770
J. Phys. Chem. 1901, 85,778-782
Isotope Effects in the Lanthanum Dihydrides Dean H. W. Carstens Los Alamos National Laboratory, Los Alamos, New Mexico 87545 (Received: August 12, 1980)
Equilibrium pressures of protium, deuterium, and tritium over lanthanum metal at atomic ratios up to 2.2 have been measured for temperatures ranging from -800 K to slightly over 1000 K. The equilibrium pressures in the plateau regions increased as expected from the protide to the deuteride to the tritide. Sieverts’ constants were determined at the higher temperatures and decreased in this same order. Within experimental error the terminal solubilities of the three isotopes in the metal, defined by the separation of the dihydride phase from the metal, are the same. Introduction Because they form extremely stable hydrides, lanthanide metals and their alloys have been proposed as tritium Additionally, alloys getters in fusion-reactor te~hnology.~~ of the lanthanides have been suggested by numerous authors as hydrogen-storing media for various energy applications utilizing hydrogen. Because a sound, experimental data base for the pure metals is needed before the alloys can be totally understood, and because no one to our knowledge has simultaneously studied all three isotopes of any lanthanide hydride, we felt that a closer look at the three lanthanum dihydrides was in order. The emphasis of this paper was on the low-concentration region, the so-called Sieverts’ region, of the pressure vs. composition isotherms. Several workers have previously measured the equilibrium pressures of hydrogen over lanthanum, among them Mulford and H ~ l l e ywho , ~ looked only at the protide? and Korst and Warf,6 who studied both the protide and the deuteride. In ref 2 we also reported data on the lanthanum dideuteride which are compared to similar data on the deuterides of several alloys of lanthanum and nickel. In none of these works was the Sieverts region studied in detail. This report discusses the measurements of equilibrium pressure of protium, deuterium, and tritium over lanthanum metal over the concentration region of the dihydride (hydrogen-to-metal atomic ratios from essentially zero to 2.2). Particular care was taken to obtain enough data points in the Sieverts region to adequately define solubilities and miscibility gaps. Although background pressures led to problems in obtaining accurate data at the lower temperatures (below 850 K), at temperatures between ca. 850 and 1025 K Sieverts’ law was obeyed and values of the Sieverts constant were obtained for all three isotopes. Experimental Details Typical Sieverts-type experiments were performed; that is, small measured amounts of hydrogen gas were successively added to, and allowed to equilibrate with, the lanthanum metal held at constant temperature. Following each equilibration the hydrogen partial pressure was then (1) D. H. W. Carstens, J. Nucl. Mater., 73, 50 (1978). (2)D.H.W.Carstens, J. Less-Common Met., 61,253 (1978). (3)R.N.R.Mulford and C. E. Holley, J. Phys. Chem.,69,1222(1955). (4)Throughout this paper we will use the symbols H, D, and T to refer to the three isotopes protium, deuterium, and tritium, respectively. The use of the first for protium is somewhat inconsistent; however, it follows general usage and avoids confusion with the use of the symbol P to denote pressure. The words hydrogen or hydride will be used to denote the three isotopes in general, and, additionally,in later equations the symbol H will refer to any of the three. Again this is inconsistent but follows general usage. (5)W.L. Korst and J. C. Warf, Inorg. Chem., 6,1719 (1966). 0022-365418112085-0778$01.25/0
accurately measured. From the molar amounts of gas added and remaining at equilibrium, the atomic ratio of hydrogen in the metal could be readily calculated by assuming the ideal gas law. The gases were measured in a calibrated volumetric apparatus kept thermostated at 40 “C. The apparatus was similar to that previously described in detai1.l Pressures were measured with either a 1000- or l-torr “Barocel” capacitance manometer (Datametrics, Wilmington, MA). Such manometers are not designed for tritium work since they are double-sided; i.e., they have sensors on both sides of the measuring diaphragm. Use of this type of apparatus leads to inaccuracies at high pressures with tritium because it is an ionizing gas. However, comparisons with a similar but single-sided manometer indicated that the error varied, approximately linearly, from zero at low pressures to only 2% at 100 torr of tritium. For the experiments described here, tritium pressures were always below 80 torr, and this error is insignificant compared to the accumulated errors in the P-V-T measurements of hydrogen content. The lanthanum metal (obtained from Research Chemicals, Phoenix, AZ)had a stated purity of 99.9%; no further analyses were made. In most cases a fresh sample was used for each run. Samples of a uniform size were prepared by cutting chips from the metal ingot with a lathe. The samples were weighed and loaded into the furnace within a few minutes of their cutting to maintain purity. The chips were contained in a molybdenum crucible inside a copper tube furnace which was externally heated and surrounded with a vacuum jacket to catch any permeating tritium (none was detected). Temperatures were measured with chromel-alumel thermocouples placed just outside of, but in contact with, the copper tube. The temperatures were read with a digital temperature indicator calibrated at the ice point and the temperature of melting NaCl(lO74 K). Tests to determine the temperature profile of the sample within the furnace indicated that the external temperature agreed reproducibly with the sample temperature to within a few degrees. During a given isothermal run, the indicated temperature varied less than 1 K. The compositions of the gases were checked before each experiment with a low-resolution quadrupole spectrometer (Inficon IQ-200, Inficon, East Syracus, NY). This instrument has a large mass discrimination for the hydrogen isotopes and was cross-checked against a more conventional sector mass spectrometer (CEC Model 21-201 equipped with solid-state electronics). The protium used was isotopically pure, whereas the deuterium contained 1-2% protium. The tritium gas used was isotopically impure as it had been recovered from hydrides used in previous experiments. It assayed 88% tritium, 10% deu0 1981 American Chemical Society
The Journal of Physical Chemistry, Vol. 85, No. 7, 1981 779
"*!
'1020'K' 971 K A 913K v 876 K + 829 K 725K o
IO' F
I
'
I
'
'
I
L j
I I
7
-
-
IO'
2 ,--
IO0
?/A;[ io0
-
-: L
10-1
n
a
10-2
10-2
....
IO-^
I
I 10-1
I 00
D/La
Flgure 2. Pressure vs. composition isotherms for deuterium over lanthanum metal at six temperatures.
.
1018 K 976 K A 918 K
I
E
/
4
/
T/ La
Flgure 3. Pressure vs. composition isotherms for tritium over lanthanum metal at four temperatures.
Each experimental isotherm, as expected for lanthanide dihydrides, consists of three main p a r k 7 The first of these, the so-called Sieverts region, is the linear (on the log-log plot), rising portion of the curve a t low H/La values. In this region hydrogen is dissolving in the metal as atoms with no separation of a second phase. As more and more gas is added to the sample, the metal becomes saturated and lanthanum dihydride begins to precipitate. According to the phase rule, because of the existance of two solid phases, the LaH2-La system becomes invariant and the pressure does not change with added gas. This is the second main region of the experimental curve known as the plateau region. The terminal solubility of the dihydride in the metal is defined here by the intersection of the rising Sieverts curve with the plateau. Each of these three, the Sieverts region, the plateau region, and the terminal solubility, will be discussed in more detail in later sections. As the H/La ratio approaches 2, a point is reached at which all of the La metal is converted to the dihydride and (6) T. Takaishi and Y. Sensui, Trans. Faraday SOC.,59,2503(1963).
(7) D. H. W. Carstens, LOBAlamos National Laboratory Informal Report LA-7602-MS, 1979.
700
The Journal of Physical Chemistry, Vol. 85, No. 7, 1981
Carstens 1
-
0.e
Y
l-
z
Ez
srn k2 W
>
w rn
D/La Flgure 4. Sieverts plot thanum metal.
vs. composition) for deuterium and lan0.I
the system again returns to a single solid phase. At this point the equilibrium pressure again rises but at a much faster rate than in the Sieverts region, as is evident in the figures. With increasing temperatures, the two phase boundaries approach each other until, as reported by Peterson and Straatmann,8 at a temperature of 1233 K and a t 45 at. % hydrogen, the two merge. At the lowest concentrations (H/La below the data points did not fall on the linear lines but rather leveled off at pressures of 10"-104 torr. This effect is believed to be due to the background pressure. Because of possible confusion, data points for such low compositions were not included in the graphs. Above -850 K, the increasing equilibrium pressures rose to the point that the background was less bothersome, and, as will be seen, from later Sieverts plots of the data it was possible to measure and correct for the background pressure at these higher temperatures. Sieverts' Region. For the Sieverts region the equilibrium reaction is 1/2H2(gas)= H(so1ution in metal) with an equilibrium constant K defined by K = aH(in metal)/aH,'/2 N (H/La)/PH:/2
(1)
(2)
The approximation arises from the assumption that the activity of hydrogen gas is its pressure and that the activity of the dissolved hydrogen is its concentration. Both assumptions are approximately correct at these low pressures and concentrations. Rearrangement of eq 2 leads to Sieverts' law (H/La) = KP1/z (3) where the equilibrium constant K is the so-called Sieverts constan t. (8)D. T. Peterson and J. A. Straatmann, J. Phys. Chem., 70, 2980 (1966).
IOOOIT
Flgure 5. Variation with temperature of the Sieverts constant for the three hydrogen isotopes.
From eq 3 it is evident that plots of the square root of pressure vs. H/La should lead to a straight line if Sieverts' law is obeyed. This was indeed found to be true for all three isotopes for data above 850 K, and a representative case is given in Figure 4 for the deuteride. The other two isotopes gave similar plots and are not shown to conserve space. Because of the finite background pressure at the start of the runs, the Sieverts plots did not pass through the origin. This is evident from Figure 4 and was seen in all cases. In order to take this into account analytically, I fitted the data for each temperature to an equation of the form (H/La) + (H/La)o = KP112 (4) where the zero subscript refers to the initial background concentration. The fit to the curve was determined by eye since the data did not warrent a least-squares treatment. Because of the falloff from the straight line as the terminal solubility was approached, the fit to the data points was somewhat arbitrary, leading to additional uncertainty in the values obtained for K. In Figure 5 the logarithms of the Sieverts constants obtained from eq 4 are plotted against 1/T for the three isotopes. In the tritium case the data point at the highest temperature seems grossly in error, and this point was not used in the fit. The trend in the value of K for the three isotopes is definite despite the errors, and it decreases from protium to deuterium to tritium. The three lines as drawn fit eq 5-7. In these equations for protium (5) In K = 8700/T - 10.1 In K = 8700/T - 10.3
for deuterium
In K = 8300/T - 10.0 for tritium the units of K are atom ratio X torr-'f2.
(6)
(7)
The Journal of Physical Chemistry, Vol. 85, No. 7, 1981 781
Isotope Effects in the Lanthanum Dihydrides
0.4
I
I
I
I
\
. . 'V
-I
0,041
Flgure 6. Terminal solubilities of hydrogen isotopes in lanthanum metal compared to the literature data (ref 7, dashed line) for protium. The solid line represents the least-squares fit to my data for the three isotopes.
Terminal Solubilities. The terminal solubility is that point at which the dihydride begins to precipitate and, as mentioned, is determined by the intersection of the linear Sieverts curve with the plateau line of constant pressure. Obviously the change on going between the two regions is more gradual than implied by this definition. The values obtained from the intersections of the curves are plotted for all three isotopes in Figure 6. For those cases in which adequate Sieverts plots were obtained, the initial concentration (H/La)o as found from the intersection of the straight lines with the horizontal axis was added as a correction. In the other cases the initial concentration of hydrogen in the metal was assumed to be 0.01. From the data, there does not seem to be any difference in the terminal solubilities of the three hydrogen isotopes in lanthanum. The solid line represents the least-squares fit to all data points neglecting the two lowest values found in two deuterium runs. Unfortunately the error in the values of the terminal solubilities is large, and the numbers are good only to several percent. Using a different experimental technique, Peterson and Straatmann* have defined the phase boundaries for the lanthanum-protium system. Their data over the temperature region of our study are also shown in the figure, and their equation for the miscibility gap, primarily defined by data at lower temperatures, is shown with a dashed line. The agreement between the two is not unreasonable in view of the different techniques used in determining the terminal solubilities. Equations for the two lines in Figure 6 are given by
+ 2.3 In S = -2900/T + 0.92
In S = -3800/T
(this work)
(8)
(ref 8)
(9)
where S is the terminal solubility expressed as atomic ratio (H/La). The approach to the terminal solubility point was also evident during the experiments and was indicated by a gradual decrease in equilibration rates followed by a noticeable increase in reaction rates after the point was passed. This effect is not unexpected: as the solubility limit is approached, fewer and fewer sites are available for the dissolving gas. Equilibration should be particularly slow since, for the chip samples used, equilibration must occur via the gas phase rather than the faster, diffusion through the solid, route. At these low pressures gas
I"
1.0
0.9
1.2
I I
13
14
1000/T Flgure 7. Variation with temperature of the equilibrium pressure of the three hydrogen isotopes over lanthanum metal.
TABLE I: Comparison of Equilibrium Pressure Data at Two Temperatures P,torr (at 1100 K) P, torr (at 800 K) isotope
H D T
this work
7.3 9.3
ref 4
this work
ref 4
6.1
2.0 x 10-3 2.2 x 10-3 2.4 x 10-3
1 . 2 x 10-3 2.5 x 10-3
7.2
12
transport would be very slow. Plateau Region. In the plateau region the equilibrium equation is H2 La(saturated with H atoms) = LaH2(solid) (10)
+
and, because of the coexistence of the two phases, the pressure must be invarient. This was indeed found, as can be seen from Figures 1-3. The data for the plateau pressure (at H/La = 1) along with the lines fitted to the points by eye are shown in Figure 7 for the three sets of data. The lines appear to cross near 800 K, above which point the pressures increase with the atomic weight of the hydrogen isotope. This is the order normally expected for the isotope effect, although Wiswall and Reillye have reported several systems for which the opposite is seen. The equations for the three lines are (pressure is expressed in torr) In P = -24000/T + 23.8 for the protide (11) In P = -24500/T
+ 24.5
In P = -24900/T
for the deuteride
+ 25.1
for the tritide
(12) (13)
The values for the coefficients for the deuteride case agree within 2% with those found in our previously reported work.2 The value of the slope, which is directly related to the heat of formation of the d i h ~ d r i d eincreases ,~ with the (9)R.H.Wiswall, Jr., and J. J. Reilly, Inorg. Chern., 11, 1691 (1972).
J. Phys. Chem. 1981, 85, 782-785
782
TABLE 11: Thermodynamic Constants for the Solution of Hydrogen in Lanthanum Metal (Reaction 1 ) isotope AH", kcal/mol AS", eu
H D
- 13
- 17 - 17
- 14
- 16
T
-13
atomic weight of the gas. This again is in direct contradiction with the work of previous authom6 In Table I the equilibrium pressure at two temperatures which embrace the extremes of the two sets of data is calculated from the equations given here and in the work of Korst and Warf.6 In view of the different samples and techniques, the agreement is probably as good as can be expected. Korst and Warf did state that much of their data was for regions in which their manometer (a McLeod gauge) had limited accuracy. The numbers in Table I indicate that slight changes in measured pressures can lead to important changes in thermodynamic quantities. Slight nonsystematic errors in temperature measurements are probably responsible for the scatter in the data and are believed to be the second largest source of errors in this work (after the effects of the background pressure). Since the same manometer was used in all pressure measurements, errors here would affect the absolute pressure values but not the isotopic trends. Small leaks or outgassing of parts is another large source of error at these low pressures. In the mass spectra of the overgas, in addition to the main hydrogen peaks a small peak at mass 28 was always seen. This could arise from either N2,indicating a small leak (a corresponding O2peak would be absent since this gas would be more reactive with La than N2),or more probably from CO. In all cases the peak at mass 28 was 2-3 orders of magnitude less than the hydrogen peaks; thus, an atmospheric leak was not indicated as a significant error in these experiments. Thermodynamic Quantities. Since the Sieverts constant is the equilibrium constant for reaction 1,from appropriate
TABLE 111: Thermodynamic Constants for the Formation of Lanthanum Dihydride (Reaction 10) isotope AH', kcal/mol ASo, eu H -47.5 - 34.0 D -48.5 -35.4 T -49.3 - 36.6 thermodynamic relations (as defined in ref 7) one can derive eq 14. By comparison with eq 5-7, it is then In K(torr)-1/2= -AHo/(RT) + A S o / R - yz In 760 (14) possible to calculate the standard heats and entropies of the dissolution of hydrogen gas in lanthanum metal. These data are given in Table 11. Apparently in only two other cases have equilibrium data been obtained in the Sieverts region for all three isotopes. The first was lithium metal, reported recently by Smith et al.l0 (Note that these authors have defined the Sieverts constant as the inverse of that here). They found similar trends in the variation of AHo: the heat of reaction decreased with increasing atomic weight of the hydrogen. The heats of reaction, of course, differed for the lithium metal, being roughly 10% more negative than the values for lanthanum reported here. Begun et al." looked at the isotope effect in the yttrium-hydrogen system, again finding similar trends. For the plateau region it can be shown similarly that In P(torr) = AH"/(RT) - A S o / R + In 760 (15) where AHo and ASoare the standard heat and the entropy of formation of the dihydride, respectively. These data are given in Table I11 for reference purposes. Acknowledgment. I thank W. R. David for help with some of the experiments. This work was performed under the auspices of the U.S.Department of Energy. (10)F. J. Smith, J. F. Land, G. M. Begun, and A. M. La Gamma de Batistoni, J. Inorg. Nucl. Chem., 41, 1001 (1979). (11)G. M. Begun, J. F. Land, and J. T. Bell, J. Chern. Phys., 72,2959 (1980).
Collisional Deactivation of O('D,) by the Halomethanes. Direct Determination of Reaction Efficiency A.
P. Force and J.
R. Wiesenfeld'
Department of Chemistry, Cornell University, Baker Laboratory, Ithaca, New York 14853 (Received: October 2, 1980)
-
The deactivation of 0(lDz)by halomethanes has been studied by monitoring the absorption-of atomic resonance radiation 33s1 23PJfollowing KrF laser photolysis of ozone in dilute helium mixtures. Both the overall rate constants for collisional deactivation and the efficiency of chemical reaction vs. physical quenching could be determined in this manner. Total deactivation rates are in good accord with previous studies, except for the molecules CFBHand CF4. The present experiments indicate that chemical reaction plays a dominant role in the deactivation process for all molecules studied except CF3H and CF4.
Introduction Chemical reactions of electronically excited oxygen atoms, O(lD,), play a major role in establishing the chemical composition of the stratosphere.' A number of experimental techniques have been exploited in recent years to examine the deactivation of this optically metastable (1)R.J. Cvetanovic, Can. J. Chern., 52, 1452 (1974).
species (7 = 140 s), special emphasis having been placed on the m m m ~ - ~ ~ofe ntotal t W IX IX~ rates of O('Dd (hereafter o*)upon collision with the atmospheric gases. Because of the potential role of chlorinated hydrocarbons in catalyzing the destruction of ozone in the stratosphere? (2) F. S. Rowland and M. J. Molina, Rev. Geophys. Space Phys., 13, 1 (1975).
0022-3654/81/2085-0782$01.25/00 1981 American Chemical Society
