June, 1057
NOTES
829
atom. I n the case of homonuclear molecules C = 2,’. Table I1 lists for the hydrides and homonuclear TABLE I molecules of Table I, c and c’’~, respectively. For IIY1)RII)WS AND HOMONUCLEAR DIATOMIC MOLECULES five elements both values are available and are gen(All quantities are cxpresscd in atomic units) erally comparable except for IC. The fourth column hloleriilc: lists as 2, either c or c’/g, or the average where both tryare available. Inspection reveals that these values dridps ke He D. a b C for effective charge show a periodicity and lie in 1.067 4.178 0.120 2.110 0,1342 0 2881 f’t1 5.652 7.235 ,1083 1.263 ,5006 1.833 011 general between the core charge as a lower limit and 1.044 7.349 12.653 2.407 .l697 IICI ,332 the atomic number as an upper limit. A more quan2.077 6,088 ,0084 0 . 5 2 8 0 4 24 KIf ,0361 titative comparison is shown in the last two columns ,0349 1.369 1 1 . 0 5 6 26.812 ,0971 3.013 Znll in which are given 2, and Ze/Za where 2, is the ef2.670 ,1440 1.030 8.454 16.554 l[l\r .2Gl ,0281 1.386 15.960 43.685 ,0774 8 33 (:,Ill fective charge acting on the valence electrons as ,1176 1.021 10,657 24.434 ,2018 8.033 IJI calculated by Slater’s rules. With the exception of 1.799 4 3 . 5 4 3 122.555 3.29 .0169 Ilgl[ .07:3P Zn, Cd and Hg the ratio Ze/Z, shows a general 1 Iornonuclesr trend as the atomic number increases. 0.1743 0,9049 1.402 1.0972 0.3GR 1.40 112 Discussion ,0386 0 , 4 8 3 2 0 , 0 1 8 4 5.05 1.524 5.459 I,ia 1.587 41.413 65.801 ,3640 1.4788 2 . 0 8 7 Nz Although the generalized semi-empirical equa61.111 l.fl02 34.143 2.280 ,1904 0% 0.758 tion 2 is not as well-founded on theoretical prin5 . 8 2 ,0272 0 . 4 8 7 8 1.785 7 . 6 8 2 N R ~ ,01103 ciples as is equation 1, which is restricted to hydro5 4 . 1 4 3 355.45 ,1867 1,131 ,357 8.58 P2 1.270 6 3 . 2 5 7 196.67 8.76 ,0923 ,2113 CI? gen, the present results show that the interpretaKz .OOG~Z 7.41 .oir112 0 , 4 5 5 6 2.448 13.094 tion of the parameter c as a product of effective 1 . 2 5 4 105.81 3 8 5 . 8 6 ,0732 ,1570 4 . 3 2 Hrr nuclear charges has some semblance of truth. It is .0572 1.207 177.47 768.17 ,1107 5.04 In particularly gratifying that the pairs of 2, values Test of the Function. Having determined the for 0, GI, Br and I agree so well. Of course no such simple potential energy funcparameters it would be desirable to make predictions of properties other than those used in deter- tion as equation 2 would be expected to be univermining the parameters. Third and fourth deriva- sally valid. In particular ionic states would retives a t the pot,ential energy minima have been quire special treatment and can be handled by an cnlculat,ed and compared with those obtained from extension of the methods of Paper I. Furthermore, roscopic constants cye and W e X e . Although since Pitsere has found that London forces may the calaiilatfed values are the right order of mag- contribute considerably to bond energies equation 2 nitudc (oft,cn within 20%) the results are no better should probably be modified to give an inversc tban obtainablc in the same way with the Morse sixth power law a t large distances. fiiiiction or witb the more recent Lippincott4 func(0) K. 5. Pitzar, J . Chem. Phys., 2 3 , 1735 (19.56). tion, or wit,h t8hareduced potential energy function.6 Therefore thcse calculations will be omitted here. More int,erestingis t,he test of the para.meter c as a THE HEAT OF FORMATION O F SILANE product of effective nuclear charges (eq. 3). For hydrides, sinw the H atom mould presumably have B Y E. 0. BRIMM AND HARRIE M. HUMPHREYS Z = 1, C = Ze, the effective charge of the other berg’s3tables for which accuratc values of all three experimental constants were known.
TABLE I1 EFFECTIVE ATOMICNUMBERS RELATED TO THE PARAMETER c Element
H Li C N 0 Na P CI IC XI1
Br Cd I Hg
Hydride
llornonuclear
C
c’/t
1.04 2.33
0.12 7.23 12.66 6.08 26.81
8.09 7.81 2.77 12.47 14.02 3.74
16.55
19.64
43.09 24.43 122.6
27.72
Effective Atornio No. zs ZI
1.04 2.33 6.12 8.09 7.52 2.77 12.47 13.34
1.00 1.30 3.25 3.90 4.55 2.20 4.80
4.91
2.20 4.05 7.00 4.35 7.GO 4.35
20.81 18.10 43.09 26.08 122.6
6.10
Ratio ZdZI
1.04 1.80 1.88 2.07 1.66 1.26 2,GO 2.19 2.24 0.16
2.38 10.1 3.44 28.1
(3) G . Herzberg, “Spectra of Diatomic hfolecules.” 2nd Edition, D . Van Noatrand Co., Inc., New York, N . Y . , 1960. (4) E. R . Lippincott, J . Chem. P h y a . , 21, 2070 (1953): E. R. Lipplncott and R . Schrocder, (bid.. 23, 1131 (1955). ( 5 ) A . A . Froet and B. hIusulin, J . A m . Chem. Soc.. 76, 2045 (1954).
Cnnfribulion from,the Ressarch Laboraloru of L i n d o A i r Produrlr Company, a Division of Union Carbide and Carbon Corporalion. l‘onnWanda, N e w York Received J a n u a r y P9, 1967
The heat of formation of silane ga,s from crystalline silicon and hydrogen a t 25” and one atmosphere pressure mas measured. by a direct, calorimetric method. The value obtained was +7.8 kcal./inole wit3han average deviation of k 3 . 5 kcal./niole. The values for the heat of formation of silane given in the vary from -8.7 to - 14.8 kcal./mole. I n addition, these values are not coiisistent with certain other observed proper~ prescnt investigation was ties of ~ i l a n e . ~ fThe undertaken in order t o providc a n a,dditional,independent determination of tJhequantity and, if possible, to reconcile the esist,ing inconsistencies. (1) H. Van Wnrtenberg, Z. a n o r g . nllgem. Chem.. 79, 71 (1913). (2) F. R . Bichowsky and F. D. Rosnini, “The TherinochemiRtry of Chemical Substances,” Reinhold Publ. Corp., New York, N. Y., 1936, p. 253. (3) U. 8. National Bureau of Standards Circ. No. 500, “Ralected Values of Chemioal Thermodynamic Proporties,” Washington, 1952. (4) T . R. Hoencss. T. L. Wilson and W ,C. Johnson, J. A m . Chem. Soc., 58, 108 (1936). (.5) -1. C . Rrmtleu and T. Suint. rinpuhliahed d n b .
NOTES
830
Previous determinations of the heat of formation have been based on indirect measurement,, either the heat of combustion of silane or the temperature coefficient of the equilibrium constant for the reaction Si(s)
+ 2 H d g ) z SiHdg)
(1)
The first method depends upon the heat of formation of hydrated silica, which is not accurately known.6 The second depends upon the establishment of a true equilibrium, and the existence of such an equilibrium is questionable.2 Also, the values obtained for the heat of formation of silane, namely, -8.7l, - 13.72and - 14.B33kcal./mole are not consistent with the following observed properties of this compound. Hogness, Wilson and Johnson (ref. 4) allowed finely divided silicon and hydrogen to remain in contact for several days a t temperatures up t o 480" and a total pressure of one atmosphere. No pressure drop was observed; no silane coiild be detected. Measurements made in this Laboratory6show that, when silane and excess hydrogen are mixed a t 425" and a pressure of 100 atmospheres, complete decomposition of the silane takes place. However, if the values for the heat of formation listed above are combined with the free energy functions (obtained from heat capacity or spectrographic measurements) for silicon, hydrogen and calculated values of the equilibrium constant for reaction 1 require large equilibrium mole percentages of silane under the conditions used by Hogness, Wilson and Johnson or by Brantley aiid Smist. I n view of these contradictory resultJs, a direct determination of the heat of formation of silane seemed necessary. The heat of formation was measured by a standard calorimetric procedure,1° purified silane being decomposed to crystalline silicon and hydrogen at 680". The calorimeter was designed to operate a t 550" but a temperature of 680" proved necessary to ensure complete decomposition of the silane. At 680" there was considerable thermal lag in the calorimeter system. Also the ratio of electrical energy t o silane decomposition energy was 60:l in contrast t o 35: 1 at 550". These facts probably account for the large average deviation in the measurements. The value obtained, based on five calibration runs and four runs in which silane was decomposed, was +7.8 f 3.5 kcal./mole. The heat of formation is positive in contrast to the negative values reported previously. Because of the limited data and large average deviation in the calorimetric experiment,s, the sign of the heat of formation was checked by another method. A 1-inch quartz tube was fitted with grtR inlet and outlet tubes, a gas flow meter in the inlet sido, and a thermocouple in the center of the quartz tube about 3 cm. downstream from the inlet tube. This assembly was placcd in a tube (6) T. L. Cottrell, "Thc Strength of Chemical Bonds," Academic Preas Inc., New York, N. Y., 1954, p. 241. (7) Calculated from data given by K. K. Kelley, Bur. Mines Bull., 371 (1934). (8) N.B.S.-N.A.C.A. Tables of Thermal Properties of Gases, Tables 7.10 and 7.11. (9) (a) C. Cerny and E. Erdos, Chem. Lis!!/. 47, 1743 (1953); (b) A. P.Altschulcr, J . Chcm. Phys., 93,:IF1 (1n55). (IO) &e, for cxarnplr, W. 11. Johnnon, It. C:. Millcr u~rd E. J. Pruwrl, N.U.S. Llc1,ort Nu. 3267, .lilituary 15, L O M .
Vol. 61
furnace and heated to 600'. The system was flushed with hydrogen and was then allowed to come to temperature equilibrium with no gas flowing throu h the system. H drogen was then allowed to flow througf the system and tXk thermocouple registered a large temperature dro (about 10'). The thermocou le returned to its originare uilibrium value when the gydrogon flow was stopped. $Text, silane was allowed to flow at the same rate, and the thermocouple registered an increase in temperature (about 2"). The system returned to equilibrium when the silane flow was stopped. This procedure was repeated several times with the same results. The temperature of the furnace was then set at 320' and the same ex eriments performed. At this temperature, which is well gelow the temperature a t which silane decomposes rapidly, a flow of either silane or hydrogen resulted in a decrease in the temperature of the thermocouple (about 4" for silane, 10' for hydro en). The temperature increase observed during the silane #ow a t 000' must then result from the exothermic decomposition of this gas.11
Calculated values of the equilibrium constant for reaction (1) based on AHf" = +7.8 kcal./mole and the data from references 7, 8 and 9 are: K p = 1.9 X a t 25' and IC, = 1.4 X lo-' a t 450'. The amount of silane present a t equilibrium a t 450" is therefore extremely small, about 1.4 X mole % in the gas phase a t one atmosphere mole % ' a t a prespressure and about 1.4 X sure a t 100 atmospheres. This explains the failure of Hogness, Wilson and Johnson t o form silane and the essentially complete decomposition of silane under the conditions employed by Brantley and Smist. The stability of silane at room temperature must be attributed t o a very slow rate of decomposition rather than to a favorable equilibrium. Acknowledgment. The authors wish to thank Dr. J. C. Brantley for his assistance in carrying out the experimental work and Dr. E. J. Prosen, chief of the Thermochemistry Section of the National Bureau of Standards, for helpful discussions on the calorimetric method. (11) Mr. E. G. Caswell of this Laboratory oarriod out these oonfirming experiments.
THE METHYL IMINO RADICAL' BY FRANCIS OWENRICEA N D CHEBTER J. GRELECKI The Cotholic Univsraity o j America, Washington, Received February 4, 1967
D. C.
The thermal decomposition of methyl azide has been studied2and found t o be a homogeneous, unimolecular reaction. Leermakers3 reported the results of a study of the reaction in a static system at temperatures between 200-240" and pressures from 0.078 $0 46.6 cm., and suggested that the methylimino radical forms, and then undergoes a rearrangement to methyleneimine. Our experiments were designed t o stabilize the methylimino radical or other active species by freezing the products of the thermal decomposition in a high speed flow ~ystem.~ (1) This work was supported in part by the United States Atomic Energy Commission, contract No. AT-(40-1)-1305. (3) 11. C. Ramsperger, J. Am. Chem. Soc., 61, 2134 (1929). (3) J. A. Lcormnkcrs, ibdd., 05, 3008 (1933). (4) For a dencriiption of thc tecliniquo sea F. 0. Rico and M. Froatno, ibid.,73, 6529 (1961).
