Anode-Electrolyte Temperature Differentials ELECTROLYTIC FORMATION OF OZONE T. R. BECK
AND G . L. PUTNAM University of Washington, Seattle, Wash.
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Because electrode reactions take place at the electrode surface, bulk electrolyte conditions such as temperature are important only in so far as they may affect the electrode surface conditions. Failure to take into account the actual conditions at the electrode surface is very probably an important factor in the well known difficulty in scaling up many electrochemical processes. For high overvoltage electrodes, the results indicate surprisingly large temperature differentials between the electrode and the electrolyte, when operated at commercial
current densities. Correlations of the heat transfer coefficient, and of the temperature differential with electrolyte viscosity, gave linear relationships for the system studied. When the electrode reaction is strongly temperaturesensitive-e.g., reactions of the type including electrolytic production of ozone and persulfates, and the plating of chromium and nickel at high current densities-the electrode-electrolyte temperature differential becomes important in controlling the course of the reaction.
T
the passage of n faradays is the sum of the reversible and the irreversible heat terms:
HE conditions for the industrially important electrolytic processes are nearly always specified in terms of bulk electrolyte conditions-e.g., bulk electrolyte temperatures and compositions and over-all cell potentials. Such conditions are important only in so far as they affect conditions at the site of the reaction, the electrode surfaces. Failure to take into account the actual conditions a t the electrode surface is very probably an important factor in the difficulty in scaling up many laboratory processes from beaker sized to plant sized equipment. Many electrode reactions of actual or potential industrial importance-for example, persulfate formation (II), chromium plating (6),and electrolytic ozone (g)-are strongly temperature-sensitive. In this paper it is shown that the temperature of an operating anode can be appreciably different from the bulk electrolyte temperature. It, has usually been assumed that the temperature of an operating electrode is essentially that of the electrolyte in which it is immersed. In general, however, heat is liberated or absorbed a t an operating electrode and the transfer of this heat through the electrolyte film adjacent to the electrode surface causes a temperature difference to develop. The heat URRENT effect a t the surface may arise in t w o w a y s : (1) reversible heat effect due to difference in heat of reaction and the free energy of reaction, and (2) irreversible heat effect due to overvoltage. In addition, there may be heat developed in the electrolyte film and the bulk electrolyte due to I R drop. The net heat that is evolved as a result of a reversible electrochemical change is: AQ = -nFE
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AH =
-nFT
(E)
The total heat liberated a t the surface of an actual electrode by
AQ' = - n F T
(g)p+
nqF
This equation applies to both anodes and cathodes. If ( b E / b T ) , for the electrode reaction is negative, heat is always liberated a t the electrode, but if ( b E / d T ) , is positive, heat may be either liberated or absorbed, depending on the relative values of ( dE/dT), and the irreversible heat. I n cases of the latter type, heat may be absorbed at low current densities and liberated a t high current densities owing to the increase in overvoltage and current. That any electrical energy applied to a cell in excess of the heat of reaction is liberated as heat was shown in early (1913) work by J. W. Richards who measured the electrical energy input in a cell electrolyzing water; the cell was operated in a calorimeter (IO). Measurements of the temperature of operating electrodes reported in the literature have been confined to systems of very low overvoltage and t o low current densities. Bruz (S), for example, LEADS measured the temperature of nickel, THERMOCOUPLE silver, lead, bismuth, cadmium, zinc, and copper electrodes in solutions of their ions in his entropy determinations. Tarasov (16)also measured the temperature difference between a copper electrode and copper sulfate solution a t current densities of 1.0 to 10 amperes per sq. dm. The maximum temperature difference reported was about 1' C. Both investigators THERMOCOUPLE worked with low overvoltage systems a t low current densities, where PLAT"UM the heat effects are essentially reversible. PLATINUM TUBE This paper deals with a high overvoltage platinum anode in 40 weight 70 perchloric acid (6). 3' c/ High current densities common in Figure 1. Cell for Determination of A T industrial practice were used, with 1123
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measured with a thermocouple in the mercury pool within the anode. Electrolyte temperatures were measured with a thermocouple in a glass thermowell 1.5 mm. in outside diameter. The thermoaell was mounted on a carriage which slid along a calibrated scale attached to the Transite cover, permitting measurement of electrolyte temperature as a function of distance from the anode. Becaupe of its intsrest in this laboratory in connection with the formation of ozone (9), the electrolyte used was 40 weigbt % (eutectic) perchloric acid. Anode overvoltages with perchloric acid are higher than with sulfuric acid (2).
TEMP.
ELECTROLYTE TEMP. PROCEDURE
The cell was operated a t current densities of 4.5, 9, 18, 36, 54, 90, 135, an: 180 amperes per sq. dm. a t bath temperatures of 40°, 25", 0 , - Z O O , and -45" C. Water baths were used a t 40' and 25' C., an ice bath at 0" C., a constant temperature refrigerated antifreeze bath a t -20" C., and calcium chloridewater eutectic mixture cooled with dry ice to obtain -45' C. At a given current density, the cell was allowed to operate for a sufficient length of time to reach thermal equilibrium before temperatures were recorded. This resulted in higher electrolyte temperatures for the higher current densities at each bath temperature. DISCUSSION OF RESULTS
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DISTANCE FROM ANODE SURFACE- CM. Figure 2. Temperature in Operating Cell the result that large temperature differences were developed. The heat from the I R drop through the anode film under the conditions employed was insignificant, but may become important with higher resistance electrolytes. APPARATUS
ELECTROLYTE TEMPERATURE GRADIENT.A plot of the electrolyte temperature zjs a function of distance from the anodes is presented in Figure 2. The data are from the run in which the cell was immersed in the ice bath. When the axis of the thermowell 1.5 mm. in diameter was more than 2 mm. from the anode surface, the temperature difference was constant, indicating that the film thickness was about 2 mm. The temperature difference, AT, is defined as the difference between the anode temperature and that of the main body of the electrolyte. EFFECT OF ELECTROLYTE TEMPERATURE AND CURREKT DENSITY O N AT. A plot of A T against electrolyte temperature with current density as a parameter is shown in Figure 3. The AT value increases as expected with current density, owing to greater heat liberation. The I R drop in the anode was considered and found t o be negligible. Calculation of the heat contribution of the I R drop in the anode film indicates that it is negligible under the conditions obtaining. However, if a higher resistance electrolyte were to be used, it might contribute significantly. The AT also increases with lower electrolyte temperatures. This can be correlated with increase in viscosity of the electrolyte.
The cell was a 500-ml. Bermelius beaker with a Transite cover 0.25 inch (6 mm.) thick (Figure I), impregnated with Tygon Clear lacquer. The anode consisted of a tube 3.01 mm. in diameter of composition 30 9 5 % p l a t i n u m a n d 501, iridium, 23.4 mm. long, sealed to a glass tube mounted in the Transite cover axially to the 25 cell. The lower end of the anode tube, sealed with a glass bead, was filled with mercury to make electrical connection. 20 The cathode was a strip of 0.002-inch (0.051-mm.) platinum foil 0.5 inch (12.7 mm.) 0; wide, fastened to the inside 15 periphery of the glass beaker. Q Electrical connection to the cathode was made through a No. 21 gage platinum wire IO spot-welded to the cathode and sealed in a glass tube with a mercury connection. 5 Temperatures were measured with calibrated copper-constant a n t h e r m o c o u p l e s and a Leeds and Northrup portable 0 precision potentiometer Model -50 -40 -30 -20 -10 0 IO 20 30 40 8662. The maximum error of t e m p e r a t u r e det-rminaELECTROLYTE TEMPERATURE- * C . t i o n w a s a b o u t 0.2' C. Figure 3. A T as a Function of Bulk Electrolyte Temperature A n o d e t e m p e r a t u r e s were
c
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of natural convection to evaluate the physical properties of the system a t the so-called film temperature, which is defined as the arithmetic mean temperature between the heat transfer surface and the main body of the fluid. The empirical plot presented in Figure 4 gives a linear relationship of AT to filin'viscosity. It is recognized, however, that anode film composition as well as temperature may be appreciably different from that of the bulk of the electrolyte. I n Figure 4, AT was plotted against the electrode film viscosity, assuming the film composition to be 40 weight % perchloric acid. HEATTRANSFER COEFFICIENT.Using the overvoltage data of Beck and Putnam (2), the heat transfer coefficient was calculated from the equation:
So
20
p = h,AAT
IO
0 0
5
IO
FILM VISCOSITY Figure 4.
- CENTlPOlSES
AT as a Function of Film Viscosity of Electrolyte
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I
The area of the anode and the A T are known, but the evaluation of g involved a number of assumptions. The total of the reversible heat, AG - A H , was assumed to be liberated at the anode surface because of uncertainty as to how it is divided between the anode and cathode. This assumption does not cause large error, because the value of AG - AH is in this case only about 10 t o 15% of the product of overvoltage and current, a8 shown by Beck (1). The heat transfer coefficient, k,, was calculated for each current density at electrolyte temperatures of 40°, 20°, Oo, -20", and -40" C. A fairly good, but probably fortuitous, empirical correlation is obtained by plotting h, against film viscosity, as shown in Figure 5. The oxygen and ozone evolved at the electrode tend to cause increased convection and relatively high values of h,. Further work with other electrodes and electrolytes is in progress to determine whether a modification of the Lorenz correlation (8) for natural convection can be employed for the correlation of heat transfer data with simultaneous gas evolution.
These two curves clearly indicate that A T may become appreciable in the range of current densities used in industrial electrochemistry. This means that in temperature-sensitive electrochemical reactions-for example, production of ozone (9)-it is not enough to specify the temperature of the electrolyte, because NOMENCLATURE the temperature of the electrode surface usually controls the electrode reaction. A = area of anode, square feet On cooling a platinum tube anode by circulating a refrigerant E = reversible cell potential, volts through it, Hornbeck and Lash obtained more than 40 weight % F = faradays constant, 96 500 coulombs per equivalent or 23,070 calories per (volt) (equivalent) ozone, as contrasted with a maximum of about 20% when the AG = free energy of reaction, calories r gram-mole anode was not cooled (6). h, = heat transfer coefficient, B.t.u./Kour) (sq. foot)(' F.) Failure to take into account AT has been a sourcy: of error in AH = heat of reaction, calories per gram-mole overvoltage I = current, amperes - measurements. Knobel and Joy (7), for example, measured the temperature coefficient of overvoltage of hydrogen at cur350 rent densities up to 40 amperes per SYMBOL AMPS./% DM. sq. dm. in the temperature range 0 " to 70" C., and stated that the temperatures were reliable to A0.05" C. At a current density of 40 amperes per sq. dm., Figure 3 shows that AT is about 5" C. from 0" to 40" C. for the oxygen electrode. While the AT for a hydrogen electrode would probably be less because of lower overvoltage, neglect of consideration of electrode heating will cause an error in the measurement of overvoltage temperature coefficients at moderate and high current densities. CORRELATION OF AT WITH VISCOSITY. The shape of the curves of AT versus electrolyte temperature closely resembled the viscosity data of Clark and Putnam ( 4 ) for 40 0 I 2 3 4 5 6 7 8 9 weight 70perchloric acid. This sugF L M VISCOSITY CENTWISES gested plotting A T against viscosity, It is the usual practice in problems Figure 5. Heat Transfer Coefficient as a Function of Film Viscosity of Electrolyte
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n = number of electrons taking part in reaction per molecule of reactant or product 4 = heat flow, B.t.u. per hour A Q = AG AH, reversible heat liberated, calories per grammole AQ' = total heat liberated atoelectrode, calories per gram-mole 2' = absolute temperature, Kelvin Ai!' = teryerature difference between anode and electrolyte,
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7 = overvoltage, volts = subscript indicating constant pressure
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ACKNOWLEDGMENT
The authors are indebted to Frank B. West and R. W. Moulton for several valuable suggestions which have improved the manuscript. L. H. Clark did some of the preliminary work on the problem. William Antonius and John Sundling constructed part of the equipment. LITERATURE CITED
(1) Beck, T. R., "Electrode-Electrolyte Temperature Differences,"
CONCLUSIONS
There is a considerable anode-electrolyte temperature difference in the electrolytic production of ozone from perchloric acid. Other reactions of industrial importance are probably significantly affected by electrode-electrolyte temperature differences. For the eutectic perchloric acid system studies, the electrodeelectrolyte temperature difference waa a linear function of the viscosity of the bulk of the solution a t the temperature of the bulk of the electrolyte. High heat transfer coefficients indicate the considerable effect of agitation by evolved oxygen a t an anode on the rate of heat transfer. Neglecting electrode-electrolyte temperature difference should cause considerable error in the determination of the temperature coefficient of overvoltage a t high current densities.
University of Washington thesis, 1950. (2) Beck, T. R., and Putnam, G. L., Trans. Faraday Soc., in press. (3) Bruz, B., 2. physik. Chem., Abt. A , 145, 283-8 (1930). (4) Clark, L. H., and Putnam, G. L., J . Am. Chem. Soc., 71, 34457 (1949). (5) Fink, C. G., U. S. Patent 1,581,188 (April 20, 1926). ( 6 ) Hornbeck, R. D., Lash, E I., Putnam, G. L., and Boelter, E. D., J . Electrochem. Soc., 98 (April 1951). ( 7 ) Xnobel, M., and Joy, D. B., Zbid., 44, 443 (1923). ( 8 ) McAdams, W. H., "Heat Transmission," 2nd ed., p. 242, New York, hIcGraw-Hill Book Co., 1942. (9) Putnam, G. L., Moulton, R. IF'., Fillmore, W.W., and Clark, L. H., 6.Electrochem. Soc., 93, 211-21 (1948). (10) Richards, J. W., Trans. Faraday SOC.,9, 140 (1913). (11) Solanski, D. N., and Sastry, P. S., J . Indian Chem. SOC., 25, 415 (1948). (12) Tarasov, G. Ya., J . Gen. Giza. (U.S.S.R.), 16, 1753-66 (1946). RECEIVED 3Tay 16, 1950.
Some Physical Properties of Sulfur exaf luoride J
H. 6. MILLER, L. S. VERDELLI, AND J. F. GALL Pennsylvania Salt Manufacturing Co., Wyndmoor, Pa.
T h e developing industrial uses of this chemically inert, high-dielectric material prompted an investigation of those physical characteristics which are inadequately treated in the literature, although of significance in the safe handling of the compound as a liquefied gas. Orthobaric densities vary from 0.14 gram per ml. (gas) and 1.47 gram per ml. (liquid) at 9" C. to 0.727 gram per ml., the critical density, at 45.5" C., the critical temperature, and 2794 cm. of mercury, the critical pressure. Pressure-temperature data for steel cylinders at various filling densities (weight of sulfur hexafluoride charged/capacity weight of water at 20" C.) indicate that pressures of 1000 pounds per square inch gage will result at about 69.7",
63.2', and 55.7" C. with filling densities of 100, 110, and 1209" respectively. Nitrogen was found to be appreciably soluble in liquid sulfur hexafluoride, the Bunsen absorption coefficient being about 2.1 at 26" C., and the effect of up to 0.807 weight % of dissolved nitrogen (0.84c/o total nitrogen) on cylinder pressures was determined. Thus the presence of 0.12, 0.42, 0.62, and 0.84% total nitrogen increased the gage pressure at 26" C. from 356 to 365,383,400, and 420 pounds per square inch, respectively. The data obtained as a result of this investigation should be of help in evaluating this industrially new chemical and promote safety in its handling as a liquefied gas in pressure cylinders.
T
of safe filling for cylinders of liquefied gases. With sulfur hexafluoride, whose critical temperature falls in this range, these data will also yield the critical constants. The sealed-tube method of Lowry and Erickson ( 1 ) for determining orthobaric densities was used in this work.
HE potential industrial use of sulfur hexafluoride as an inert material of high dielectric strength has prompted a rather thorough study of this compound (2). However, data that are of significance in its safe handling as a liquefied gas are inadequately treated in the literature, and, as predicted by Schumb ( 5 ) , the reported critical temperature of 54" C. is in error by 8" or 9'. It was the purpose of this study to correct this error and to supply the missing data, which will help in the evaluation of this industrially new chemical and to promote safety in its handling as a liquefied gas in pressure cylinders. ORTHOBARIC DENSITIES
A knowledge of the variation of vapor and liquid densities with temperatures in the range commonly encountered under normal atmospheric conditions is of importance in determining the limit
PROCEDURE. The cross-sectional areas of three glass capillary tubes approximately 2 mm. in bore and 7 to 15 cm. long were determined by the standard procedure, using known volumes of mercury. Purified sulfur hexafluoride was condensed as a solid in the tubes and sealed from the atmosphere, and the weight of the sulfur hexafluoride was obtained. The bulbs were suspended in a water bath behind a safety barrier, and the volume change of the liquid and vapor was observed with variation in temperature. The observations were made through a cathetometer telescope. With these data the liquid and vapor densities of sulfur hexafluoride under its own vapor pressure were calculated by solving the following simultaneous equations:
