Energy & Fuels 1997, 11, 931-935
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Lithium-Water Reaction Chemistry at Elevated Temperature Martin Klanchar* The Applied Research Laboratory, The Pennsylvania State University, P.O. Box 30, State College, Pennsylvania 16804
Brian D. Wintrode The Applied Research Laboratory, The Pennsylvania State University, P.O. Box 30, State College, Pennsylvania 16804
Jonathan A. Phillips The Department of Chemical Engineering, The Pennsylvania State University, 133 Fenske Laboratory, University Park, Pennsylvania 16802-4400 Received March 24, 1997. Revised Manuscript Received May 20, 1997X
Lithium-water reaction experiments were conducted at 616.5 K in a thermoelectric calorimeter in an effort to characterize the reaction in a batch-type environment. The experimental system was configured to measure both heat of reaction and hydrogen generation as small fixed-volume doses of water were repeatedly admitted to a large quantity of lithium. Heat of reaction measurements slightly declined with each successive water dose, while hydrogen generation markedly increased. These heat and hydrogen quantities, along with material and energy balances, were used to quantify gradual changes in reaction chemistry as lithium was depleted. The trends evident with heat and hydrogen generation indicated the formation of lithium hydride and lithium oxide as intermediates. These products, in competition with unreacted lithium, reacted with additional water to yield hydrogen and lithium hydroxide. In conclusion, these experiments provided a quantitative understanding of how reaction chemistry changes when water is continuously added to a lithium bath.
Introduction The lithium-water reaction has been studied primarily for electrochemical-battery applications,1,2 as a surface oxidation reaction,3-5 and to realize the effect of inadvertent lithium-water contact as in nuclear environments6 or in extinguishing metal fires.7 But it also can be regarded as a combustion reaction, since it ultimately yields substantial quantities of hydrogen and heat.8,9 A chemical reactor utilizing lithium and water can offer an efficient method of supplying hydrogen and heat, especially in space-limited applications that require relatively high storage and energy densities. For example, the hydrogen storage density of a lithiumwater reactor is approximately 25% greater than cryoAbstract published in Advance ACS Abstracts, June 15, 1997. (1) Li, Wu; Dahn, J. R.; Wainwright, D. S. Science 1994, 264, 11151117. (2) Hoenigman, J. R.; Keil, R. G. Appl. Surf. Sci. 1984, 18, 207222. (3) Deal, B. E.; Svec, H. J. J. Am. Chem. Soc. 1953, 75, 6173-6175. (4) Irvine, W. R.; Lund, J. A. J. Electrochem. Soc. 1963, 110 (2), 141144. (5) Besson, J.; Pelloux, A. C. R. Acad. Sci., Paris, Ser. C 1967, 265 (15), 816-819. (6) Greene, G. A.; Cho, D. H.; Hyder, M. L.; Allison, D. K.; Ellison, P. G. Nucl. Eng. Des. 1994, 148, 317-326. (7) Rhein, R. A. Lithium Combustion: A Review; NWC TP 7087; Naval Weapons Center: China Lake, CA, 1990. (8) Cook, L. P.; Plante, E. R. Survey of Alternate Stored Chemical Energy Reactions; NBSIR 85-3282; National Bureau of Standards: Gaithersburg, MD, 1985. (9) Markowitz, M. M. J. Chem. Educ. 1963, 10 (13), 633-636. X
S0887-0624(97)00047-9 CCC: $14.00
genic (liquid) hydrogen, assuming reaction water can be conveniently supplied from an external source. Heat of reaction and hydrogen production rates, however, are a precise function of reaction sequence. The heat of reaction when lithium hydroxide (LiOH) and hydrogen are formed as products, for instance, is about 210 kJ/mol water reacted at 1000 K.10 But the heat of reaction is substantially highersabout 490 kJ/mol waterswhen products such as lithium hydride (LiH) and lithium oxide (Li2O) form as intermediates. Hydrogen generation is also greatly reduced in this case. Previous studies have indicated that product formation during lithium-water reactions are functions of temperature, pressure, and reaction environment.11 Experience has shown that the lithium-water reaction can be difficult to control, especially when reasonable quantities of heat and hydrogen are desired. An accurate understanding of the reaction process over a range of operating conditions is required for development of a compact, controllable system. For example, in a batch-type system that initially contains pure lithium, the operating temperature, pressure, and reactor bath composition will change as lithium is depleted and reaction products accumulate. It is essential to (10) Chase, M. W.; et al. JANAF Thermochemical Tables, 3rd ed.; American Chemical Society and American Institute of Physics: New York, 1986. (11) Besson, J.; Muller, W. Compt. Rend. 1958, 247, 1869-1872.
© 1997 American Chemical Society
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Figure 1. Apparatus for conducting lithium-water reaction experiments.
realize what effect these changes have in defining reaction chemistry. This investigation examined the lithium-water reaction at temperatures higher than those of the majority of previous studies and under operating conditions in which a small amount of water was added to an excess of lithium, simulating the start of a batch process. A thermoelectric calorimeter was configured to provide heat of reaction measurements, while accompanying hydrogen production was also detected and quantified. As will be shown, these measurements allowed the exact reaction equation to be determined and also demonstrated how reaction chemistry was altered as products accumulated in the system. Experimental Section Equipment and Materials. Lithium-water reaction experiments were conducted in the apparatus outlined in Figure 1. The two primary components of the system included a Calvet-type thermoelectric calorimeter, which measured the heat of reaction, and a hydrogen gas analyzer, which determined the amount of hydrogen in the calorimeter effluent gas. The lithium-water reaction itself was conducted batchwise inside a small (18 cm3) cylindrical test cell constructed of 316 stainless steel. The lid of the test reactor contained two ports, one for an incoming inert carrier gas stream and the other for the effluent. The test cell was typically filled with about 5 g of 99% pure lithium (initially in -4 + 16 mesh shot form), sealed, and then installed entirely inside the custom-made high-temperature thermoelectric calorimeter (International Thermal Instrument Co.). This assembly was put in an insulated box and then mounted inside a convection oven that precisely maintained reaction experiments at 616.5 K. A twoconductor wire attached to the calorimeter transferred the calorimeter voltage (heat) signal to an amplifier outside the oven and then to a data acquisition computer. The carrier gas section of the instrument, detailed in the upper left section of Figure 1, was used to maintain a continuous argon purge (1.6 cm3/s at 99.999% purity) through the calorimeter and, more importantly, to deliver a small water dose (2.0 µL or 111 µmol) to the test cell at specific times for reaction with lithium. This plumbing network included a septum-sealed injection port outside the oven that enabled the
introduction of the water dose to the carrier stream by means of a precision gas-tight syringe (Hamilton Model RN1701). In addition, a coiled length of carrier gas tubing inside the oven allowed the incoming carrier stream and water sample to reach oven temperature before entering the calorimeter assembly. The effluent stream from the calorimeter, which contained the inert carrier and any generated hydrogen, was directed into a thermal conductivity-type gas analyzer (MSA, Thermatron) just after exiting the convection oven. The gas analyzer produced an output voltage in a manner much like a thermal conductivity detector (TCD) in a gas chromatograph. This detector was carefully calibrated to provide the total quantity of hydrogen in the stream and was also used to detect unreacted water in the effluent, although none was evident in these experiments. Calibration and Operation Equations. The governing equation used to determine heat generation in a heat flux calorimeter is12,13
∫
Q ) Cc ∆Vc dt
(1)
where Q is the total heat generated, ∫∆Vc dt is the integrated calorimeter voltage signal, and Cc is a calibration constant determined from a standard electrical calibration procedure.13,14 During each lithium-water reaction experiment, the quantity ∫∆Vc dt is calculated for each water dose, and then the experimental enthalpy of reaction, ∆Hrxn, is determined by
∆Hrxn ) Q/nH2O
(2)
where nH2O is the molar water dose. The amount of hydrogen generated is similarly calculated from the analyzer detector voltage over time. In this case, the governing equation (given constant flow rate) has the power law form
∫
nH2 ) Cd[ ∆Vd dt]b
(3)
where nH2 is the number of moles of hydrogen generated, (12) O’Neil, M.; Louvien, R.; Phillips, J. Rev. Sci. Instrum. 1985, 56, 2312-2318. (13) Gravelle, P. C. Adv. Catal. 1972, 22, 191-263. (14) Berger, R. L. In Biochemical Microcalorimetry; Brown, H. D., Ed.; Academic Press: New York, 1969; p 221.
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∫∆Vd dt is the integrated detector voltage signal, and Cd and b are constants determined by regression of calibration data. Equation 3 offered a slightly better fit to the calibration data (correlation coefficient ) 0.999) than a linear fit over the expected range of hydrogen quantities (0-125 µmol). For all experiments reported here, Cc ) 0.162 J V-1 s-1, Cd ) 0.583 µmol H2 V-1 s-1, and b ) 1.184, as determined from regression of calibration data. Both the construction and calibration of this calorimeter are discussed in greater detail in a previous report.15 Procedure. Lithium was loaded into the test cell in an argon atmosphere, placed inside the calorimeter and oven, and then plumbed into the calorimeter flow system. A constant argon purge was maintained through the instrument at all times. The oven was then set to the desired reaction temperature; however, it typically required 16-24 h of thermal equilibration before the calorimeter produced the steady heat signal (baseline voltage) required for reaction experiments. When proper conditions were sustained, water doses of 2.0 µL were injected into the carrier gas stream about once an hour. This interval allowed ample time for complete calorimeter and gas analyzer signal response, including return of base line voltages to original levels. Both heat and hydrogen voltage signals were digitally recorded at 10 Hz for each dose interval. A total of 12 water doses were injected for each of three “runs”sdesignated runs A, B and Cswhere each run represented a fresh charge of lithium in the test cell.
Figure 2. Heat of reaction for lithium-water reaction experiments. Table 1. Enthalpies of Formation of Reactants and Products at 616.5 K compound
∆Hf (kJ/mol)
compound
∆Hf (kJ/mol)
LiH Li2O
-95.2 -606.5
LiOH H2O
-487.1 -244.9
The coefficient w is found by first applying an enthalpy balance to eq 4, resulting in
Results The lithium-water reaction equation at 616.5 K can be generally characterized in the following form:
∆Hrxn ) w∆Hf(LiOH) + x∆Hf(Li2O) + y∆Hf(LiH) ∆Hf(H2O) (9)
vLi(l) + H2O(g) f wLiOH(s) + xLi2O(s) + yLiH(s) + zH2(g) + ∆Hrxn (4)
where ∆Hf(x)’s are enthalpies of formation for each compound at reaction temperature (enthalpies of formation of Li and H2 are defined to be zero). Solving eq 9 for w and substitution of x, y, and v from eqs 5-7 yields
This reaction model assumes all possible intermediate and final reaction products including LiOH, Li2O, LiH, and H2. The compound lithium hydroxide monohydrate (LiOH‚H2O) was not included as a product because the reaction temperature is substantially above the reported decomposition temperature of LiOH‚H2O.11,16 Stoichiometric coefficients v, w, x, y, and z in eq 4 represent molar amounts of each compound, normalized by the molar amount of water reacted, nH2. Through atom balances, the coefficients are related by the following three equations:
x)1-w
(5)
y ) 2 - w - 2z
(6)
v ) 2w - 2z
(7)
Note that coefficients x, y, and v can all be evaluated in terms of w and z, and, as will be explained below, experimental heat and hydrogen measurements allow direct calculation of w and z. The coefficient z, for example, is computed from the measurement of generated hydrogen (eq 3), and the amount of injected water, or
z ) nH2/nH2O
(8)
(15) Bradford, M. C.; Phillips, J.; Klanchar, M. Rev. Sci Instrum. 1995, 66 (1), 171-175. (16) Popescu, C.; Jianu, V.; Alexandrescu, R.; Mihailescu, I. N.; Morjan, I.; Pascu, M. L. Thermochim. Acta 1988, 129, 269-276.
w) ∆Hrxn + 2(z - 1)∆Hf(LiH) - ∆Hf(Li2O) + ∆Hf(H2O) ∆Hf(LiOH) - ∆Hf(LiH) - ∆Hf(Li2O) (10) Enthalpies of formation are available from the JANAF Thermochemical Tables10 and are listed in Table 1 for these compounds at 616.5 K. Since the experimental heat (enthalpy) of reaction, ∆Hrxn, is determined via eq 2 for each water dose, all the right-hand side terms of eq 10 are known. This allows numeric determination of w followed by calculation of coefficients x, y, and v from eqs 5-7. Thus, a unique quantitative solution of eq 4 can be obtained for each water dose. Note also that the model can accommodate negative product coefficients. This simply implies that materials initially formed as products could eventually serve as reactants. This is certainly the situation in a batch process where experience has shown that intermediate products such as LiH and Li2O will react further with water as their overall concentration in the reacting mixture increases.17 Experimental heats of reaction from runs A-C are plotted in Figure 2 according to water dose number. Most of the values fall between 525 and 425 kJ/mol water reacted, which is generally consistent with (17) Phillips, J.; Bradford, M. C.; Klanchar, M. Energy Fuels 1995, 9 (4), 569-573.
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Figure 3. Average heat of reaction from runs A-C.
Figure 6. Stoichiometric coefficients of the overall lithiumwater reaction equation as determined by heat and hydrogen measurements.
Figure 4. Hydrogen generation for lithium-water reaction experiments.
Figure 5. Average hydrogen generation from runs A-C.
enthalpy balance calculations. Equations 11 and 12, for example,
Li(l) + H2O(g) f LiOH(s) + 0.5 H2(g)
(11)
4Li(l) + H2O(g) f 2LiH(s) + Li2O(s)
(12)
are two likely but different reactions identified in previous studies8,11,18 that yield calculated heats of 224 and 552 kJ/mol water reacted, respectively. The experimental values shown in Figure 2 represent some combination of eqs 11 and 12, as generally modeled by eq 4. (18) Chan, S. H.; Tan C. C. Combust. Flame 1992, 88, 123-136.
Although the data in Figure 2 are fairly consistent between runs, the heats vary on average about 50 kJ/ mol water reacted for each water dose. This variation arises because the measured heat flow quantities are extremely smalls about 50 J. Thus, the heat measurement can be extremely sensitive to small calorimeter signal errors such as a slight shift in the base line or a miniscule change in oven temperature over the dose interval. Additionally, slight inconsistencies in water dose size or lithium quality/purity could introduce further variations. Given these considerations, the data remain in a fairly tight band. Some experimental variation is eliminated by averaging the heats from all three runs at each dose. These average values, plotted in Figure 3, clearly demonstrate that higher dose numbers produce slightly smaller heats of reaction. The average heats are initially about 525 kJ/mol but decrease to around 475 kJ/mol by the final few doses. The corresponding hydrogen generation data for these runs are shown in Figure 4. All three runs show an increase in hydrogen generation with each water dose. Again, some variation exists in the data between runs because of the tiny amounts measured (about 30 µmol), so it is helpful to consider the average hydrogen generation for each dose. Figure 5, which plots these average values, conclusively shows a steady rise in hydrogen generation per dose. In fact, hydrogen quantities triple between the first and last dose in Figure 5. The trends of decreasing heat of reaction and increasing hydrogen generation indicate that reaction chemistry is changing as more and more water reacts. This can be further examined by computation of the stoichiometric coefficients as outlined in eqs 4-10. Figure 6 displays these coefficients when calculations are based on the average heat and hydrogen generation (Figures 3 and 5) at each water dose. Notice that the trends of heat and hydrogen are reflected in slight alteration of stoichiometric coefficients and thus reaction chemistry. For example, the first few water doses produce the approximate reaction equation
3.7Li + H2O f Li2O + 1.7LiH + 0.15H2
(13)
which is fairly similar to eq 12 except for the generation of a small amount of hydrogen. For the final dose, however, the reaction
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3.1Li + H2O f 0.9Li2O + 1.2LiH + 0.1LiOH + 0.35H2 (14) is apparent. Aside from the appearance of an inconsequential amount of LiOH, the primary difference between eqs 13 and 14 involve the reaction/formation of Li, LiH, and H2. As reaction experiments progress, the coefficients for lithium (v) and LiH (y) decrease, while that for hydrogen (z) increases. The coefficients for Li2O (0.9-1.0) and LiOH (0-0.1) remain approximately the same throughout the experiments. In addition, it can be concluded that Li2O is the favored oxygen-bearing product under these reaction conditions. Discussion The formation and identification of intermediate products LiH and Li2O, along with a clearer understanding of reaction progression in a batch-type system, are the significant findings of the experiments. At the start of each experimental run (when lithium is purest), the initial water doses produce primarily LiH and Li2O rather than hydrogen gas or LiOH. Once some LiH is present, both lithium and LiH react with subsequent water doses. Therefore, the situation eventually arises where part of the water dose will react with remaining lithium and part with accumulated LiH according to
LiH(s) + 0.5H2O(g) f 0.5Li2O(s) + H2(g) (15) This explains the trends of Figure 6 where coefficients of lithium and LiH are relatively high with initial water doses but then gradually decrease with additional doses. The corresponding increase in hydrogen generation is attributed to LiH further reacting with water as in eq 15 or lithium reacting with water to form primarily Li2O and hydrogen as in
2Li(l) + H2O(g) f Li2O(s) + H2(g)
(16)
Note that eqs 13-16 all include the formation of Li2O, which is consistent with its steady production throughout experiments. In other words, Li2O is a product no matter which reactant is involved. At these experimental conditions, Li2O appears to be the primary oxygen-containing product rather than LiOH. This observation is supported by previous lithium-water studies where Li2O is shown to be the favored product over LiOH at higher temperatures11 or when water exposures are relatively small.19 In addition, it has been reported that LiOH is not stable in molten lithium,20 which corresponds to the situation in the present study. (19) Hoenigman, J. R.; Keil, R. G. In Lithium: Current Applications in Science, Medicine, and Technology; Bach, R. O., Ed.; John Wiley & Sons, Inc.: New York, 1985; pp 233-255.
Although little LiOH formed in the present study, it is believed that some of the accumulated Li2O will eventually react with water according to
Li2O(s) + H2O(g) f 2LiOH(s)
(17)
The existence of this particular reaction was confirmed in a previous calorimetry investigation at roughly the same temperature.15 Equation 17 was also identified as a second-stage reaction in a study where pure LiH reacted with water to initially yield the products of Li2O and hydrogen.17 It is believed in the present study that the total amount of lithium converted was not sufficient for this final reaction to take place. These experiments were conducted well above the melting temperature of lithium (mp ) 453.7 K), and the overall purity of the lithium remained relatively high, since the total lithium conversion was only about onehalf percent of the original amount. However, lithium’s high surface tension and its tendency to form a “skin” prevented it from forming a regular liquid pool in the test cell. Rather, the exact state of the lithium in the test cell can probably best be described as a viscous slurry. The argon carrier stream (and water dose) entered the bottom of the test cell and traveled through the bulk of this slurry, but it was apparent that the carrier stream did not convectively mix the lithium. Rather, it likely made contact with the same region dose after dose. This channeling effect was visually evident in the test cell after experiments were completed. However, the exact configuration during experiments was difficult to realize because of the phase change during cool-down. If the test cell had been well mixed, it is expected that more water doses would have been required to observe the reported trends of heat and hydrogen production. The oven used in these experiments limited the operating temperature to 616.5 K, although the thermoelectric calorimeter itself can be operated up to 811 K. A lithium-water reactor in a combustion application is likely to operate at temperatures beyond this study; thus, it would be beneficial to extend experiments at even higher temperatures to determine if the same trends are evident. Additionally, reaction chemistry in a batch system could be explored at later stages when greater amounts (or doses) of water have been reacted. This would provide an accurate model of heat and hydrogen generation over the complete batch reaction process, that is, from pure lithium to pure LiOH. Acknowledgment. This work was performed under the sponsorship of the Office of Naval Research. EF970047E (20) Schreinlechner, I. E.; Sattler, P. F.; Kozuh, J. In Lithium: Current Applications in Science, Medicine, and Technology; Bach, R. O., Ed.; John Wiley & Sons, Inc.: New York, 1985; pp 207-216.
