COMMUNICATIONS TO THE EDITOR
597
just this type of behavior in the Ce-Cd system (although it should be remarked that their interpretation of their results is quite different). DEPARTMENT OF CHEMISTRY ARIZONA STATEUNIVERSITY TEMPE, ARIZONA RECEIVED NOVEMBER 23, 1965
M. O'KEEFFE
0.1
:
:
The States of Nitrogen Adsorbed on an
d
Ammonia Synthetic Iron Catalyst and the Reactivity
Sir: For understanding heterogeneous catalysis, the information of the activity of chemical species in various adsorbed states is valuable since such an adsorbed species is expected to take part in the reaction as an intermediate. Recently, Emmett and the present author' found that the state of the chemisorbed nitrogen on an ammonia synthetic catalyst changes with the temperature a t which chemisorption took place. It is widely accepted that the ammonia synthesis on this catalyst proceeds through the steps of the nitrogen chemisorption and hydrogenation of the chemisorbed nitrogen. 2, If the nitrogen chemisorbed in different states is brought in contact with hydrogen, the rate of ammonia formation by the latter step is expected to be different at a given experimental condition. An iron synthetic ammonia catalyst (2.6 g, 2.03% alumina, 0.81% potassium oxide, 0.16% silica as promoters) was reduced in a stream of hydrogen at a flow rate of 500 cc (STP)/min a t 450 and 600" for 45 and 40 hr, respectively. Between the runs, the catalyst was reduced a t 450 and 600" for 2 and 16 hr, respectively. At the end of the reduction, the catalyst mm. Puriwas evacuated at 600" for 3 hr to 2 X fied nitrogen was adsorbed a t temperatures ranging from 208 to 445" and a t pressures of 20 to 120 mm. After the nitrogen adsorption, the temperature of the catalyst was decreased to a reaction temperature and nitrogen in the gaseous phase was removed by a Toepler pump. The amount of nitrogen adsorbed during the process of decreasing the temperature was less than 0.15 cc (STP). Thereafter, purified hydrogen a t a flow rate of 450 cc (STP)/min was passed over the catalyst on which nitrogen had been chemisorbed. The amount of ammonia thus formed was successively followed by absorption in a sulfuric acid solution. The rate of ammonia formation gradually decreases with time due to a decrease in the amount of adsorbed
0.5
1.5
1.0
Adsorbed amount of nitrogen, cc (STP).
Figure 1. Plots of the rate of ammonia formation us. the amount of nitrogen adsorbed: reaction temperature, 162' ( X ) and 208' (0); adsorption temperature is shown in the figure.
"
V
'8 rd
0.5
-1.24
9 0.973
200
300 400 Temperature of adsorption, "C.
Figure 2. Plots of the rate of ammonia formation a t a given amount of nitrogen adsorbed us. temperature of adsorption: reaction temperature, 162' ( X ) and 208" (0). Numbers designate the amount of nitrogen adsorbed.
nitrogen. The initial rates a t several temperatures for the nitrogen adsorbed a t different temperatures are shown in Figure 1 as a function of the amount of adsorbed nitrogen. The plots of the rate at the given temperatures and given amount of nitrogen adsorbed against the temperature of nitrogen adsorption are shown in Figure 2. As seen from Figure 2, it is clear that the nitrogen adsorbed a t lower temperature has ~~
(1) P. H. Emmett and N. Takesawa, to be published. (2) W. G. Frankenburg, "Catalysis," Vol. 111, P. H. Emmett, Ed., Reinhold Publishing Corp., New York, N. Y . , 1955, p 171. (3) C. Bokhoven, C. van Heerden, R. Westrik, and P. Zwietering, ref 2, p 265.
Volume YO, Number 8 February 1966
COMMUNICATIONS TO THE EDITOR
598
higher activity for ammonia formation than that at higher temperature. This suggests that the state of adsorbed nitrogen is varied by the temperature a t which the adsorption was made.
Acknowledgment. The author wishes to express his sincere thanks to Professor P. H. Emmett of Johns Hopkins University and Dr. I. Toyoshima of Hokkaido University for their profound interest and valuable discussions on the present work. RESEARCH INSTITUTE FOR CATALYSISNOBUTSUNB TAKEZAWA HOKJUIDOUNIVERSITY SAPPORO, JAPAN RECEIVED NOVEMBER 23, 1965
19.5 22.6
0
19.3
7 22.3 E!
0 a
I
3
19.1
X
gI
; 22.1 1 1
-x”I
3
b
18.9 21.9
Charge-Transfer Spectra in Nonpolar Solvents 18.7
Sir: The problem of the effects of solvent interactions on charge transfer (ct) spectra is an extremely complex one and various aspects of it have been studied in the past.’ From these it appears that existing theories which are based on the Onsager model2 of the solute point dipole embedded in the dielectric continuum of the solvent do not account for the observed facts. Even for pure compounds the electronic spectra were found to deviate from the predicted linear relation between solvent refractive index function and band shift in the region of low refractive index.3 One reason for the confused situation of the effects of solvent on ct spectra has been the total lack of gas phase ct data. Recently, a few such measurements were publi~hed.*~5They allow us now to test existing solvent theories as to their applicability to ct spectra. We wish to report here preliminary results of a comprehensive study on solvent interactions in ct complexes which will be published in full somewhat later.6 We have measured the ct energies of complexes of five nonpolar aromatic donors with tetracyanoethylene (TCNE) as acceptor in a set of six selected nonpolar solvents (three perfluoro and three hydrocarbon solvents with a refractive index range between 1.26 and 1.47 and which are not expected to show specific interactions). The results are shown in Table I. The data were plotted in terms of a refractive index function, (n2- 1 ) / ( 2 n 2 l), as derived in McRae’s solvent perturbation theory.’ Examples of such plots are shown in Figure 1 for two of the complexes measured. The measurements were made on a Cary-14 recording spectrophotometer. Matched, stoppered absorption cells of varying path lengths were used to obtain spectra
+
21.7
0.14
o.i$
0.18
(MI-1)/ (2M*f 1).
0.20
0.22
Figure 1. Charge-transfer energies of tetracyanoethylene (TCNE) complexes with aromatic donors us. refractive index function of the nonpolar solvents.
of optical densities between 0.5 and 0.9 (and relative donor-acceptor concentrations from approximately 50: 1 to 100: 1). The cells were thermostated at 22”. The solvents and materials used were spectrograde where available; where not, they were purified appropriately, Details of this and further experimental information will be given in the forthcoming paper.6 The results are striking. For each ct complex the data could best be analyzed in terms of two straight solvent lines of different slope: one of greater slope deriving from the saturated hydrocarbons and one of lower slope from the perfluoro solvents (Figure 1). These pairs of lines extrapolate in each case to the same intercept with the ordinate, (n2 - 1)/(2n2 1) = 0
+
(1) (a) G. Briegleb, “Elektronen Donator Acceptor Komplexe,” Springer-Verlag, Berlin, 1961, p 38; (b) C. C. Thompson, Jr., and P. A. D. deMaine, J . Am. Chem. SOC.,8 5 , 3096 (1963); (c) K. M. C. Davis and M. C. R. Symons, J . Chem. SOC.,2079 (1965); (d) H. M. Rosenberg and D. Hale, J . Phys. Chem., 69, 2490 (1965). (2) L. Onsager, J . Am. Chem. Soc., 58, 1486 (1936). (3) E. A. Bovey and S. S. Yanari, Nature, 186, 1042 (1960). (4) (a) F. T. Lang and R. L. Strong, J . Am. Chem. SOC.,87, 2345 (1965); (b) J. M. Goodenow and M. Tamres, J . Chem. Phys., 43, 3393 (1965); (c) J. Prochorow, ibid., 43, 3394 (1965). (5) M. Kroll and M. L. Ginter, J . Phys. Chem., 69, 3672 (1965). (6) C. Reid and E. M. Voigt, to be published. (7) E. G. McRae, J . Phys. Chem., 61, 562 (1957).
