cell wouId in practice probably approach the lowest available value-Le., that t TERUINAL VOILTA determined by composition of the exiting 100 o> 40 c gases. P Figure 1, based on this assumption, f 60: CURRENT DENSITY shows that nothing is gained by trying 5 r 20g to achieve more than about 97y0 utilization. Efficiency is greater if conversion is carried out in stages in separate cells. For example, using Figure 1, a single cell consuming 97y0 of the fuel and operating at voltage A may be compared with a two-cell unit whose first cell a t voltage B consumes and feeds a second cell at voltage A, which takes the consumption up to 97%. In the latter case there is a net gain of energy proporFigure 2. Efficiency o f a fuel battery tional to the shaded area. With 97% having a nominal output o f 2l/2 kw. utilization, a single-stage oxidation will a t 50 volts has been realistically give only 8391, of the available energy, assessed whereas 4 stages give 94yo and 8 stages give 36y0. I n practice, of course, the Battery Efficiencies efficiency is always lower than this because of ohmic polarization. Because a fuel, for example hydrogen, A realistic assessment of the efficiency is progressively consumed when passed of a fuel battery, taking into account all through an electrode, the ratio, PH20/PF12 probable energy losses, has been made increases; hence P',,, the equilibfor a battery with nominal output of rium oxygen partial pressure, also 21/2 kw. a t SO volts. (See Figure 2.) increases and the equivalent cell electromotive force (Equation 4) is expected to Energy efficiency V F = decrease. Because the potential over a electrical energy delivered highly conducting electrode must be to external circuit x 100% constant, the electomotive force of the -AF 0
I-.
3
1.0
-
2,- 0
3.0
4 0
' I
1ia.a
60
U
'Thermal efficiency = electrical energy delivered to external circuit -AH
x
100%
With a reasonable amount of lagging, a battery of this size would be selfsustaining in temperature a t outputs of 1 to 1.5 kw. Based on thickness of the present laboratory cells, the powervolume ratio of the battery a t 2.5 kw. would be 1 kw. per cubic foot of active volume and a t the maximum output of 4.1 kw. this figure would be 1.6. Considerable improvements may be expected to result from reduced cell thickness and internal resistance. literature Cited (1) Bacon, F., British Patent 667,298 (1952). (2) Davtyan, 0. K., acad. xi. U.R.S.S. c l a m sci. tech. 1946, pp. 107, 125. RECEIVED for review September 15, 1959 ACCEPTED January 4, 1960 Division of Gas and Fuel Chemistry, Symposium on Fuel Cells, 136th Meeting, ACS, Atlantic City, N. J., September 13-18,1959. Work was started under the joint sponsorship of the Ministry of Power and the Central Electricity Authority. Additional support was provided by Shell Research Ltd.
H. H. CHAMBERS and A. D. S. TANTRAM Sondes Place Research Institute, Dorking, England
Hydrogen-Oxygen Fuel Cells with Carbon Electrodes C A m o r ELECTRODE fuel cells are characterized by simplicity, reasonable initial cost, low maintenance expense, and an operating life in the range of years. Also, they have high power output per unit weight and volume, a conversion efficiency of about 7070, and a capacity to carry high overloads for short peak demands. I n the cell developed in these laboratories (Figure l), the electrolyte is 3070 potassium hydroxide, and electricity is produced when hydrogen is fed into the inner porous carbon tube. The outer tube is exposed to air, and its larger surface offsets the lower current density of the air electrode. With pure oxygenhydrogen cells, equal-surface electrodes are preferable to obtain proper cell balance; in this instance, tube bundle cells or plate cells were selected. The transportation of oxygen through the wall of the carbon tube determines current of the electrode. Pressure drop between gas side and electrolyte side of the carbon wall, calculated, using Fick's law for linear diffusion, amounts to several per cent of the applied gas pressure, depending on the load. No gas escapes into the electrolyte in a properly operating cell. The pore structure is chosen such that a large pressure differ-
296
ential is required to produce gas bubbles on the electrolyte-carbon interface. Penetration of the electrolyte into the carbon is effectively stopped by a special carbon repellency treatment. Oxygen molecules, adsorbed on the carbon surface, are ionized in accordance with the 2-electron transfer process :
+ H20
On(ads.)
-I- 2e
+
HOn-
+ OH-
Using special peroxide-decomposing catalysts (70, I I), hydrogen peroxide concentration is reduced beyond sensitivity of analytical tests to an estimated 10-'OM. In this connection, it is remarkable that for caustic electrolytes, the minimum half life of peroxide occurs at about pH 14 (4, 5, 75). Different catalysts change the half life by several magnitudes, but the minimum stays in the same pH region. The low concentration of peroxide corresponds to the open circuit potential of 1.10 to 1.13 volts against the hydrogen electrode in the same electrolyte. The oxygen formed by decomposition of the hydrogen peroxide is entirely re-used. This changes the 2-electron process to an apparent 4-electron mechanism. The 0.1-volt difference to the open circuit potential of the oxygen-water electrode
INDUSTRIAL AND ENGINEERING CHEMISTRY
(1.23 volts) reveals that the electrode is not following the equation, 0 2 2H20 4e + 4 0 H - . According to the theory, the oxygen electrode potential must depend on alkali concentration of the electrolyte. Slope of the curve for the oxygenhydrogen peroxide electrode is about 30 to 32 mv. per p H unit, which agrees with the estimated 29 mv. for a 2-electron process (Figure 2). For the oxygencarbon electrode, potential follows the Nernst equation, and as a result, such electrodes can be used for determining oxygen partial pressures. Hydrogen is not active on untreated carbon electrodes and therefore a catalyst was deposited on the surface of the hydrogen electrode. The reaction occurring at the catalytically active sites is
+
+
HZ(gas)
-
2H
ads. On O&talyat
2H (ads.) -I- 2 0 H -
--c
2Hz0
+ 2e
Like the oxygen electrode, structure of the hydrogen electrode is important for the best gas diffusion rate. A permanent three-phase zone (solid-gasliquid) has to be established by wetproofing the carbon material. Also, precautions has to be taken against "internal
FUEL CELLS
RRENT COLLECTOR YDROGEN ELECTRODE
Figure 1 . Concentric hydrogen-air cell. Hydrogen is fed into the inner porous tube
drowning" of the hydrogen electrode by the reaction product water which forms at the anode and creates a second currentlimiting situation. At least this is true for low temperatures. T h e hydrogen electrode also follows the theoretical p H function closely (Figure 2). Good reproducibility of measurements makes the carbon-hydro-
gen electrode a tool for determining activity coefficients. Electrode equilibria are reached in minutes instead of many hours as is required with a platinum-platinum black electrode. I t is not easy to poison these carbonhydrogen electrodes. I n 4 years of testing, no electrode has failed as the result of catalyst poisoning, except where large amounts of cyanide were introduced. Oxygen mixed with the hydrogen is detrimental only if the quantity is sufficient to form large amounts of water catalytically. This recombination prevents dangerous gas mixtures from forming above the electrolyte. This is important when gas is leaked accidently. Under load, the current output increases rapidly with temperature and between 20' and 70' C., at 0.85 volts, it doubles. T h e pressure sensitivity on open circuit follows the Nernst equation. Under heavy load conditions, pressure effect is magnified because of the faster gas diffusion and higher adsorption values reached under pressure. I n principle there are four ways to dispose of reaction water: 1. Operation near 100" C., or when above this temperature, under higher pressure. 2. Operation at low temperatures under reduced pressure; current densities even a t 100 mm. of mercury are above 20 ma. per sq. cm. a t 0.8 volt.
Figure 2. pH function of the oxygen and hydrogen electrode
0.1
, I
1 5
2 3 4,5 , IO ,I5 20 N-KOH 10 20 304050 O/o-KOH
mA/cm2
Figure 3. Voltage-current curves for hydrogen-oxygen carbon fuel cells operating at 60" C. A. Atmospheric pressure, 1958 6. 1 5 0 p.s.i.g., 1958 C. Atmospheric pressure, 1959
3 . Use of gas circulating principle. Water from the electrolyte evaporates through the porous carbon wall especially if a temperature difference is set up. T h e water removal speed depends also on gas flow rates and is limited by the saturation value of water vapor. With a cell temperature of 70' C. and a condenser temperature of 20" C., 180 grams of water are transferred by each cubic meter of gas streaming through the electrodes. Evaporation of water occurs on both electrodes; however, more water appears at the anode if the cell is operating. 4. Operation at low cell temperatures, allowing all the water to enter the electrolyte, with concentration of the electrolyte in a separate thermal or low pressure unit. For low power applications considerable dilution of electrolyte can be tolerated. T h e cell operates as well in 20y0 as in 50% potassium hydroxide--e.g., a 1-amp. cell can be operated for 1000 hours with the production of less than 1 pound of water. Five basic arrangements of electrodes were compared in numerous tests : 2-electrode, parallel tube cell; 4-electrode tube bundle; 9-electrode tube bundle cell; concentric tube cell; and parallel plate cell. T h e increase in current output calculated for the same geometrical electrode area, but arranged in these configurations is remarkable. At a current density of 100 ma. per sq. cm., improvement factors for the arrangements in the order mentioned are 1 to 1.6 to 2.0 to 2.5 to 2.5. For Figure 3, the ohmic resistance is eliminated by means of the pulse current (interrupter) technique. All curves on the graph can be compared on an equal polarization basis. T o calculate actual terminal voltages in special cells, electrolyte resistance of 1.0 to 2 ohm-cm. should be used, depending on temperature and concentration. Electrode spacing should be 0.1 to 0.3 cm. VOL. 52, NO. 4
APRIL 1960
297
For example, voltage drop caused by ohmic resistance in cell components is about 0.02 volt a t 100 ma. per sq. cm. for a parallel plate battery; thus, terminal voltage of the cell can be determined by combining this internal resistance loss with appropiate polarization value from Figure 2. Low temperature, low pressure cells are not subject to electrode attack by electrolyte or oxidation. T h e only lifelimiting factor is wettability of the carbon electrodes, which seems to depend on the potential a t which the electrode operates rather than on current density. Two years' intermittent service has been achieved on 10 ma. per sq. cm. and over 1 year's continuous service on 20 ma. at 0.8 volt. This was at atmospheric pressure, between room temperature and 70" C. In the meantime better repellency treatments and more active catalysts have brought expectations u p to 30 to 50 ma. per sq. cm. over 0.8 volt or at least the same time period. By using high pressures, high currents a t low temperatures can be obtained if more auxiliary equipment is used. Hydrogen is an ideal fuel--'/, pound produces 1 kw.-hr. However for everyd a y purposes, hydrides decomposed by water are more convenient. O n e pound of lithium hydride is equivalent to 1 kw.-hr.
T h e cells developed in these laboratories operate with high current densities on air with only a small potential difference to the pure oxygen-hydrogen cell. With carbonaceous fuels, such as carbon monoxide, alcohols, and aldehydes, good results have been obtained a t low temperatures, but the need for removing carbonate from the alkaline electrolyte complicates these systems. Unfortunately, the present oxygencarbon electrode does not function well in acid. Only current densities of 20 ma. per sq. cm. are obtained a t 0.8 volt. A redox-chemical intermediate such as bromine is necessary for high current outputs. All halogens operate on carbon electrodes with high current densities in acid systems. As a result. hydrogen-chlorine fuel cells can be operated a t high power outputs for extended periods. However, despite higher voltages and current densities, energy output per pound of combined fuel is less than that for the hydrogen-oxygen cell.
References
(4) Hunger, H., Dissertation, Univ. of Vienna, 1954. (5) Hunger, H., Marko, A,, 5th World Power Conf., No. 275 (K-ll), Vienna, 1956.
(6) Justi, E., others, Jahrbuch Akad. Wiss. Mainz (1955). (7) Ibid.,No. 1, (1956). (8) Kordesch, K., Marko, A., Oesterr. Chemiker-Ztg. 52, 125 (1951). (9) Kordesch, K., Martinola, F., Monatsh. Chem. 84, 1, 39 (1953). (10) Marko, A , , Kordesch, K., U. S. Patent 2,615,932 (Oct. 28,1952). (11) Ibid.,2,669,598 (Feb. 16, 1954). (12) Proc. 12th Annual Battery Research and Develon. Conf.. U. S. Armv, Signal . , Research &'Devel. Lab., 1958. 3) Proc. 13th Annual Power Sources Conf. U. S. Army Signal Research and Development Laboratory, Ft. Monmouth, New Jersey, 1959. 4) Spengler, H., Angenw. Chem. 68, 689 (1956). 5) Witherspoon, R. R., Urbach, H. B., Ycager, E., Hovorka, F., Tech. Rept. 4, Western Reserve University, Office of Naval Research Contract Nom 581, (1954). RECEIVED for review July 29, 1959 ACCEPTED December 29, 1959 Division of Gas and Fuel Chemistry, Symposium on Fuel Cells, 136th Meeting, ACS, Atlantic City, N. J., September 13-18, 1959.
(1) Baur, E., Tobler, J., 2. Electrochem.
39, 148-80 (1933). (2) Davtyan, 0.K.. Bull. acad. sci. U.R.S.S. C l a m sci. tech. 1946, p. 107. (3) Ibid.,1 9 4 6 , ~215. .
KARL KORDESCH Union Carbide Consumer Products Co., Parma, Ohio
Catalysis of Fuel Cell Electrode Reactions R E S m R C H ox FUEL cells over the past few years has resulted in the development of commercial prototypes of fuel gas cells operating on such gases as hydrogen, carbon monoxide, and hydrocarbons. Depending on projected applications and power requirements, fuel gas cells have been designed to operate at low and medium temperatures using aqueous electrolytes ( 7 , 2, 6 ) and at higher temperatures using molten salt electrolytes (3, 7). Fuel gas cells, particularly those operating at lower temperatures, are subject to a n irreversible free energy process resulting from the interaction of the reactant gases with the electrode surfaces (3, 8 ) . T h e reactant gases are chemisorbed by the electrode catalyst or the electrode surface and the reaction established is between the chemisorbed species and the electrolyte. T h e potential developed by the cell therefore depends on activity of the chemisorbed species, which is (8) inversely proportional to the heat of chemisorption a t high surface coverages. Thus, the catalyst surface can play a dual role in fuel cell electrode reactions-if chemical
298
kinetics are rate-controlling, it can enhancereaction rate, and it can influence potential of the cell by minimizing- the free energy loss caused by chemisorption.
Fuel Electrode T h e role of the cataIyst a t the anode in a fuel gas cell is twofold: It must rapidly chemisorb the fuel gas in such a manner as to make it more susceptible to oxidation by the active species of the electrolyte and a t the same time it should act to minimize the free energy loss due to chemisorption. Thus, the criteria for a n active catalyst generally will be a weak, but rapid chemisorption of the fuel gas. I n the selection of a catalyst, these conditions must be met in addition to the requirement that the catalyst surface must preferentially chemisorb the fuel gas species over the reaction products so as to limit self poisoning. Hydrogen. Chemisorption of hydrogen, particularly on metal surfaces, has been studied more extensively than other fuel gases. At normal temperatures, chemisorption of the type required for high catalytic activity presumably
INDUSTRIAL AND ENGINEERING CHEMISTRY
involves a partially covalent surface bond between hydrogen atoms and d electrons of the metal. Thus, one requirement for high catalytic activity of a metal in simple gas reactions of hydrogen appears to be that it possesses d band vacancies. This limits the active metal catalysts to the transition elements. T h e early members of the transition series, w-hich have vacancies in both the first and second subbands, chemisorb hydrogen strongly and are in general not as active catalysts in hydrogen reactions as the latter members of the three transition series, which have vacancies only in the second subband. These metals exhibit the lowest heats of chemisorption a t the surface coverages involved in heterogeneous reactions, and are recognized as highly active catalysts for hydrogen reactions. Thus, it appears that the most active metal catalysts for fuel cell electrode reactions where hydrogen is the fuel gas should be selected from those transition metals with d-band vacancies only in the second subband-e.g., group VI11 metals. These views are confirmed by data (Figure 1) obtained with a low-tempera-
