Thc current efficiency for the formation of formate ions tends to decrease with increasing acidity of the solution because of the competition from hydrogen e v o l ~ t i o nIo. ~ The current efficiency also decreases at very high current densities because of amalgam formation. Wan Rysselberghe and coworkers6l5 demonstrated that I-ICOZ - and C0,2- ions are not reduced but that CO, must be present. Smith and Jordan13 confirmed the above results and also showed that the electroreductivo species is dissolved COz rather than HzC03
Current
Density
Figure 1. The log current-overvoltage polarization curves under various partial pressures of carbon dioxide at 30’ : 1, 95.6 mm; 2, 219 m m ; 3, 347 m m ; 4, 497 m m ; 5 , 615 mm. Current density in amps cm-2; 9 in volts.
(which either exists in solution or is formed by the rapid combination of C02 and HzO prior to the charge transfer). These authorsI3 also obtained polarographic diffusion currents which substantiated the idea that COZ molecules are the primary reactants rather than one of the hydrated C02 species, of which there are several in low concentration under the present conditions. 16-19 The rate of COz reduction is so slow that it can only be studied in the high overvoltage range. A revealing feature of this reaction on Hg in neutral aqueous media is the occurrence of two Tafel regions-one a t lower current densities having a slope of approximately 91 mV and another at high current densities with a slope of approximately 224 mV a t 30” (see Figure 1). These Tafel lines are independent of stirring. The diagnostic tests for the reaction products and the current efficiencies in the present work were carried out in the low current region (region 1))with 0.1 M XaRC03 solution a t room temperature under 625 n m Hg of pure carbon dioxide. The resulting solutions gave positive results for the chromotropic acid testz0and, hence, the reaction product was confirmed to be the formate ion. Titration of the solution with a permanganate solution showed that the current efficiency was xhich is in good agreement with results of previous research.1° The two Tafel lines with their relative respective slopes (15) P. Van Rysselberghe, J . Rmer. Chem. Soc., 68, 2050 (1946). (16) E. L. Quinn and C. L. Jones, “Carbon Dioxide,” Reinhold, New York, N. P., 1936. (17) W. Heinpel and J. Seidel, Ber., 31, 2997 (1898). (18) P. Villard, C. R. Acad. Sci., 119, 368 (1894); P. Villard, Ann. Chim. Phys., 1 1 , 355 (1879). (19) R. M. Barrer and D. J. Ruzicka, Trans. Faradoy SOC.,58, 2239 (1962); R. M. Barrer and D. J. Ruzicka, ibid., 58, 2253 (1962). (20) F. Feigl, “Spot Test in Organic Analysis,” 6th English ed, Eisenia, New York, N. Y., 1960, p 368. The Journal of Physical Chemistry, Vol. ‘76,N o . $9,1978
3280
J. RYU,T. N. ANDERSEN, AND H. EYRING
lead us to believe that thc reaction occurs in two consecutive steps, which W B conclude are 602
+ RzO + e- --+HCOz. (ads) + OHHCrOz. (ads) + e- +HC02-
(I)
(11)
The overall reaction is hence written as
C02 IH,O
+ 2e +NC02- + OH-
(111)
I n region I of the Tafel curves (Figure 1) step I1 is considered to be rate controlling while in region 2 step I i s rate contrding. Several past ~ t u d i e s l ~have ,~~,~~ also proposed an initial step which involves electron transfer. The present! mechanism, however, suggests .0.98 0 20 40 60 80 100 I20 140 160 SEC. that water is involved in the initial step and that an Figure 2. Open circuit potential decay curve at 30” adsorbed formate radical intermediate (HCO,. (ads)) and fcoZ = 497 mm. is formed in this step. I n support of water participating in the first step, it has been observed’l that the chronopotentiograms and potentials for the reduction verse process I while AG’2 and AGi-z are those for of CO, in dimethyl sulfoxide containing various conreaction 11;f represents the partial pressure of the COS centrations of VI ater change with the water content. in the gas phase; 0 is the surface coverage of the adPrevious ~ o r l r ’has ~ provided evidence for the exsorbed reaction intermediate (the formate radical, istence of the reaction intermediate in region 1 from HCO,. (ads)); R’s are the rates for The appropriate regalvanostatic charging curves and from the difference actions for which the specific rate constants are denoted exhibited by these curves in CQz- and K2-saturated by IC” Iyith the proper subscript; q52 is the potential of solutions. Thc~charge necessary t o produce this inthe outer Helmholtz plane, the position a t which the termediate was found to be 32 pC/cm2. From specreactants of step I are assumed to be located. The troscopic data2‘ the area of the adsorbed formate radireaction order with respect to COZ i s denoted by n. cal is ealcuIated to be 11.3 liz.14 Assuming closeIf the presently proposed mechanism is corr(Act, n should packed coverage is present when 0 = 1, one calculates equal one, but we leave it unspecified in the equations that the 32 piC/cm2 corresponds to 0 N 0.23. The t o be determined as a criterion of the mechanism. The open-circuit potential decay curves, which were meareaction orders of the other species (OR-, H20,HC02-, sured in the present work (cf. Figure a), are also conand 8) in the steps 1, - 1,and 2 are assumed t o be equal sistent with this general range of values for 8. I n this to one, in accordance with the proposed mechanism. case if one considers reaction I1 to be the predominant The present study does not specifically serve to determeans of desorbing the intermediate, then one obtains mine such values. W is defined by 8 from the total charge passed by step I1 in changing the potential by the amount shown in Figure 2. The remainder of this paper serves to substantiate K&b(Kw ~ o H -.._______ K~) and quanbify I he proposed mechanism. Equations 0,02K,aowK, 0.02KCK, 0.1 - aoH- are derived for the rai,e of reaction 111 under different, aozM, ~oH-K, conditions for the heat of adsorption of the intermediwhere K , is the ionization constant of water and the ate, and the experimental results are compared with 0.02 and 0.1 appear since our solution connumbers these formulations. tained 0.02 M formate and 0.08 &’ bicarbonate. K,, 23. Langmuir Adsorption of Intermediate. 1. Rate Equations. If we apply absolutc reaction rate theoryz2 Kb, and K , are the equilibrium constants for the followK, to the proposed mechanism and derive the standard ing reactions: COJgas) H,O(I) e &C03 (in soluelectrochemical free energy of activation in the usual Kb KC H+ HC08- and HC02€I HC02Hi. tion) way,23,24 rate equations 1-6 in Chart I are obtained A&, is the inner potential difference between the elec(see ref 24 for the detailed derivations of similar equations), The notation in Chart I is as follows: a’s (21) 6. Heraberg, “Molecular Spectra and Molecular Structure,” are the transfer coefficients, K’S are the transmission D. Van Nostrand, Princeton, N. J., 1966. coefficients and will be taken as equal to unity through(22) 6 . Glnsstone, K. J. Laidler, and H. Eyring, “Theory of Rate Processes,” MoGraw-Hill, New Y o r k , N. W., 1940. out this paper, 7 equals the overvoltage, a’s are the (23) P. Delahay, “Double Layer and Electrode Kinetics,” Wiles, activities; AG *I and AG *-I are the chemical part (nonNew Yorlr, N. Y., 1965. voltage-dependent part) of the standard electrochemi(24) B. E. Conway, “Theory and Principles of Electrode Processes,” cal free energies of activation for the forward and reRonald Press, New York, N. Y . , 1965.
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The Journal of Physical Chemistry, Vol. 76, No. 22, 1972
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ELECTBODE REDUCTION KIIYETICS OF COz
3281
Ch4artI:: The Rate Equations for Reactions I and I1 under Langmuir Adsorption for the Intermediate
_ I
Table I : The Standard Reversible Potential of Reaction 111 Compared to the Standard Hydrogen Electrode a t Several Temperatures Temp, "C Rev. potential (EDh)
2
10
20
25
-0.699
-0.708
-0.718
-0.723
30 -0.728
40
50
-0.739
-0.750
60 -0.760
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trotle and solution at standard state. The other symbols in Chart I have their usual meanings. The overvoltage is defined by 11 = Eh - E h r where h h is the potential of the working electrode and E h r is the reversible or Kernst potential with respect to the standard hydrogen electrode. The reversible potential of the Rg electrode can be written in the form
The reaction is too slow for E h r to be measured directly, but the values of _Fob (eq 8) a t various temperatures were computed from thermodynamic data26 and are tabulated in Table 1. The rates for R-2 are not formulated as there was no indication that HC02- ions influenced the overall rate. 2 . Reaction Orders Using Langmuir Adsorption Condition. a. Reaction I as the Rate-Determining Step. At high overvoltage the net current density is given by i
=
2i0,1exp(-alFq/RT)
r p ~
(9)
where io,l is the exchange current density for reaction I which is given by z0,1 = F1col. From Figure 1 log io,l = -8.81 at! = 497 mm and T = 30"; a1 = 0.27 independent of j a t 30". From eq 9 with the aid of eq 1 and 2 in Chart I we can derive the reaction order differentials in the forms
by assuming that 0
