8.
660
phase. The concentration of oxygen in solid silver has been investigated by Sieverts and Hagenacker7 and Steacie and Johnson.8 These workers confirm that Sieverts' law holds for this system, Le., that c =
.
Kd/P
VOl. 56
CHU-LIANG
(9)
Differentiating (9) and substituting values of c and dc into equation (8) gives for the surface in the oxygen-silver system
The change in surface tension with the logarithm of oxygen partial pressure can be obtained from Fig. 6 and is -188. The surface excess a t 930' is then I' = 1.98 X 10" atoms oxygon/cm.*
This figure can be more easily visualized by coliverting it to an excess oxygen surface concentration in atoms of oxygen per atom of silver. Assuming a close packed arrangement on the surface, a (111) face of silver will have 1.4 atoms of oxygen adsorbed per silver atom. A (110) face will have 1.65 atoms oxygen per silver atom. These values are uncertain by 15%. The oxygen is undoubtedly chemisorbed to the silver surface. This conclusion is supported by the fact that the sorption is not, withiti the limits of the precision of the data, affected by a change in oxygen partial pressure. While the adsorption isotherm gives the amount of adsorption, i t says nothing about the form of the (7) A. Sieverts and J. Hagenacker, Phys. Chem., 68, 125 (1909-lo). (8) E. W. R. Steacie and F. M. C. Johnson, Proa. Roy. SOC. ( L o n d o n ) , A l i P , 429 (1926).
layer. However, it is certain that the chemisorbed layer does not have the thermodynamic properties of bulk silver oxide since bulk AgZO could only exist at the test temperature at an oxygen pressure of 1.7 X lo4 atmospheres. One may speculate about the decomposition pressure of the surface film by observing that the oxygen partial pressure a t which the presence of the oxygen becomes ineffective on the surface tension. Extrapolation of the line in Fig. 6 shows that in the atmospheres of oxygen the surface vicinity of appears tension would become 1140, so that to be a decomposition pressure of the surface film at932'. There have been no previous studies of adsorption of oxygen on silver at high temperatures reported in the.literature. The results of this study are, however, in general qualitative with the results of Benton and Elging who found that oxygen is chemisorbed on silver in the range 0 to 200'. Armbruster'O also studied the sorption of oxygen on silver and concluded that a t 20' about 0.3 X l O I 5 molecules of oxygen adsorbed to silver per square centimeter and that adsorption increases with temperature. Thus the work of Armbruster provides an order of magnitude confirmation of the work presented in this paper. Acknowledgment.-The authors are indebted to the Office of Naval Research for funds made available under contract number N5ori-07841 for equipment and supplies used in this research. We are also indebted to Prof. J. Wulff for his guidance and counsel. (9) A. I". Beoton arid J. C. Elgiri, J . Am. Chem. Soc., 61, 7 (1029). (LO) M . H. Armbruster, ibid., 64, 2545 (1042).
LOW VAPOll PRESSURE MEASUREMENT AND THEllMAL TRANSPIRATION BY S. CHULIANG Division of Chemistru, National Research Council, Ottawa, Canclda Received July 10, 18.51
The thermomolecular pressure ratio, the R-factor, of ethylene has been measured and found that it can be fitted to the nitrogen values with the ressure shifting factor f = 0.5. The vapor pressure of ethylene a t 77.5 and 90 "K. has been measured. The results of Tic&ner and Lossing on methane, ethylene and acetylene have been corrected for the R-factor. The B/T. Thc constants A and B , and the applicable temperacorrected values can be represented by an e uation log p = A ture range for each gas are given. A metho1 is suggested by which!, and therefore the R, factor of one gas can be estimated from that of another.
-
Recently, Tickner and Lossing re-examined the vapor pressure data of hydrocarbons by means of a mass spectrometer for the purpose of providing some reliable values.'e2 Their effort will be defeated if the results so obtained are not corrected for the thermal transpiration effect. These authors being aware of such correction, used a connecting tube of large diameter (16 rnm. i.d.), and contended that only the pressures below 0.01 mm. were uncertain. While it is so for some of the larger molecules, it is not so in the cases of methane and '
(1) A. W. Tickner and F. P. Losaing, J . Cham. Phya., 18, 148 (1950). (2) A . 'w. Tickner and P.P.I,o~eing,Trrr'd JoonNAr.. 66, 733 (1951).
acetylene. The present paper will treat the cases of these two gases and ethylene as examples and outline a method by which such a correction may be estimated without actually determining the effect for individual gases. Experimental Part Experimental work was carried out for ethylene to determine its vapor pressures at 77.5 and 9O"K., and also to measure its thermomolecular pressure ratio R as a function of pressure p . The apparatus used was that of the relative method described elsewhere.8 The small tubing was a 1.6mm. capillary and the large tubing was 32 mm. inside diameter. The etliylenc used was withdrawn from a com' ( 3 ) 8.
C. Liang.
J . A p p l . P h ~ n . 22, . 148 (19511,
May, 1952
LOW VAP6R P R E S S U R E
MEASUREMENT AND THERMAL 'TRANSPIRATION
GG1
mercial tank and rurified by repeated operations of freezing the ethylene at 78 K., pumping off the residual gas and then revaporizing the solid. The finally purified gas was found t? be better than 99.9% pure by mass spectrometric analy818.
The R-value.-In the work with non-condensable gases, it was found that the R us. pol curves of different gases could be reduced to a single curve by introducing a "pressure shifting factor" f, and the f-values for several gases were given by arbitrarily choosing fN, = 1. The present determination revealed that Rc2=( fitted into this scheme very well withfc," = 0.5. In other words, with the cold and warm ends at the same temperature differentials and tubing being of the same size, R N ,at p is the same as R c , ~ at , 0.5 p. This is obviously because ethylene has larger cross-sectional area than nitrogen. The exact functional dependence of f on r, the radius of the gas molecule, is not. known. However, in several instances we have found that the empirical relation
holds quite well. Table I gives the comparison of experimental values and that calculated from eq. ( 1 ) with r values compiled by Hirschfelder, Bird and Spotz4 from the viscosity data.
TABLE I COMPARISON
O F f-FACTORS
Experimental"
I
NZ
I1
1 1.25 2.2 2.80
Calcd. by eq. (1)
1 A 1.30 1.3 H2 1.80 2.0 He 2.12 2.4 Ne .. 2.0 2.0 the experimental values were taken from Tahle VII, ref. 3, with neon value the unpublished work of the present author. 1
4.5
,
+
x
10
%
I
I
15
20
x103.
Fig. 1.-Vapor pressures of methane, ethane and ethylene: X, Delaplace; Bourbo; A, Liang; 0,Tickner and Lossing; ,. corrected T-L data.
+,
transpiration effect,are plotted in Fig. 1to compare with the uncorrected values and the values of other investigations. As the corrected values are believed to be the most reliable, for reasons to follow, and they can be represented by a straight line relationship B log P mm. = A - - (2) 1'
Considering the accuracy of the r-values, the agreements between the experimental and calculated values are excellent indeed with the exception of helium. With the aid of eq. (I), it is thus possible to estimate the R-factor for a given gas if its collisional diameter can be estimated from some other sources. The Vapor Pressure.-The vapor pressure of ethylene has not before been measured experimentally a t as low as 77.5"IC. the constants A and B for all three gases are given To serve as a double check, the pressure was measured si- in Table 111. multaneously with two McLeod gages connected to the cold-bath by tubing of 1.6 and 32 mm. i.d., respectively. TABLE I11 The 90°K. value was measured in the same manner. The CONSTANTS FOR VAPOR PRESSURE AND APPLIC.4BLE TEMreadings from the two gages, as may be expected, are difPERATURE ferent; but become the same when corrected for R. I n Methane Ethylene Acetylene Tahle I1 are summarized the experimental results. . A 7.655 9.526 9.630 TABLE I1 B 515 998 1250 VAPORPRESSURE OF ETHYLENE Highest applicGAGEI ( d = 1 . 6 mm.) able T,"K. 90 115 140 Experimental reading, mm.
RFactor
Corrected pressure, mm.
Discussion The results of Tickner and Lossing on methane, ethylene and acetylene, corrected for the thermal
To use .'q. (2) and Table IV, attention is called to the fact that they are only applicable up to a certain temperature and the highest applicable temperature for each gas is given in Table I11 as well. Methane.-The vapor pressure of methane seems to be one of the least investigated. The data from tJheBureau of Standards6extend only down to 10 mm., below which Tickner and Lossing's work appears to be the only one. Equation (2) agrees with the Bureau of Standards value up to about 100 mm., above which the straight line takes a new slope. Ethylene.-Within the temperature range where eq. (2) is applicable, the only other work is due to
(4) J. 0.Hirschfelder, R. B. Bird and E. L. Syotz, J . Chem. Phys., 16, 975 (1948).
(5) National Bureau of Standards, A. P. I. Project 44, "Selected Values of Properties of Hydrocarbons." (1848).
77.5%. 00 O K , 77.5"IC. 90 "IC.
1.04 X l o F 3 0.51 (0.53 f 0.01) X 10-3 0.036 0.82 0.030 GAGEI1 ( d = 32 mm.) 0.77 X 0 . 7 2 (0.55 f 0.03) X 10-8 0.030 1.00 0.030
The importance of thermal transpiration is clearly iilustrated. I t is of interest to notice that the vapor pressure of ethylene at liquid nitrogen temperature has alwa s been designated as 0.73 X 10-9 mm. The discussion ottained by correlating experimental data with this wrong value should therefore all be re-examined.
662
C. S. ROHRER, R. C. CHRISTENA AND 0. W. BROWN
Delaplace who calculated the vapor pressure by thermal conductivity.6 His data are also shown in Fig. 1. His values appear to be much too low as the temperature lowers. Inasmuch as his method is an indirect one, we incline to ignore his data. Acetylene.-The vapor pressure of acetylene has attracted unusual attention. Aside from the data of Tickner and Lossing, Delaplacee and Bourbo' supply another two sets. The results of Delaplace again appear to be too low. The comparison between the results of Bourbo and the corrected values of Tickner and Lossing is of interest. The values plotted in Fig. 1 are taken from the table given by Bourbo. There is a small but definite difference between the T.-L. and Bourbo results. In terms of pressure, the two sets differ by a factor of 1.4 in the higher pressure region and 1.8 in the lower pressure region. I n terms of temperature, the two sets differ by 2 degrees a t all pressures. The temperatures recorded by T.-L. are accurate to f0.5' ; if the Bourbo temperature scale is off by about, lo, (6) R. Delaplace, Comp. Tend., 204, 493 (1937). (7) P. S. Bourbo, J . Phys. (U.S.S.R.), 7, 286 (1943).
Vol. 56
then the difference between the two can be accounted for by other experimental errors. Since we know the accuracy of T.-L. measurements, we choose their values as standard. While this is rather arbitrary, there is a somewhat side line evidence. Bourbo's data in such low pressure ranges fit exactly into the extrapolation of Stull's valuess which extend down to 10 mm. only. This is not the general behavior of other hydrocarbons which all exhibit a change of slope with temperature, The results of Tickner and Lossing, when corrected for R-effect, follow this general rule. Other Gases.-The data of Tickner and Lossing on other gases, including carbon dioxide, may all need small corrections. However, owing to the low accuracy in their temperature measurements (3=0.S0),the corrections are about the same as their experimental errors. Acknowledgment.-The author is indebted to Drs. Tickner and Lossing for much of the detailed information about the vapor pressure measurements. (8) D. Stull, I n d . Eng. Clrem., SQ, 517 (1947).
T H E CATALYTIC ACTIVITY OF SOME REDUCED VANADATE SALTS BY C. S. ROHRER, R. C. CHRISTENA AND 0. W. BROWN Contribution No. 657from the Department of Chemistry,Indiana University, Bloomington, Indiana Received July I O , I951
The o timum operating conditions for each of the catalysts are summarized in Table VI11 for the reduction of nitronaphthakne to 1-naphthylamine. Of the catalysts studied copper gave the highest yields of the amine. Reduced nickel vanadate is too active for reduction to the amine giving appreciable amounts of the over reduction product, naphthalene. Small amounts of 1,2,5,6-dibenzophenazinewere produced by the use of reduced copper. vanadate and lead vanadate cata1.ysts.
Many studies have been carried out on the properties of reduction catalysts, particularly on the metals prepared by the reduction of their oxides. These metals have been classified into various systems, in attempts to predict the best catalysts for hitherto unexplored systems. Many other investigations have been made on addition agents and their effects in modifying the activity or selectivity of these metal catalysts. On the other hand little has been reported as to the manner in which these catalysts may be altered by reduction from their salts. Studies are now being made on t,he classified effects of acid radicals as moderators or promoters on these reduced metal or reduced compound catalysts. Reports of the catalytic activity of copper chromate, nickel tungstate2 and cobalt molybdates in the reduction of nitro compounds t o their respective amines have been published. This paper will deal with a systematic study of the variables concerned in obtaining the maximum activity of reduced nickel, copper and lead vanadates for the gas phase reduction of 1nitronaphthalene to 1-naphthylamine. These cations were chosen as representatives from each of three groups, nickel being of extreme activity in the (1) H.A. Doyal and 0. W. Brown, THIBJOURNAL, 36, 1549 (1932). (2) C. 8. Rchrer, J. Rooley and 0. W. Brown, ibid., 55, 211 (1951). (3) F. A. Griffitta and 0. W. Brown, ibid., 42, 107 (1938).
reduction of nitrobenzene, even giving reduction within the benzene ring; copper, from those metals of moderate reducing activity; lead, of a group with considerably less activity. Apparatus.-The apparatus was the type usually employed in such shdies. A horizontal type electrically heated, metal block furnace containing a reaction tube of . 1/2-inch iron pipe was employed. The furnace was thermostatically controlled, and temperatures were measured in the block and in the 10-inch catalyst bed at point,s one inch from each end. The feed was through a capillary tube with the rate regulated by a variable head of mercury. An oven thermostatically controlled a t 77" was placed at the head of the furnace to contain the 1-nitronaphthalene feed tube, in order that the reagent be maintained as a liquid with R nearly constant viscosity. Hydrogen flow was measured in liters per hour through a calibrated flow meter. Preparation of the Catalysts.-Copper oxide was prepared by themethodof Brown and Henke.4 The salts were prepared by adding solutions of copper sulfate pentahydrate, nickel acetate tetrahydrate and lead acetate trihydrate to hot solutions of ammonium metavanadate. Vanadium pentoxide was prepared by adding glacial acetic acid to a boiling solution of ammonium metavanadate. Each of the above precipitated compounds was washed repeatedly with distilled water except that the nickel vanadate, which is appreciably soluble, was washed with ethyl alcohol. The compounds were then dried overnight at 145'. Ten grams of a compound prepared as above, was evenly distribded in a ten-inch bed in the cold catalyst furnace and reduced under a stream of 14.5 liters of hydrogen per hour. (4) 0. W. Brown and C. 0. Henke, $bid., 26, 715 (1922): F. A. Madenwald, C. 0. Henke and 0. W. Brown, ibid., 31, 862 (1927).
