R. J. ACKERMANS, E. G. RAUH,R. J. THORN, AND &I.c. CANNON
762
of a t least one process whose a value depends upon Q-in the present example, one or a group of unimolecular gas phase decompositions of radical intermediates competitive with radicsil termination processes. Before this conclusion is accepted, it is important that some additional hiformation coiicerniiig this (these) unimolecular reaction(s) be developed. Let us assume, for the sakes of argument and some crude computations, that kh is sufficiently large that, under conditions of illumination like those employed in the experiments described abovf:, (the fraction of radicals terminating a t the wall) may be neglected; further, assume that kb has no associated activation energy, and so has a high value such as that calculated above, eq. 13: i’ch = 2-88 X ~ m molecule-1 . ~ see.-’ = 1.73 X l O I 4 cm.3 mole-1 set.-'. When u = 0.5, the steady-state condition is equivalent to the approximation
kuR
+ khR2
=
2khR2
=
Q
(19)
from this equality, k h , and the kiiown range of Q, we , calculate R = (1.2 - 20.4) X 10-14mole ~ m , - ~The corresponding range of ku is (2.0 - 35.2) sec.-l; this is the specific rate constant for a relatively slow unimolecular reaction-certainly, it is much too low to be
1‘01. 67
characteristic of the decompositions of excited radicals. Additional complexity of the photolysis mechanism can be considered in terms of equations similar to those derived above. Mention will be made here of oiily one other possibility: that in which just one of the reactions H, occurs when the in eq. 18, say COOH + C 0 2 radicals are hyperthermal, while the others may or may not be effective competitors with the earlier described radical termination processes. In such a situation, provided f,“ and fh differ significantly, the character of the plot of f ( C 0 , ) us. Q - ’ / 4 depends upon the relation of lchR to the mean k , for the non-hyperthermal unimolecular decompositions exactly as discussed a t length above, because afa for the hyperthermal reaction is independent of Q. Polychromaticity and nonuniformity of illumination seem to be much more attractive explaiiations of the possible deviation from linearity of the plot in Fig. 1 than the assumption of slow uiiimolecular radical decomposition processes. Acknowledgments -This research was sponsored by the U. S. Atomic Energy Commission. We are indebted to Professors Aron Kuppermaiin and R. Linn Belford for helpful interest and discussion, and to a referee for his suggestion that we consider our original neglect of uiiimolecular reactions in the gas phase.
+
A THERMODYSAlLTIC STUDY OF THE THORIUM-OXYGES SYSTEM AT HIGH TER/IPERATURES1 BY R. J. ACKERMASS, E. G. RAUH,R. J. THORN Argonne Sational Laboratoiy, Argonne, Illinois AND
&I.c. CANNON
Department of Chemistry, Ctah State University, Logan, Utah Received August 8, 1968 The evaporation behavior of the thorium-oxygen system in the temperature range 2000-3000°K. has been studied. The results of the effusion measurements of the dioxide phase and liquid metal-dioxide mixtures and a mass spectrometric investigation of the former have been combined to yield an internally consistent set of thermodynamic data for the system. The solid dioxide evaporates congruently a t all temperatures. Above 2800’K. there iai a meahrable but thermodynamically insignificant substoichiometry, Th01.998. The “effective” pressure based on: the assumption that the vapor is comprised entirely of dioxide molecules is given by the equation log p e (atm.) = (8.26 f 0.13) - (3.55 =t0.03)104/T. The average heat and entropy of sublimation of the solid dioxide to ThOz(g) are 158.7 =k 2.5 kcal. mole-’ and 35.3 =t 1.0 e.u., respectively, and the respective values for the reaction ThOz(s) = ThO(g) O(g) are 347.0 rrt 3.1 kcal. mole-’ and 76.9 =t 1.4 e.u. The 11.4T standayd free energies of formation of the gaseous thorium oxides are AFP(Th02) = -138,600 (2000-3000°K.) and AFro(Th0) = - 10,300 - 14.4T (2000-3000’K.). Upper limits t o the dissociation energies a t absolute zero of ThOz(g)and ThO(g) are Do(ThOz)5 16.3 e.v. and Do(Th0) 5 8.3 e.v. The reaction occurring in thoriated tungsten filaments is discussed.
+
Introduction Although thorium is generally considered to be the first member of the 5f “actinide” series, the high teniperature chemical behavior of its oxides closely parallels that of the oxides of the group IV transition metals, zirconium and hafnium, in that the +4 valence state in the solid and both +4 and +2 states in the vapor exhibit paramount stability. This similar behavior is perhaps to be expected since the electronic configuration nd2 (n l)s2 is common to all three metals. Because of the refractory nature of thorium dioxide, vaporization studies utilizing the effusion method are a convenient means of obtaining chemical bond energies
+
(1) Based on work performed under t h e auspices of t h e U. 9. Atomic Energy Commission.
+
and, more importantly in a thermodynamic sense, the standard free energies of formation of the gaseous molecules which exist at high temperatures. Previous determinationsz-4 of the vapor pressure of Tho, have been critically reviewed5 and appear to be in error because of inaccurate measurement of the temperature. reduction of the sample by the effusion cell, and nonequilibrium conditions resulting from too little sample. ~ 6 the vapor pressure of thorium Darnell, et ~ ~ 1 .report (21 E. Shapiro, J . Am. Chem. Soc., 74, 5233 (1952). (3) 31. Hoch a n d H. L. Johnston, ibid., 76, 4833 (1954). (41 R. J. dckermann and R. J. Thorn, “Vaporization Properties of Thoria,” 133rd National Meeting of the American Chemical Society, San Francisco, Calif., 1958. (5) R. J. Ackermann and R. J. Thorn, Ch. 2 of “Progress in Ceramio Science,” Pergamon Press, Oxford, 1961.
THERMODYXAMIC STCDYOF THORIUM-OXYGEN SYSTEMS
April, 1963
metal and the effect of oxygen contamination thereon from which the dissociation energy of ThO(g), 8.5 e.v., mas estimated. More recently, Darnell and McCollum7 have measured effusion rates for both solid dioxide and mixtures of inet,al and dioxide. I n the case of the dioxide the method of determining the effusion rates over a 300” temperature range yielded data which when treated in accordance with the second law of thermodynamics resulted in a precision of Ct3 kcal. mole-l in the heat of sublimation. An attempt to increase the precision in the heat of sublimation by a third law treatment of the data is precluded since there are no reliable free energy functions for the gaseous thorium oxides. I n the case of the mixtures the monoxide was shown mass spectirometrically to be the principal gaseous species, but since the solubility of the dioxide in liquid thorium was found to be rather extensive and significantly temperature-dependent, the activities of the components in the condensed phases cannot be accurately obtained by the assumption of Raoult’s law and hence an accurate evaluation of the thermodynamic properties of the gaseous monoxide is not possible. The experimental goals of the present investigation were these: (1) To measure more precisely the effusion rate of thorium-containing species from solid dioxide over a large range of temperature in order to miiiimize the effect of experimental errors in the subsequent treatment of the data by means of the second lam of therniodynamics. (2) To measure the effusion rate of gaseous monoxide from a metal-dioxide mixture under experimental conditions where the solubility of the dioxide in the liquid metal is minimized and practically insignificant. (3) To examine mass spectrometrically the composition of the vapor and the temperature dependence of the predominant species in equilibrium with the dioxide phase. (4)To demonstrate that the dioxide phase evaporates congruently and essentially stoichiometrically. From these measurements it will be possible to derive the standard free energies of formation and dissociation energies for ThO(g) and TliOz(g) and to demonstrate the extent of reduction of thoria by tungsten in a discussion of the reaction occurring in thoriated tungsten filaments. It will be shown that tungsten does not significantly reduce thoria a t high temperatures and it is suggested that carbon is present in tungsten filaments as an impurity and accounts for the reduction of thoria and the evaporation behavior observed by previous investigators. Experimental Methods and Apparatus A variety of experimental adaptations of the effusion method were utilized in the present investigation in order to minimize systematic errors. Basically, the effusion method allows one to calculate the vapor pressure p of a condensed phase from the measured mole effusion rate z (moles cm.-2 sec.-l) of the vapor having a molecular weight M and an absolute temperature T v m the equation
p
=
Z(Z~TMRT)”~
(1)
All effusion cells used i n the present investigation were made of ( 6 ) A. J. Darnell, W. A. McCollum, a n d T. A. Milne, J . Phys. Chem., 64, 341 (1960). (7) A. J. Darnell a n d 1%’. 8. McCollum, NAA-SR-6498, September 15, 1961.
763
tungsten. Temperature measurements were carried out by means of a calibrated optical pyrometers and it is believed that the absolute accuracy of these measurements is within 10’ over the range of the present investigation. All temperatures reported herein correspond to the cavity of the effusion cell and have been corrected for an interposed window and prism. The samples of thorium and thorium dioxide were 99.97, pure as established by spectroscopic analysis. Mass spectrometric analysis confirmed that any volatile impurities were less than 1% of the thoriumcontaining vapor species. Subsequent to the evaporation studies spectroscopic analysis showed that the dioxide samples contained less than lY0tungsten. The mass effusion rate of Tho:, from a tungsten cell was measured in a collection-type apparatus essentially the same as described previously.8 The amount of thorium collected on the plates was determined by neutron activation analysis shown by the scheme
23.3m p27.4d p-+ Pa233 ___ +r 2 3 3 Th232(11,y)Th233\313
kev. y
‘I
The condensation plates were 0.75 in. in diameter and 0.005 in. thick and were made from Corning’s ultrasonic grade of fused silica which contains less than a few parts per million total impurities. After exposure in the vapor pressure apparatus a series of the plates containing unknown amounts of thorium were packaged for thermal neutron irradiation by separating the targets with fused silica annuli and by interspersing plates containing known amounts of thorium as standards. The standard plates were prepared by transferring t o the plate with a micropipet a volume of an aqueous solution of known concentration of thorium and then evaporating to dryness in a desiccator. In general, for plates containing of the order of lo-’ g. 2-hr. irradiations were sufficient to produce reliably measurable amounts of Pa233. The amount of Th232present on the plates was then determined by quantitatively counting the 313 kev. yl0 of Pa233by means of a y-ray spectrometer. The plates with known amounts of Th232provided the necessary relationship between counts per minute and weight of Th232. It also was possible to utilize the 98.4 kev. K a X-ray” of U233which is emitted simultaneously with the 313 kev. y of Pa233. The total mass effusion rate of thorium dioxide and of mixtures of thorium metal and dioxide were measured over a relatively short range of temperature by means of the vacuum balance and effusion cell assembly similar to that previously described.12 Relative rates of evaporation of Tho2 and T h o were measured by means of a Bendix Model 12-101 time-of-flight mass spectrometer. The tungsten effusion cell assembly and the operational characteristics of the instrument have been described previously.12
Results Three series of collector plates were exposed to the vapor effusing from tungsten effusion cells containing approximately 0.5 g. of Tho2 and these plates along with standard plates subsequently were analyzed for thorium by neutron activation analysis. The results in order of exposure of one of these series, which is typical of all the measurements, are given iii Table I. The weight of ThOz on each plate determined by activation analysis is given in column 2. The “effective” pressure given in column 5 was calculated from eq. 1 in which 2 = ( W / M a t ) [(a? +)/r21, M = 264 based on the assuinptioii that the vapor was comprised en-
+
(8) F. Hoffmann and C. Tingwaldt, “Optisohe Pyrometrie,” Edwards Bros.. Inc., Ann Arbor, Mich., 1944. See also Natl. Bur. Standards (U.S.) Tech. New Bull. 43 ( 6 ) . 114 (1959). (9) R. J. Ackermann, P. W. Gilles, and R. J. Thorn, J . Chem. Phys., 26, 1089 (1956). (10) D. Strominger, J. M.Hollander, and G. T. Seaborg, Rev. Mod. Phys., SO, 814 (1958). (11) C. E. Crouthamel, “Applied Gamma-Ray Spectroscopy,” Pergamon Press, Oxford, 1960. (12) R. J. Ackermann a n d E. G. Rauh, J. Chem. Phys., 36, 448 (1962).
R. J. ACICERMAKN, E. G. RAUH,R. J. THORN, ISD 11. C. C a m o s
764 I
I
I
I
I
Yol. 67 I
I
D Vacuum Bolonce
0 Series
4.0
x Series B Activotion Analysis A Series G Activation Analysis Series E Activation Analysis
5.0
F4
log pe = (8.26f 0.13)-(3.552 k 0.031) h
c E
6.0
0 Y
Q
-
0
0
0
7.0
8.0
9.0
1 3.4
3.6
3.8
4.O
4.2
4.4
4.6
4.8
5.0
!Q4. T
Fig. 1. -The t,emperature-dependence of the “effective“ pressure and the partial pressures of predominant vapor species in equilibrium with solid thorium dioxide.
tirely of ThOz(g), and a,r, and d have the values given in Table I. The total mass effusion rate of Tho, from a tungsten effusion cell containing a channel orifice (diameter
= 0.158 cm., length = 0.299 cm.) mas determined by means of the vacuum balance.l2 The dimensions of the orifice necessitated the use of a Clausing factor,13 K = 0.417. The resulting data are shown in the order of determination iii Table 11-A. The weight losses of TABLE I the empty tungsten effusion cell were determined at six TEE E ~ s r j s ~ oRATE s OF T h o 2 : SERIESE WITH STANDARD arbitrary temperatures over the range of the effusion PLATES measurements and the particular weight loss at a chosen Counts/min.a temperature then was calculated from a least squares 313 kev. y ThOn, T, t, - log P o , sec. atm. (I. Pa238 equation. The calculated weight loss a t each tem3600 7.064 2326 683 5.72 X l o u 7 perature is shown in column 2, and in column 3 is 900 5.418 5.84 x 2233 6981 given the corresponding net weight loss of Thoz. The 2700 6.766 1007 8 . 4 3 X lo-’ 2374 “effective” pressures were calculated from eq. 1 in 8400 7.475 2264 627 5.25 X which 2 = 1,04W/MKat. The factor of 1.04 was 1200 5.731 3.94 X l o M 6 2543 4704 necessary to correct the observed weight losses for the 10800 7.980 2180 257 2.15 X lo-” fraction of the effusate that condensed on the support 5.138 2661 600 8967 7.50 X rods and hence was not weighed by the balance. The Counts/ results of all measurements of the effusion rate of ThOz inin./lO-G g . are shown in Fig. 1. If one calculates an effective 4034 3.42 X 1180 14276 1 . 2 1 X 10-6 1179 pressure p e by assuming the vapor to be comprised en1204 Standard 1715 1.43 X tirely of dioxide molecules, the experimental data can 1221 plates 17397 1 . 4 3 X 10-6 be represented by the equation OK.
16972
1 . 4 3 X 10-6
1191 Av. = 1195 i 8
Orifice area ( a ) = 8.04 X Effective target radius ( T ) = 0.960 cm. Orifice-to-target distance ( d ) = 7.746 cm. a Corrected for background count of the silica targets, which is 39 i 1 counts/min.
log p , (atm.)
=
(8.26
=t0.13)
-
(3.53
* 0.03)104/T
(2)
4 limited investigation of the effusion rate of ThO(g) (13) S. Dushman, “Scientific Foundations of Vacuum Techniques,” John Wiley and Sons, Kev Yoik, N. Y., 1949.
April, 1963
765
THERMODYNAAIIC STUDY O F THORIGM-OXYGEK SYSTEMS
from a mixture of liquid metal and solid dioxide was carried out. The tungsten effusion cell, similar t o that TABLEI1 THE EFFUSIOK RATEOF Tho$: A. SERIESD-VAccuIvf BALAXCE
wx
I
I
I
I
Tho:
px
Thoe:
I
I
I
x
A H 0 = 175.5 f 2.7 o v e 173.5 0 AH' i 169.8 f 3.7
X
AH' = 158.4 .t 5.3
0
A H 0 = 160.0 f 2.7
_
f 2.2
159.5
2.5
4.0
2', OK.
103 (empty cell)
W X IO8, g . ThOz
sec.
atm.
2642 2720 2822 2694 2871 2544
0.4 0.7 0.9 1.0 1.4 0.5
2.2 3.2 3.1 4.3 5.6 3.1
2520 1800 720 3660 600 12120
5.101 4 785 4 393 4.967 4.053 5.642
t,
4.5
--log P o
3.5
'\
\
Tho
Log I + T ,
RATEOF Th-ThOn: R. THE EFFVSION SERIEB F-VACUUMBALASCE
3.0 -log
TI' x 103, g . Tho
t,
see.
pT110,~
atm.
2337 5 2 2700 4 614 2381 2 5 900 4 451 2369 3 0 1200 4 498 2366 3 0 1200 4 498 2362 2:' 1200 4 544 Uncorrected for the non-unit actlvity of liquid thorium.
described in a previous study,12 (orifice diameter = 0.139 em., length == 0.38 cni.) contained a thoria cup into which approximately 1 g. of thorium metal was added. The present authors have observed, as have Darnel1 and nIcCollum,7 that liquid thorium is difficult to contain in the effusion cell. After several hours at high temperature the liquid is able to "dissolve" through the thoria cup, wet the tungsten, and then creep out of the cell through the orifice. When the metal is first melted it does iiot wet the thoria as evidenced by its pronounced convex meniscus. However, after a few hours time the meniscus becomes concave as a result of the change in physical properties of the metal phase as it slo~vlydissolves thoria. It therefore is difficult to determine reliably the heat of sublimation for the process Th(1, activity
< 1)
+ ThOz(s) +2ThO(g)
(3)
in which the activity is both time and temperature dependent. However, if the measurements are coiifined to sufficiently short time intervals during which the meniscus of the metal phase always remains convex, measurements of limited success can be achieved. The results are given in Table 11-R. Samples of thorium metal that had been heated in a thoria cup at 237OOK. for 30 min. and had retained a convex meniscus upon cooling contained 0.395 and 0.380 weight % oxygen or an average of 2.8 mole % Tho, as determined by the method of S~miley.'~Several samples (of the order of 5 g.) of the fragments of the thoria cups that were used to contain the liquid thorium were heated in air a t 1000° in an attempt to measure deviation from ideal stoichiometry of the dioxide phase. Since no measurable weight increases ( > 2 X g.) were observed one concludes that the dioxide phase did not dissolve a significant amount of the liquid metal and that the stoichiometric composition was maintained throughout the effusion measurements. Since the vapor pressures of thorium metal6 and dioxide are sufficiently small a t 237OoK., the weight losses given in (14) JT. G. Smiley, And C h e m , 27, 1098 (1955)
2.5
3.5
3.6
3.7
3.8
3.9
4.0
4.1
Fig. 2.-Mass spectrometric measurement of the temperaturedependence of Tho + and Thoz+ in equilibrium with solid thorium dioxide (30 e.v. ionizing electrons.)
Table 11-I3 correspond essentially to the pressures of gaseous monoxide calculated froin eq. 1. If Raoult's law is assumed, the liquid thorium phase in eq. 3 has an activity of 0.97 based on the concentration of dissolved thoria. Therefore, the monoxide pressures given in Table 11-I3essentially correspond to a liquid metal with an activity 0.97 and a dioxide phase a t unit activity. The mass spectrometric examination of the vapor effusing from a tungsten cell containing thoria established that both ThOz(g) and ThO(g) are of comparable importance. The dependence of the respective intensities on the absolute temperature is shown in Fig. 2 for 30 e.v. ionizing electrons. For the purpose of plotting, the data of different runs were normalized a t 104/T = 3.750. The heats of sublimation 173.5 f: 2.2 and 159.5 .& 2.5 kcal. mole-l for ThO(g) and TliOz(g), respectively, were shown to be essentially independent of the ionizing electron voltage within experimental error for 25, 30, 35, and 50 e.v. electrons. Ion currents of T h + also were observed but were severely dependent on the electron energy and hence appeared to have resulted from fragmentation of the Tho2 and T h o by the ionizing electrons. Subsequent thermodynamic calculation will confirm this observation. There was no evidence of volatile tungsten oxides such as WO, WO2, WO3, or polymers thereof. The limit of detection of such species is approximately 1/100 the T h o + intensity. The failure to observe such gaseous reduction products is concordant with previous obs e r v a t i o n ~of~ the ~ lack of reduction of Tho, by tungsten. However, other investigators,2p16have reported the significant reduction of Tho2 in tungsten filaments to yield gaseous reduction products. The cause and (15) R. J. Ackermann and R. J. Thorn, .Irgonne National Laboratory Report, ANL-5824, January, 1958. (16) P. Schneider, J . Chem. P h y s . , 28, 675 (1958).
5.0
k°K I
2825
OK
I
I I 1 I 2825 O K (2745 I Tho
I
.
2725 O K l T h 0 2
I
I
I! i\ThO I I
I
I I
I
&OK
I
I
I
2615OKl I\
I
0.1
-
0.05 -
I I I
I I 2490°K I
\
IL
Tho 2525 O 1 Tho2
Tho2
0.02
-
extent of the observed reduction will be discussed in some detail later. Some evidence has been found for a slight deviation from stoichiometry of T h o z samples heated in vacuo. All of the residues from the evaporation studies acquired a medium gray color but rapidly reverted to the original white color when heated in air to 1200OK. Both materials had the fluorite cubic structure and lattice parameters identical within theo error of measurement, i.e., a. = 5.5973 =k 0.0003 A. Several 5-g. samples were heated in vacuo and analyzed by determining the weight gain when heated iii air a t 1200'K. Thoria which had been heated to 28OOOK. was found to have the composition Th01.gg8,whereas no measurable deviation from stoichiometry was observed in samples heated to temperatures below 2600°K. It was possible to demonstrate the effect of slight substoichiometry and the change of composition with temperature by the mass spectrometric observations shown in Fig. 3. The Tho2+ and T h o + ion currents were recorded simultaneously and continuously as the temperature of the cell was varied. The temperature of the cell could be changed from an initial value to within a few degrees of a final value in less than 1 min. In all cases the TliOz+ reached a constant value within 2 min. or essentially as soon as the temperature attained the final value. On the other hand the T h o + did not
come to equilibrium immediately but decreased toward a constant value when the temperature was lowered and increased when the final temperature was higher than the initial. The rate of approach to the equilibrium value was considerably more rapid a t higher temperatures and for smaller temperature increments. All of these observations appear to be correlated with the previously described composition change occurring in the sample which influences essentially only the gaseous monoxide. Since both species ultimately reached steady-state values, the solid phase must evaporate congruently, and in view of the very small subStoichiometry actually detected, it will be considered to evaporate stoichiometrically. Thermodynamics of the Thorium-Oxygen System Using the heat capacity data of Southard" and Hoch and Johnston's for ThOz(s), the heat capacity data of Stull and Sinke19for liquid thorium and diatomic oxygen, the heat of formationZ0ofThOz(s), AHOzss = -293.2 i 0.4 kcal. mole-l, and the absolute entropy,21Xom (17) J. C. Southard, J . A m . Chem. SOC.,63, 3142 (1941). (18) M. Hoch a n d H. L. Johnston, J . Phys. Chem., 66,1184 (1961). (19) D. R. Stull and G. C. Sinke, "Thermodynamic Properties of the Elements," American Chemical Society, Washington, D. C., 1956. (20) E. J. Huber, Jr., C. E. Holley, Jr., a n d E. H. Meierkord, J . A m . Chem. f l o c . , 74, 3406 (1952). (21) D. W. Osborne and E. I?. Westrum, Jr., J . Chem. Phys., 21, 1884 (1953).
THERMODYNAMIC STUDY OF THORIUM-OXYGEN SYSTEMS
April, 1963
= 15.59 e.u., one obtains for the standard free energy of formation of TliOz(s)the equation
AFf'(ThO2,s) = -2'96,000
+
46.387' (2000-3000°K.)
(4)
Based on the mass spectrometric identification of vapor species and the congruency of evaporation of the stoichiometric ThOz(s), there are two principal evaporation processes, namely Th(&(s) +Tho&)
(5)
and
+
ThOe(s:I + O(g) (6) The importance of each process can be ascertained in the following manner. The standard free energy change for reaction 3 can be evaluated from the pressure of ThO(g) at 2369OK. (Table 11-B) corrected for the assumed activity of liquid thorium, 0.97. When this quantity is combined with the standard free energy of formation of ThOz(s) (eq. 4) and the standard free energy of formation of atomic oxygen,l9 one obtains the standard free energy for eq. 6, AF02369= 164.9 kcal. Combining this value with the average heat of sublimation of ThO(g) corresponding to eq. 6 reported in Fig. 2 one obtains for reaction 6 the standard free energy 347,000 - 76.9T (7) Since it has been established that the dioxide phase evaporates essentially via reactions 5 and 6, it necessarily follows that the mole effusion rate of ThO(g) equals that of atomic oxygen. Therefore, po = (16,' 248.12)'/'p~ho, which fact when combined with eq. 7 yields the partial pressures of monoxide and atomic oxygen AF6'
log p T h O (atrn.)
=
8.70 - 3.79 X 104/T
(8)
and log p o (atm.)
8.10 - 3.79 X 104/T
(9) It can be easily shown that the "effective" pressure given by eq. 2 is related to the partial pressures of the dioxide and monoxide via the relationship = I)ThOt
=
+ (264.12/248.12)"'~~h0
(10) from which one can obtain the partial pressure of the gaseous dioxide pe
log pThOl (atm.)
=
7.64 - 3.44 X 104/T
(11) which is shown in Fig. 1. The heat of sublimation corresponding to rea,ction 5 is 157.4 f 2.5 kcal. mole-' and is in concordance with the mass spectrometric value of 159.5 f. 2.5 shown in Fig. 2, which fact illustrates the internal consistency of the mass effusion and the mass spectrometric measurements. It is interesting to note that the internal consistency demonstrated by the two kinds of measurements allows the evaluation of the ratio of the ionization efficiencies of ThOjg) and ThOz(g)of 3.5 to 1. By combining eq. 7 and eq. 4 with the standard free energy of formatioil of atomic oxygen,Ig one obtains for the standard free energy of formation of gaseous thorium monoxide the equation
767
AFfo(ThO,g)= - 10,300 14.47' (2000-300OOK.)
(12)
The combination of eq. 4 and 11 yields for the standard free energy of formation of gaseous thorium dioxide the equation Al"fO(Th02,g) = - 138,600
+
11.47' (2000-3000°K.)
(13)
The estimation of the dissociation energies of ThO(g) and ThO,(g) a t absolute zero can be accomplished by means of the molecular constants and thermodynamic functions given in Table 111. The respective values of the dissociation energy, 8.3 and 16.3 e.v., shown in Table 111-A represent upper limits since the electronic states of the molecules are unknown. For the case of ThO(g) the calculated values of the absolute entropy a t 2000 and 30OO0K. are 3 and 5 e.u., respectively, less than those evaluated from the measured entropy of formation and the absolute entropies of liquid thorium and atomic oxygen. This discrepancy between the calculated and measured entropies probably reflects to some extent the existence of unknown low-lying electronic states. The vibrational constants for ThO(g) mere estimated from the valence-bond model described by Herzberg.22
Discussion The over-all agreement between the mass effusion rates of thoria reported herein (eq. 2) and those reported by Darnel1 and McCollum7 [log p(atm.) = (8.16 f 0.47) - (35,500 i 1100)/T] is quite consistent within the precision of the respective measurements. The heat of formation and dissociation energy of ThO(g) given in Table I11 are in agreement with the values (-8.0 kcal. and 8.5 e.v., respectively) reported by Darnell, et aL6 The extent of reduction of Tho2 by tungsten has been a subject of considerable discussion. The presence of approximately 1% thoria in tungsten filaments has been shown to produce electron emission which is orders of magnitude greater than that of il, pure tungsten filament. The increased emission generally is ascribed to a layer of thorium atoms adsorbed on the surface of the tungsten filament, and this layer of thorium atoms is said to result from the reduction of ThGz to metallic thorium by the tungsten and the subsequent diffusion of the metal t o the surface,2a2 4 Shapiro2has reported relatively large losses of tungsten from filaments a t temperatures above 24OOOK. which cannot be attributed to the vapor pressure of pure tungsten. It was concluded that the loss of tungsten results from the reduction of ThO, by tungsten to produce volatile oxides of tungsten. Measurements of the evaporation rate of thoriumbearing species from thoriated tungsten filaments both carbonized and non-carbonized have been reported recently by Schneider,l8 who has demonstrated that heating at 215OOK. for 5 min. causes an appreciable fraction of the thoria content of the filament to be reduced to metal without producing appreciable evaporation of the metal. It is, however, significant to note (22) G. Herzberg, "Infrared and R a m a n Spectra of Polyatoinic .Molecules," D. Van Nostrand, New York, N. Y . , 1945. (23) I. Langmiur, Phys. Rev., 22, 357 (1923). (24) C. J. Smithells, J. Chem. Soc., 121, 2236 (1922).
R. J. AcKEK~~ASS, E. G. 1?AUH, R. J. THORN, BND AI. C. CANSON
768 A.
MOLECULAR CONSTAKTS AND
Property
TABLE I11 THERYODYKAMIC PROPERTIES O F GASEOUS THORIUM ThOz(d
ThO(g)
OXIDES
----References----Tho
Intera toniic distance, -4. Vibrational constant(s), em.-'
1.93 740
Nl.9 715,305(2), 763
Electronic ground stat e Heat of formation, cal. (2000-3000°K.) Entropy of formation, e .u. (2000-3000°K. j Dissociation energy (0°K.j, e.v.
("1
('2)
-10,300
6
ThOz
a
Estimated Estimated
- 138,600
Eq. 12
Bssumed Eq. 13
f14.4
-11.4
Eq. 12
Eq. 13
t-o,= 2.5 x 10-15 (4) ThOs(s) W(s) -L Th(g) M'0dg)' pTh I)\T02 = 1.15 x 10-" 8 . 2 x 10-9 5, 6, and eq. 4 5 . 3 x 10-6 5 and eq. 4 and 12 2P\T02 = 7 . 1 x lo-' (5) 2 T h 0 ? ( ~ ) W(S) + 2ThO(g) JVOp(g)' PThO ... c and eq. 4 p c o = 0.18 (6) ThOa(s) 2C(s) Th(1) 2CO(g) ( 7 ) Th(l) Th(g) PTh = 8 . 0 x 3 . 6 x 10-4 6 (8) ThOZ(9) Th(1) -+ 2ThO(g) PThO 2.6 x 2 . 0 x lo-? eq. 4 and 12 a The weight losses of thorium were calculated from eq. 1 and the assumption that the eflective evaporating area is the surface area multiplied by the fritction of Tho2 contained in the filament (0.0189). Even if this latter factor is omitted of the filament ( S4.7 J. P. Coughlin, Bur. Calculations involving WO(g) and WOa(g) yield similar conclusions. the conclusions reached are unaltered. Mines Bulletin KO.542, 1954.
+
- + + +
+
+
+
+
that all of the thoria could not be reduced, but oiily a fraction thereof. At approximately 2550'K. it is possible to evaporate all of the thorium metal (or reduction product) in less than 1 min. I n attempting to explain this behavior, it is instructive to compare the evaporation rates for noli-carbonized filaments with those calculated for possible reactions occurriiig in the filament. I n order to make these calculations it is necessary to recall certain features of the filaments employed by Schneider. Each filament weighed approximately 1.5 g. and initially contained 1.89% Tho2 by weight and 0.20% thorium by weight or 2.8 X g. of thorium. From an g. of Tho, and 3 X inspection of the experimental diode employed by Schrieider it is reasonable to assume that a t least half of the length of the filament attains the uniform reaction temperature. However, in the present discussion the calculations can be simplified if it is assumed that the entire filament (0.06 cm. diameter, 25.0 cm. long) attains the reaction temperature since a t most an error of a factor of two is involved, which is insignificant for the present purpose. ,4t 2550'K. (see Fig. 2 , ref. 16) about half of the initial amount of Thoz is reduced to
thorium metal or a volatile species all of which evaporates from the filament in less than 1 miii. The results of the calculations of various reductioii and evaporation processes are shown in Table IV. Clearly there is no reduction process involving tungsten metal that is capable of producing thorium metal or a volatile species containing thorium which can account for the extent of reduction and subsequent evaporation observed by Schneider. An explanation of Schiieider's observations, therefore, demands the presence of a reducing agent in the filament which is capable of reducing thoria to thorium metal a t temperatures greater than approximately 2000'K. Reaction 6 in Table IV demonstrates that carbon is quite capable of reducing thoria to thorium metal. Furthermore, it also has been demonstrated by Schlierz5that carbon is a common impurity in tungsten filaments. According to his measurements a large amount of carbon monoxide is produced when an emitting tungsten filament is heated in an oxygen atmosphere of the order of lo-' mm. Actually there are few, if any, reducing agents in addi( 2 5 ) R. E. Schlier, J . A p p l . Phys., 29, 1162 (1958).
April, 1963
ACIDITYOF SOLIDCATALYSTS
tion to carbon capable of reducing thoria a t high temperatures. It appears, therefore, to be a necessary conclusion that the filaments employed by Schneider contained a carbon impurity sufficient to reduce some of the thoria according to reaction 6 followed by the subsequent evaporation of gaseous monoxide via reaction 8. Only these two of the proposed reactions are capable of explaining quantitatively the extent of reduction and subsequent loss in weight observed by
7 69
Schneider16if the carbon content of the filaments was of the order 0.05-0.1% by weight. This value is in reasonable agreement with the findings of Schlier. * 5 Since it has been shown that the extent of reduction of thoria by tungsten is indeed small, the previous discussion suggests that a carbon impurity in tungsten filaments is implicated in the production of thorium metal which is necessary for the enhancement of the electron emission of thoriated filaments.
DETERR/IINATION OF ACIDITY OF SOLID CATALYSTS BY AMMONIA CHEMISORPTION BY
YUTllKA
KVBOKAWA~
Department of Applied Chemistry, University of Osaka Prefecture, Sakai, Osaka, J a p a n Received A u g u s t 37, 1962 The determination of the strength of acid sites on silira-alumina and alumina catalysts has been carIied out by measuring the rate of desorption of ammonia chemisorbed on the catalysts. The activation energy of desorption is increased with decreasing amount adsorbed from 10 to 50 kcal./mole on the silica-alumina, indicating the heterogeneous nature of the distribution of acid sites. The acid strength of the alumina is found to be comparable to that of the silica-alumina. The energy distribution of acid sites on silica-alumina deteimined by the desorption method is markedly different from that obtained by the indicator method proposed by Benesi. The reason for such discrepancy is discussed.
Introduction In view of the importaiice of acidic catalysts in the petroleum industry, a considerable amount of research on measurements of the acidity of solid catalysts has been carried out by various workers. The most general method seems to be the titration of solid catalysts in a noli-aqueous medium with amines as proposed by Tamele.2a Benesi2tl expanded this method to make it possible to determine the acid strength distribution using a complete set of available Hammett indicators. As pointed out by Benesi, such a method is based upon some assumptions and it also has a number of limitations in its actual application. For measuring surface acidity, the investigation of the chemisorption of a basic gas such as ammonia at elevated temperatures seems to be very promising. Most of these studies are a t present restricted to measurements of the amount adsorbed a t a particular temperature.3 If, however, the energy values of ammonia chemisorption were determined over a wide range of coverage, the information concerning the acid strength distribution would be much improved. At present, the only work along these lines seems to be that of Zettlemoyer, et al.,4 who have determined the acid strength distribution from heat of immersion measurements. In previous work5 it has been shown that measurement of the desorption rate can give the energy relation for chemisorption 011 oxide catalysts over a wide range of coverage. It therefore has been undertaken to in( 1 ) Department of Chemistry, The Johns Hopkins University, Baltimore 18, Md. (2)(a) 31. W.Tamele, Discussions Faraday Sue.. 8, 270 (1950); (b) H. A. Benesi, J . Am. Chem. Soc., 78, 5490 (1956): J . P h y s . Chem.. 61, 970 (1957). (3) G. A . Mills, E. R . Boedecker, and A. G. Oblad, J . Am. Chem. Sue., 74, 1564 (1950). (4)(a) J . J. Chessick and A. C. Zettlemoyer, J . Phys. Chem., 64, 1217 '. C. F. Holm, and D. M. Blackburn (1058); 64, 1131 (1960); (b) A . Clark, 1 have recently determined tlie energy values of ammonia chemisorption on silica-alumina catalysts from adsorption equilibrium measurements ( J . Catalvsis, 1, 244 (1962)). (5) Y. Kubokau-a, Bull. Chem. S U CJ. a p a n , 33, 546, 550, 5 5 3 , 739, 747, 036 (1960).
vestigate the rate of desorption of ammonia chemisorbed 011 solid acid catalysts such as silica-alumina and alumina, Experimental Materials.-The silica-alumina catalyst containing 13% alumina was obtained from the Shokubaikasei Co. The alumina used in this work was the material for chromatographic use manufactured by the Wako Junyaku Kogyo Co. The impregnation of the catalysts with sodium hydroxide was carried out as follows: a solution containing the desired amount of sodium hydroxide was added to the catalysts, dried at looo, and sintered at 400". The final catalysts contained 3 mmoles of NaOH/g. By a similar impregnation procedure, sulfuric acid was mounted on pure silica gel for chromatographic use obtained from Mallinclirodt Chemical Works. The acid concentration was 1 meq./g. Ammonia was obtained from the thermal decomposition of ammonium chloride and purified by fractional distillation. Prior to the chemisorption experiments, the silica-alum ina and alumin a catalysts were evacuated a t 500" for 12 hr., and the sulfuric acid mounted on silica gel a t 250" for 8 hr. The surface areas determined by B E T method using nitrogen adsorption were 448 m.2/g. for the silica-alumina and 190 m.2/g. for the alumina catalysts. Apparatus and Procedure.-The amount adsorbed was determined by using a conventional constant volume apparatus. The method for measuring the rate of desorption has been described in the previous paper,5 and will be repeated here only in outline. The chemisorbed gas was desorbed by malting use of a mercury diffusion pump and the desorbed gas was collected in a McLeod gage whose pressure was follofi-ed at definite intervals. It was confirmed that the oflserved rate of desorption is unaffected by the reverse process, i.e., readsorption. The activation energy of desorption was determined as follows. The temperature was lowered abruptly during the desorDtion experiment and the rates before the temperature drop were extrapolated to those for the smaller amounts adsorbed when the measurements were carried out after the temperature drop. Thus, the rates at the two temperatures corresponded to the same amount adsorbed and, the activation energy of desorption could be obtained.
Results and Discussion Acid Strength Distribution by Desorption Method.After ammonia was allowed to adsorb at about 250°, the temperature of the specimen was raised from - 50" to 450" in stages, at each of which the activation enerry of desorption was determined in the manlier given
