NOTES
296
we can calculate a secondary solvent kinetic isotope 4 X 3314 effect from eq. 3 using ZYH = 1246 3600 3 X 3000 = 27,102and ZVH*= 4 X 3314 2 X 3000 3600 2 X 3400 = 29,656 where 2 X 3400 has been added to the second sum to conserve bonds. We then have
+
+
+
+
+ +
27,102 - 29,656 (kH/kD)2,, = antilog = 0.30 (12) 3734 The assumed structure and frequencies of the transition state, eq. 11, are by no means unique. However all plausible alternatives give rise to secondary kinetic isotope effects that are significantly less than unity. Thus, we may calculate a primary kinetic isotope effect from the relation which for cis-3,3-diphenylcyclopropane-l,2-dicarboxylic acid is
1.6/0.30 = 5.3 (14) Analogous values for the other acids are shown in Table I. Thus, we see that a secondary solvent kinetic isotope effect masked the expected large primary kinetic isotope effect in the experimental data of Table I. Furthermore, the similarity of the primary kinetic isotope effects indicates a similar reaction mechanism for the two acid series. In the past, such comparatively large kinetic isotope effects have been associated with symmetric transition states in, which the transferring proton is bound equally strongly to both reactants.lB However, Willi and W~lfsber@~ have recently rejected this conclusion as superficial; for instance, high hydrogen-bending force constants in the transition state could produce a low experimental kinetic isotope effect even though the extent of bond making and breaking was identical. While the systems considered here axe not of the protonated ether or carbonyl type for which Willil* felt it necessary to replace the Bunton-Shiner method with an approach requiring an estimation of ratios of partition functions, it may eventually be interesting to carry out an analysis of our kinetic data using Willi’s method19 since it would yield limiting values of kH/kD that could be compared directly with experiment. (16) C. A. Bunton and V. J. Shiner, Jr., J . Am. C h . SOC.,83, 44 (1961). (16) C. A. Bunton and V. J. Shiner, Jr., dbid., 83, 3216 (1961). (17) A.V. Willi and M. Wolfsberg,C h .Ind. (London),2097 (1964). (18) A. V. Willi, Z.Naturjorsch., 19b, 461 (1964). (19) A. V. Willi, “Siiurekatalytische Reaktionen der organischen Chemie,” Verlag Vieweg, Braunschweig, West Germany, 1966, p. 91 ff.
The J o u d of Phy&
Chmtktru
In principle, one would then be able to identify the transition state as being linear or nonlinear. As our eq. 10 clearly indicates, we prefer the latter possibility.
The Catalytic Reduction of Nitric and Nitrous Oxide by R. J. Kokes Depart& o j Chemistry, The Johm Hopkins University, Baltimore, Maryland 91818 (Received July 86, 1966)
Recently,’ it has been suggested that on supported metallic catalysts hydrogen can be activated by the metal component and migrate from particle to particle. If this be so, it would seem that the catalytic reduction of nitric oxide with hydrogen should yield the same products as observed for the reaction of hydrogen atoms with nitric oxide. Harteck2 has shown that, at low temperatures, the latter reaction proceeds via the formation of (HNO),, a solid species that decomposes rapidly to yield nitrous oxide and water when the temperature is raised above -100”. Similar overall reactions have been suggested8 involving the intermediate HNO, a well-characterized species14in the gas phase. Recently in this laboratory5 we have found that the reaction reported by Harteck also occurs between nitric oxide adsorbed on Cabosil and hydrogen atoms produced in the gas phase. On the other hand, the available data in the literaturee-10suggest that the catalytic reduction yields water and ammonia, nitrogen, or hydroxylamine. Only in a very old reference’’ WBS the formation of nitrous oxide noted. Most of (1) H. W. Kohn and M. Boudart, Science, 145, 1949 (1964); S. Khoobiar, J . Phys. C h m . , 68, 411 (1964). (2) P. Harteok, Ber., 66, 423 (1933). (3) A. R. Knight and H. E. Gunning, Can. J . Chem., 41, 763 (1963). (4) M.A. A. Clyne and B. A. Thrush, DiScussWna Faraday SOC.,33, 139 (1962). (6) R. Gonzales and R. J. Kokes, unpublished results. (6) P. Neogi and B. B. Adhikary, C h m . Abstr., 5, 1031 (1911): B. B. Akhikary, ibid., 10, 24 (1916); L. Andrussov, ibid., 21, 1872 (1927); H. Tropsch and T. Bahr, ibid., 24, 4983 (1930). (7) R. J. Ayen and M. 8. Peters, Ind. Eng. Chem. Process Design Deudop., 1, 204 (1962). (8)L. Duparo, P. Wenger, and C. Unfer, Helv. Chim. Acta, 11, 337 (1928). (9) P. Sabatier, “Catalysis in Organic Chemistry,” translated by E. E. Reid, D. Van Nostrand Co., New York, N. Y.,1922,pp. 137, 181, 186. (10)A. J. Butterworth and J. R. Partington, Trans. Faraday SOC., 26, 144 (1930). (11) S. Cooke, PTOC. Phil. SOC.(Glasgow), IS, 284 (1887).
NOTES
297
these studies, however, were carried out at or above room temperature, whereas the formation of nitrous oxide by the direct reaction of hydrogen atoms with nitric oxide21Khas been directly observed only at low temperature. In view of this, we have carried out a brief investigation of the catalytic reduction of nitric oxide between -80 and 100".
Results and Discussion The catalyst used for this study was prepared by impregnation of alumina (from the hydrolysis of aluminum isopropoxide) with chloroplatinic acid followed by drying and air calcination at 400". A 0.25-g. sample was used in these runs. Prior to each series of runs, it was reduced in hydrogen for 2 hr. at 400". The surface area was 110 m.2/g. In the studies with nitric oxide, the reactant stream was hydrogen at 1 atm. passed through a saturator containing nitric oxide at -183" to give a nominal 250:l H2:NO composition. I n the studies with nitrous oxide, the reactant stream was hydrogen at 1 atm. passed through a saturator at -132" to give a nominal 1OO:l H2:N20 composition. Temperature control at or below room temperature was achieved by liquid baths manually controlled. Above room temperature an oven was used. Products collected in a trap at - 195" were analyzed mass spectroscopically with a few check runs by gas chromatography. Figure 1 shows the analysis of condensables (except for water) found in the effluent stream as a function of catalyst temperature. Points for ammonia, and nitrous oxide very near 0% are shown only when these products were detected. The solid symbols indicate points obtained on increasing the temperature; the open-circle point indicates the result of a check run performed after the above sequence; the triangular points show results obtained after the catalyst was reduced at 400" for 2 hr. The results definitely show that nitrous oxide is the product formed first as the temperature is increased. The occurrence of a maximum is consistent with the initial formation of HNO followed by one of the subsequent steps 2HNO -+
H20
+ N20 -% 2NH3 + 2H20
or
TEMPERATURE
"K
Figure 1. Composition of the exit gas (water excluded) as a function of temperature in the catalytic hydrogenation of nitric oxide.
nitrous oxide over the same catalyst yielded no detectable ammonia. (Nitrogen would not be detected by our method of analysis, but since reduction did occur, as evidenced by water formation, it is likely that the sole product is nitrogen and water as reported by other 12)
Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research. (12) See ref. 9, pp. 135, 181.
On the Refractive Indices of
Aqueous Solutions of Urea by John R. Warren' and Julius A. Gordon Biochemistry 8&ion, Depart& of Pathology, University of Colorado Medical Center, Denver, Colorado 80220 (Received August I , 1966)
11[1 NHs + H20
Recently, it has been reported28that the refractive index, viscosity, and density of aqueous urea solutions undergo identical sharp transitions below 1 M ; the refractive index becomes completely linear with con-
The first possibility is unlikely for two reasons. First, variation of the flow rate by a factor of 10 caused only slight variation in the composition of products. Second, at room temperature and at 125O, hydrogenation of
(1) Trainee under U. S. Public Health Services Grant GM-97703. (2) (a) V. K. Venkateaan and C. V. Suryanrtrayana, J . Phys. Cbm., 60, 776 (1956); (b) these authors report their data, in terms of normality. They apparently assume the equivalent weight of urea to be half the gram molecular weight for reasons unclear to us.
HNO +H20
+ N20
Volume 7'0,Number 1 January 1066
