C. E. JOHNSON, R. R. HEINRICH, AND C. E. CROUTHAMEL
242
Thermodynamic Properties of Lithium Hydride by an Electromotive Force Method'
by Carl E. Johnson, Robert R. Heinrich, and Carl E. Crouthamel ChemkaJ Engineering Division, Argonne National Laboratory, Argonne, Illinois (Received August IS, 1966)
The standard free energy of formation of solid lithium hydride has been determined from e.m.f. measurements over the temperature range 675 to 885°K. The cell utilizes an Armco iron flag over which hydrogen is passed as the cathode and a molten lithium anode. The e.m.f. of the cell over this temperature range was expressed as E" = 0.9081 - 7.699 X 10-4T. Derived values of the standard free energy, enthalpy, and entropy at 800°K. are AGt" = -6740 cal. mole-', AHto = -20,940 cal. mole-', and A&" = -17.7 cal. deg.-l mole-'. The values at 298°K. calculated with the use of the best available heat capacity data are AGtO(298) = -16,160 k 50 cal. mole-', AHf"(298) = -21,790 f 290 cal. mole-', and Aij'f"(298) = -18.9 f 0.4 cal. deg.-l mole-'.
Introduction Although the anionic character of the hydride ion has been established,2 and the electrode reaction 1/2Hz e- -+ H- has been utilized in a regenerative cell system,a no systematic thermodynamic data are available on the lithium hydride cell. The purpose of this investigation was to determine the standard values of the thermodynamic functions (AGt", AHf", As,") for the formation of solid lithium hydride over as wide a temperature range as possible by an e.m.f. method.
+
Experimental Section All experimental work was carried out in an inert helium atmosphere of very high purity. The maintenance of the high-purity atmosphere was accomplished by the development4 of a helium gas purification unit which continuously recirculates purified helium through the box. Impurity levels during noma1 operation in the box were below 5 p.p.m. for both oxygen and nitrogen and below 1p.p.m. for water. Lithium metal, a special grade low in sodium content, was obtained from Foote Mineral Co. of Philadelphia, Pa. Before use, the lithium metal was trimmed free of any surface oxide or nitride coating. Reagent grade lithium chloride was purified by the method of Maricle and H ~ m e . The ~ pure material melted at 606.8'. LiBr-2H20 N.F. material from The J o u d of Phyeical Chembtw
Fisher Scientific Company and LiI-3H20 from MalIinckrodt Chemical Co. were purified by a modification of the procedure described by Laitinen, et aL6 The bulk of the water of hydration was removed from the finely ground crystals by programmed heating in vacuo before treatment with their respective anhydrous halogen acids to remove the Iast traces of moisture from the molten salts. Melting points of these purified salts were found to be 549.8 and 468.8O, respectively. Lithium hydride was prepared by contacting liquid lithium metal at 750' with purified hydrogen at 1 atm. for 12 hr. The material was cooled in hydrogen and transferred under vacuum into the high-purity helium-atmosphere box. Chemical analysis of the material gave a composition greater than 99.8% lithium hydride. The melting point of this material was found to be 686.4'. (1) Work performed under the auspicea of the U. S. Atomic Energy Commission. (2) D. C. Bardwell, J. Am. Chem. Soc., 44, 2499 (1922). (3) J. M. Fuscoe, S. S. Carlton, and D. P. Laverty, Regenerative Fuel Cell Systems Invastigations, WADD Technical Report 60-442, AD-249256, May 1960. (4) M. 8. Foster, C. E. Johnson, and C. E. Crouthamel, Report USAECANL-6652, Dec. 1962. (5) D. L. Maricle and D. N. Hume, J . Electrochem. SOC.,107, 354 (1960). (6) H. A. Laitinen, W. S. Ferguson, and R. A. Osteryoung, ibid., 104, 516 (1957).
THERMODYNAMIC PROPERTIES OF LITHIUM HYDRIDE
FURNACE WELL
INLESS STEEL PPORT TUBES
CHROMEL-ALUME THERMOCOUPLE
B e 0 INSULATORS
STAINLESS STEE BUBBLER TUBE
ARMCO I BUBBLER
-
NLESSSTEEL
TERED METAL IBER SPONGE
ELECTROLYTE
LITHIUM METAL
Figure 1. LiH cell aasembly.
Commercial hydrogen gas was purified before use by passage through a silver-palladium alloy' held at 400'. The beryllium oxide tubing used in construction of the anode and cathode was purchased from National Beryllia Corp., North Bergen, N. J. The metal fiber sponge used for the anode consisted of pressed and sintered 50-p diameter fibers of Type 430 stainless steel with a theoretical density of about 10%. The metal sponge was purchased from the Huyck Equipment Co., Milford, Conn. The essential parts of the e.m.f. cell are shown in Figure 1. E.m.f. measurements were made with the cell contained in a furnace well attached to the floor of the helium-atmosphere box. When in use, the cell was isolated from the box atmosphere by a cover which contained insulated holders for the electrodes. The lithium anode was prepared just prior to each experiment. The anode consisted of about 0.25 g. of liquid lithium metal retained on a sintered metal fiber sponge which was held in a small stainless steel crucible. The interfacial tension of the lithium on the sponge was such that the liquid metal was retained in place below the surface of the electrolyte. The cathode comtruction provided for the passage of purified hydrogen over the active Armco iron surface of a bubbler cap held below the surface of the electrolyte. The hydrogen pressure was maintained at 1 atm. during each experiment. Beryllium oxide insulators were used around the center wire of the anode and the central tube of the cathode (see Figure 1) in order to shield each electrode from its surroundings. Cell temperatures were measured by a chromelalumel thermocouple which had been calibrated against the melting points of N.B.S. pure Zn (m.p. 419.5') and N.B.S. pure A1 (m.p. 660.0'). The accuracy of
243
the absolute value of the temperature was estimated to be h0.3'. Lithium chloride, lithium bromide, and lithium iodide, each saturated with lithium hydride, were used as electrolytes. The use of the three pure salts permitted a range of temperatures from 400 to 600' to be studied. The compositions of the electrolyte^^^^ were calculated from the weights of the pure components and were chosen such that for each cell, at its particular temperature of interest, solid lithium hydride was always present. Prior to each cell run,the electrolyte was melted and mixed to obtain the saturated molten salt. Hydrogen gas was present at all times over the electrolyte in order to keep lithium hydride from dissociating. In operation at 550°, the thermal gradient in the fused salt electrolyte (7.6 cm. deep) was a maximum of 0.5'. This low gradient was achieved by proper shunting of a marshall resistance furnace and by introducing baffles above the cell electrolyte surface. The baffles prevented cooling of the electrolyte surface by convective movement of the hydrogen gas in the furnace well. The cell was assembled with both electrodes and a thermocouple in their respective positions in the cell cover. With the electrolyte molten, the thermocouple and the cathode were lowered to a position of about 1.3 cm. below the surface of the electrolyte (see Figure 1). The anode was held in the top cool zone of the cell assembly until the desired temperature of the electrolyte was reached. The anode was then introduced into the electrolyte and the e.m.f. followed immediately by a recording potentiometer. An instantaneous e.m.f. of greater than 1 v. was observed when the cold anode was introduced into the melt. This thermoelectric effect was of short duration, decaying to a more stable e.m.f. within 60-80 sec. After the cold anode was introduced, about 30 min. was required for the electrolyte to return to its equilibrium temperature. At thermal equilibrium, the observed e.m.f. was very stable. Reproducibility of the e.m.f. was good with a standard deviation of about 2 mv. In all cell experiments, pure lithium hydride solid was present in the electrolyte and was chosen as its standard state. For hydrogen, the usual standard state of the pure, ideal gas at 1 atm. pressure was chosen. The deviations of the gas from ideality under the experi-
(7) Purchased from Englehard Industries, Inc., Newark, N. J. (8) C. E. Johnson, S. E. Wood, and C. E. Crouthamel, IWTQ. Chem., 3, 1487 (1964). (9) C. E. Johnson, 9. E. Wood, and C. E. Crouthamel, J. Chem. Phya., in press.
Volume 70, Number 1
January 1066
244
mental conditions are negligible in comparison to the accuracy of the e.m.f. measurements. Pure liquid lithium metal at 1 atm. pressure was chosen as the standard state of lithium. The actual standard state is lithium metal in equilibrium with the satuated molten salt and solid lithium hydride. Experiments have shown that the solubility df the lithium halide in lithium metal up to 1000’K. is small. However, the solubility of lithium hydride in lithium metal may be significant. I n order to test this, experiments were run to determine the effect, if any, of the fused salt electrolyte on the activity of the pure liquid lithium anode. A cell was operated with successive introductions of lithium anodes. The first anode was allowed to remain in the electrolyte for a period of 24 hr., after which time a second anode (preheated to cell tempem ture) was placed in the electrolyte. The same cell voltage was attained within 1 min. after introduction of this preheated anode. In the next experiment, another preheated anode was substituted for the cathode of the cell. Except for a small thermoelectric effect occurring in the first 30 sec., the two anodes showed no appreciable difference in potential. Within the limits of our experimental accuracy, it appears that if lithium hydride does dissolve, its solubility does not appreciably alter the activity of the lithium anode. With this assumption, there is no difference between the operative standard state and that of the pure liquid metal. Initially, attempts were made to carry out the thermodynamic e.m.f. studies utilizing hydrogen gas diffusion diaphragms. Although thin (0.13-0.25 mm. thick) diaphragms of both pure iron and pure vanadium used as electrodes gave a voltage response to the hydrogen pressure with the expected theoretical Nernst slope, the behavior of the electrode potential was somewhat erratic and gave low voltages by as much as 20 mv. The difEculty was traced to the metalfused salt-hydrogen gas interface of the metal tube leading to the diffusion diaphragm. Thus, this electrode possessed two active sites for the catalytic reduction of hydrogen to hydride ions: the site at the metal-fused salt-hydrogen gas interface and the diffusion diaphragm below the electrolyte surface. With the recognition of this problem, a cathode made of an Armco iron flag (0.25 mm. thick, 3 cma2area) which dipped into the electrolyte was used-the active site now being the metal-fused salt-gas interface. Voltage response of this electrode to hydrogen gas pressure in the furnace well was with the expected theoretical Nernst slope and the e.m.f. was extremely stable. The J o u r d of Physical Chemistry
C. E. JOHNSON, R. R. HEINRICH, AND C. E. CROUTHAMEL
With this ‘(flag” cathode, it was found that the temperature coefficient of the e.m.f. became more negative in passing from a saturated to an unsaturated electrolyte for a given loading of the cell. Such behavior implies an activity of lithium hydride greater than unity in the unsaturated electrolyte. An electrode in which hydrogen was passed over an active Armco iron surface beneath the surface of the electrolyte resolved this problem and this form of cathode was used in all subsequent experiments. The solubility of lithium in the electrolyte saturated with lithium hydride also caused some difficulty. When the bubbling cathode was used, the rate of dissolution of lithium from the anode increased, presumably because of the stirring of the electrolyte, and the measured e.m.f. was lower than had been observed in previous experiments. These irreversible effects may be caused by the polarization of the cathode by lithium metal or possible electronic conduction. Several different experiments were conducted to study and overcome these effects. When the voltage of a cell began to drift downward over an extended period, the electrolyte was scavenged by passing hydrogen through the electrolyte, with the electrodes removed, for several hours at 650-700’. A reproducible e.m.f. was obtained once again after this treatment. Apparently, lithium metal which had dissolved in the electrolyte was converted to lithium hydride during this time. The irreversible effects were more noticeable at higher temperatures, an observation which is in agreement with the concept of an increase in the solubility of both lithium and lithium hydride with increasing temperature. In order to reduce the rate of dissolution into the electrolyte, a thin stainless steel cover plate, through which a I/&n. hole was drilled at the center, was placed over the anode cup. This reduced the problem caused by the lithium metal solubility but, although the e.m.f. was more reproducible, it was lower than the best values obtained with the open anode. This effect increased with rising temperature. To determine this effect, the e.m.f. between an open and a covered lithium electrode immersed in lithium chloride saturated with lithium hydride was measured at different temperatures. These e.m.f. values were reproducible to 2 mv. and are given in Table I. The e.m.f. becomes zero below 510’. Also, the differences in the e.m.f. between these electrodes was zero when only lithium chloride was used as the electrolyte. Thus the effect was caused primarily by the presence of lithium hydride and the presumed increase of the solubility of lithium in such an electrolyte. These effects were not observed when lithium bromide and lithium iodide
THERMODYNAMIC PROPERTIES OF LITHIUM HYDRIDE
245
Table I: E.m.f. us. Temperature for a Closed us. Open Lithium Anode Temp., OK.
E.m.f., mv.
852 86 1 856 857 853 843 83 1 837 815 814 812 810
20 16 15 17 14 14 13 10 7
799
3 3
795
6 5 5
were used, presumably because of the lower temperatures involved. The values of the e.m.f. of the lithium chloride electrolyte cells given in this paper have been obtained with the covered anode and bubbling cathode and have been corrected by values taken from a smoothed graph of data of Table I. Values of the e.m.f. of the lithium bromide and lithium iodide electrolyte cells were obtained with an open anode and a bubbling cathode and are used without any corrections.
Results and Discussion The lithium hydride cell may be represented as Li(I)lLiX(l) saturated with LiH(s)IHz(g) 1 atm., Fe
+
for which the cell reaction is Li(1) ‘/2HZ(g) = LiH(s). The data for cells of differing lithium halidelithium hydride compositions are summarized in Table I1 and a composite curve of e.m.f. US. temperature for these cells is given in Figure 2. A comparison of the individual cell data with the composite set showed no apparent inconsistencies and the data are treated as a whole. However, because of corrections, the uncertainty in the data obtained with the chloride cells is somewhat greater than that obtained with the bromide and iodide cells. Therefore, in order to obtain an equation giving the e.m.f. as a function of tempem ture, the e.m.f. data obtained for the bromide and iodide cells were combined and fit to a linear equation in the temperature as were the data for the chloride cells. For the temperature range considered, a linear function seemed to fit the data best. The intercept and slope of the final equation were obtained by using an average of the values in the two equations, weighted according to the number of experimental points in
1.
\
0.35oc
f Y
030/1
O
25850
\-
I
I
,
, 700
I
I
I
,
,,
l\o,o,
750 800 T E M P E R A T U R E , *K
I
,
850
Figure 2. LiH cell e.m.f. us. temperature: V, open Li anode: 40.2 mole % La-59.8 mole % LiI; 0, open Li anode: 40.2 mole % LiH-59.8 mole % LiBr; A, closed Li anode (corrected): 44.9 mole % LiH-55.1 mole % LiC1; 0,closed Li anode (corrected): 60.0 mole % LiH40.0 mole % LiCl. Hydrogen pressure, 760 mm.
each set of data. This equation, which is represented by the solid line in Figure 2, is E o = 0.9081 - 7.699 X lo-%!‘ for the temperature range 673-853OK. with a standard deviation of 1 2 mv. Calculations were made of A G y O , AH!’,and ASio at 800OK. and these data are given in Table 111. Since the literature data for the thermodynamic functions A G f O , myo,and A&’ of lithium hydride are for 298’K., it was necessary to transform our data before comparisons could be made. This was done by integrating the best available heat capacity data over the temperature interval 800-298’K. Heat capacity data for lithium metal and hydrogen gas were taken from Kelley.lo Vogt’l has tabulated similar data for lithium hydride. The values of the heat capacity in cal. deg. -l mole-’ used in the calculations are
+ 11.33 X 10-aT Li Cp(l) = 1.64 + 11.10 X 10-aT + 0.84 X Li Cp(s) = 6.78 + 0.99 X lO6T-2 H~ c,(g) = 6.52 + 0.78 x + 0.12 x 1 0 5 ~ 4 LiH Cp(s) = 7.263
io-3~
The heat of fusion data for pure lithium metal was obtained from JANAFI2 tables. The experimental values and calculated values of A G t O , A H f O , and A S f o are given in Table 111. (10) E. E. Eelley, “High TemperatureHeat Content, Heat Capacity and Entropy Data for the Elements and Inorganic Compounds,” U. 8. Bureau of Mines Bulletin 584, U. S. Government Printing Office, Washington, D. C., 1960. (11) J. W. Vogt, TAPCO Report NP-11888, Oct. 1961. (12) “JANAF Thermochemical Tables,” D o w Chemical Co., Midland, Mich., 1962.
Volume 70, Number 1 January 1966
C. E. JOHNSON, R. R. HEINRICH, AND C. E. CROUTHAMEL
246
Table 11: E.m.f.-Temperature Data for Lithium Hydride Cell 69.8 mole 40.2 mole E.m.f.,
69.8 mole % LiBr 40.2 mole % LiH Temp., E.m.f., K. V.
70Li 70LiH
V.
0.3934 0.3915 0.3864 0.3822 0.3797 0.3739 0.3698 0.3628 0.3598 0.3522 0.3520 0.3440 0.3387 0.3325 0.3318 0.3228
Temp., OK.
0.3440 0.3412 0.3368 0.3324 0.3275 0.3253 0.3247 0.3197 0.3183 0.3183 0.3178 0.3120
673.2 674.2 683.2 687.2 693.2 698.2 704.2 708.2 716.2 722.2 723.7 734.2 742.7 747.2 756.2 766.2
733.7 738.7 744.2 749.7 754.2 758.2 762.2 764.2 767.7 769.7 773.2 778.2
Table 111 : Calculated Values" and Literature Values for the Enthalpy, Free Energy, and Entropy of Lithium Hydride This work, 800OK.
This work, 298'K.
AHto, cal. mole-'
-20,940
-21,790 & 290
AGrO, cal. mole-'
-6,740
-16,160 f 50
Asto, cal. deg.-l
-17.7
-18.9 1 0 . 4
Lit., 298'K.
-21,340b -21, 67OC -16,45ObSd 16,800C'd -16.3d
-
mole-'
-
A H r O = nF[T(a&O/aT), E O ] , AGtO nFE0, A&" = nF(aEO/aT),, where E O ~ ~ = O ~0.2922 . v., (bEo/dT), = -7.699 X 10-4 T v. deg.-'. C. E. Measer, L. G. Fasolino, and C. E. Thalmayer, J. Am. Chem. SOC.,77, 4524 (1955). S. R. Gunn and L. G. Green, ibid., 80, 4782 (1958). K. K. Kelley and E. G. King, U. S. Bureau of Mines Bulletin 592, U. S. Government Printing Office, Washington, D. C., 1961.
The agreement of the values reported here with the literature values is judged to be good. The literature value of the standard entropy of formation of lithium hydride is based on the absolute entropy calculated
The Joumd of Phyadcol Chemialry
66.1 mole % LiCl 44.9 mole % LiH E.m.f., Temp.,
40.0mole % LiCl 6 0 . 0 mole Yo LiH E.m.f., Temp.,
V.
OK.
V.
OK.
0.3076 0.3033 0.2993 0.2951 0.2908 0.2870 0.2828 0.2828
777.2 782.2 788.2 792.2 798.2 803.2 808.7 813.2
0.2845 0.2822 0.2764 0.2747 0.2690 0.2690 0.2605 0.2629 0.2520 0.2517 0.2510
803.2 810.2 813.2 819.7 823.2 828.2 838.2 843.2 844.2 847.7 852.2
from heat capacity data from 74 to 93'K. and a single point at 293'K.Ia The available enthalpy data were derived from 'calorimetric measurements and appear to be quite reliable. In transforming the standard free energy of formation from 800 to 298'K., the accuracy of the result is highly dependent upon the reliability of the heat c ~ r pacity data. This is quite apparent when comparing the standard free energy of formation at 800'K. with that at 298'K. since the value changes by more than a factor of 2. Similarly, the enthalpy of lithium hydride at 298'K. calculated with the use of the e.m.f. data reported here and the JANAF heat capacity data is some 2 kcal. lower than the literature data. In spite of these problems, comparison with literature data is quite good.
Acknowledgment. Helpful discussions with Professor Scott E. Wood of Illinois Institute of Technology are gratefully acknowledged. (13) P.Guenther, Ann. Physik, 63,476 (1920).
