ETHER-BORON HALIDEADDITIONCOMPOUNDS IN DICHLOROMETHANE
macroscopic dielectric constant. In a solvent such as ethylene dichloride, there appears to be a significant difference between the macroscopic dielectric constant and the effective LLmicroscopic”dielectric constant which exists near the ionic and molecular species in solution. Inami, Bodenseh, and Ramsey’o have indicated that the effective dielectric constant of ethylene dichloride may be as large as 12.4 in solutions of n(C4He)4NC104 a t 25” (eniscro = 10.232). Glueckauf’s’l corrections of the conventional DebyeHuckel expression t o include short-range changes in
89
the dielectric strength are insignificant a t the temperatures and concentrations used in this study. As both A and K in eq. 3 are functions of the dielectric constant, our reported degree of dissociation may be slightly in error. This could account for the variation between the calculated and observed minima in the degree of dissociation. Further work will be necessary to clarify this point. (10) Y.H.Inami, H. K. Bodenseh, and J. B. Ramsey, J . Am. Chem. SOC.,8 3 , 4745 (1961). (11) E.Glueckauf, Trans. Faraday Soc., 60, 776 (1964).
A Nuclear Magnetic Resonance Znvestigation of Ether-Boron Halide
Molecular Addition Compounds in Dichloromethane
by Ernest Gore and Steven S. Danyluk Department of Chemistry, Univereity of Toronto, Toronto 6 , Ontario, Canada
(Received M a y 87, 1964)
A study has been made of the stabilities of a number of ether-boron halide addition compounds in dichloromethane a t 23”. Equilibrium constants were determined for the reaction, RzO BX3 e RzO.BX8, by a least-squares analysis of the chemical shift-concentration curves for these systems. A shift to low field was observed for all of the ether protons on complexing with boron halide. The most marked deshielding (- 1.25 p.p.m.) was noted for the 1 : l diethyl ether-boron trichloride compound and has been attributed to the formation of ethyl ethoxychloroborate, CzHs+C2HSOBCl3-. Boron trichloride was found to be a stronger acceptor than boron trifluoride while the relative donor strengths of the ethers studied decreased in the order: (CzH&O 2 [(CH3)&H],0 > (CH3)zO. The disagreement between the present n.m.r. results and earlier infrared and gas-phase dissociation studies has been attributed to the influence of the solvent on the mean chemical shift of ether-boron halide solutions in dichloromethane.
+
Introduction 14olecular addition compounds formed with group 111-A acceptors such as boron and aluminum halides (A!&) have been widely investigated and recent reviews emphasize the scope and importance of these compounds as intermediates in many organic reactions. 1-5 A variety of techniques including gas-phase dissociation,6 c r y o s ~ o p y ,electrical ~*~ conductivity,7~e~10 and infrared
spectroscopyll have been used to establish the structure and stabilities of group I11 addition compounds. Re( I ) W. Gerrard, “The Organic Chemistry of Boron,” Academic press, New N. y., (2) A. V. Topchiev S V. Zavgorodnii, and Ya. M .Pauskin, “Boron Fluoride,” Pergam;n London, 1959.
pres,,
(3) F, G, A, Stone, Chem, R e v , , 58, lol (1958), (4) R.s. Mulliken, J . p h y e . Chem., 5 6 , 801 (1952).
Volume 69, Number 1
January 1966
90
cently, n.1n.r. measurements have provided additional information about relative s t a b i l i t i e ~ ' ~and ~ ' ~ exchange rates14 of boron addition compounds. As part of a general study of solvent effects on the stabilities of addition compounds, we have determined the equilibrium constants for boron halide-ether complexes in dichloromethane. Chemical shifts were measured at 23' for the ether protons at varying boron halide concentrations and equilibrium constants were obtained by fitting the resultant chemical shift-concentration curves by a least-squares procedure. The relative stabilities for the complexes, determined by this method, are in disagreement with the results of earlier work.l18l3 This discrepancy can be attributed to the effect of the solvent niolecules on the stability of the 1:1 complex. Experimental Reagent grade gases, BFI, BC4, and (CH&O (Matheson Co.) , were further purified by the freeze-thaw technique in the vacuum system prior to passage into the reaction tube. Ethyl ether (Baker, analyzed) was used without further purification. The isopropyl ether (Fisher, certified reagent) was treated with ferrous sulfate to remove peroxides, dried over fresh Drierite, and then distilled in a 40-cni. vacuum-jacketed column packed with glass helices. The center portion boiling at 67.3' (uncor.) was collected for use. Dichloroniethane (Fisher, certified reagent) was dried over Drierite and distilled in the 40-cm. column; the fraction boiling at 38.9" (uncor.) was retained. Solutions were prepared in a vacuum system and standard vacuuni techniques were used to transfer the gases into calibrated volumes and then into the reaction tube. Diethyl ether and diisopropyl ether solutions in dichloromethane were prepared gravimetrically and were degassed before addition of the boron halide. Dimethyl ether was introduced by vacuum transfer into the reaction cell. The final ether concentrations were approximately 3 mole %. The shifts observed a t this concentration differed by less than 0.1 C.P.S.from the shifts a t infinite dilution. The reaction tube was equipped with a magnetic stirrer and was attached to the vacuum system oia a number of flexible joints. These permitted the tube to be tipped on its side in order to transfer a portion of the solution to an attached 5-nim. spinning tube. The solution in the side arm was cooled to liquid nitrogen temperature and sealed off prior to measurement. A slow irreversible reaction was noted in diethyl etherboron trichloride solutions at room temperature, In order to minimize the effects of such reactions, the solutions were stored in Dry Ice until their measurement. The Journal of Physical Chemistry
ERNEST GOREAND STEVENS. DANYLUK
The n.m.r. spectra were recorded with a Varian DP60 spectrometer a t 23 f 1'. All of the chemical shifts were measured relative to internal dichloromethane by the usual side-band technique. A minimum of four spectra were recorded for each concentration and the mean value was taken as the chemical shift. Chemical shifts for the methylene protons in ethyl ether complexes and the tertiary proton in diisopropyl ether addition compounds were taken as the midpoints of their respective multiplets. The measured chemical shifts are accurate to *0.2 c.p.s. Calculations The relative stabilities of ether-boron fluoride addition compounds were determined by Craig and Richards13from an analysisof the mixing ratio, L e . , n ~ ~ ~ / n f t ~ o , chemical shift plots obtained for binary mixtures of the parent compounds. Satisfactory straight-line plots were obtained for each of the two component systems studied. In dichloromethane solutions, however, a linear relationship was only obtained for the diethyl ether-boron trichloride system and an alternative procedure was therefore adopted for evaluating the equilibrium constants. The equilibrium constant for formation of the 1 : l addition compound RzO
+ BX3 e RzO*BX3
is given by
K =
+
z(a b - 2) (a - z)(b - 5 )
(1)
where the initial moles of boron halide and ether are given by a and b, respectively, and z is the moles of 1: 1 compound at equilibrium. Since the chemical shifts of the ether protons were independent of concentration a t low concentrations in dichloromethane (3 mole %) the (5) H.H.Prekampus and Th. Kranz, 2.physik. Chem. (Frankfurt), 34, 213 (1963). (6) D.E. McLaughlin and M. Tamres, J . Am. Chem. SOC.,82, 5618 (1960). (7) N. N. Greenwood and R. L. Martin, Quart. Rev. (London), 8 , 1 (1954). (8) H.E.Wirth, M. J. Jackson, and H. W. Griffiths, J . Phys. Chem., 62, 871 (1958). (9) A. V. Topchiev, Ya. M. Pauskin, T . P . Vishny, and M. V. Hur, Dokl. Akad. Nauk SSR, 80, 381 (1951). (10) E. Gore and S. S. Danyluk, unpublished results. (11) H . E. Wirth and P. I. Slick, J . Phys. Chem., 66, 2277 (1962). (12) T.D. Coyle and F. G. A. Stone, J . A m . C b m . SOC.,83, 4138 (1961), (13) R.A. Craig and R. E. Richards, Trans. Faraday SOC..59, 1972 (1963). (14) S. Brownstein, A. M. Eastham, and G . A. Latermouille, J . Phys. Chem.. 67, 1028 (1963).
ETHER-BORON HALIDEADDITION COMPOUNDS IN DICHLOROMETHANE
solvent concentration was not included in the calculations. If the initial mole fraction of boron halide is c b ) ; setting a equal to 1 mole, the then c = a / ( a equilibrium constant can be rewritten in the form
+
I n the present work only one set of signals was observed for the ether, indicating a rapid exchange of the ether molecule between free and complexed states. The observed chemical shift, dobsd, therefore corresponds to the weighted average of the chemical shifts for the addition compound, Stomp, and the free ether, and is given by
Solving (2) for 2 and substituting in (3), the observed chemical shift is then given by aobsd
= 6R,O-+
4K (6comp
-
~R,o)
2(1 - c )
c ( l - c)
(4)
The experimental and calculated shifts were analyzed by a least-squares procedure to obtain K and Stomp values which gave the best fit of the experimental data. Calculations were carried out with an IBM 7090 computer using a program written for this purpose. The limits of errors for the calculated equilibrium constants and chemical shifts could not be evaluated satisfactorily by Gutowsky’s method. l 5 An alternative procedure was therefore adopted and consisted basically of an analysis of the effect of random errors on K and beomp. A Fortran program was written to apply a random error to each chemical shift &bsd in turn. The random numbers were chosen to lie in the interval -1 to +l. For example, if the number lay within the interval 1) an “error” of 0.5 C.P.S.was added to the shift; for a number in the interval (- l/3, the “error” added was zero, and for theinterval ( - 1) an “error” of -0.5 c.p.s was added. In this manner a series of random errors was generated and a new set of data was calculated in which a random error of f 0 . 5 C.P.S. had been introduced. With this set of data a new K and Stomp were calculated. The procedure was repeated a minimum of 25 times and from the 25 dif-
91
The effect on the chemical shift &bsd of an additional reaction involving 2 moles of BX3
RzO
+ 2BX3
RzO.2BX3
to form the 1: 2 addition compound was also investigated. However, the chemical shift-concentration curves could not be fitted satisfactorily when this reaction was included in the calculation and the chemical shift measurements therefore indicate that this reaction does not occur to any appreciable extent in the present systems over the concentration ranges studied.
Results 1 . Dimethyl Ether Addition Compounds. A single signal was observed for the methyl protons in both the BF3 and BC13 systems at the concentrations studied. The boron halide is therefore undergoing a rapid exchange between ether molecules at 23’. A slight broadening of the line widths was noted for (CH3)20BC13 solutions with increasing BC13 concentration; no change in line width was noted in BF3 solutions. No new bands were observed to form in both systeiiis over a period of several days. The variation of chemical shifts with boron halide concentration is illustrated in Figure 1. Concentrations are given in terms of the mole ratio of boron halide to ether in dichloroniethane as solvent and the cheiiiical shifts are relative to the solvent as reference. The solid curve represents the calculated chemical shifts obtained with the best values for K and Seomp. A summary of chemical shifts, ,S and equilibrium constants for the dimethyl ether compounds is given in Table I. The standard deviations between calculated and observed curves are also given in Table I. A slight but measurable low-field shift of the pure ether signal was noted in dichloroniethane a t concentrations greater than 5 mole %. Table I: Equilibrium Constants and Chemical Shifts for 1 : 1 Complexes
K
Complex
BF3.Me20 BCl$.MezO BFa.EtO BCla.Et20
BFa.i-PrpO
’,,S
3.36 f 1.06 90.9 f 6 . 3 171.4 f 7 3 . 3 472.6 f 7 7 . 8 106.8 f 19.1
6,0,p,5
c.p.8.
61.2 f6.5 61.9 f 0 . 8 65.1 f 0 . 5 36.6 f 0 . 2 30.1 f 0 . 5
is given relative to internal CH&12.
Standard deviation between calculated and experimental curves, c.p.8.
2.17 2.09 1.20 1.50 1.31
ERNEST GOREAND STEVENS. DANYLUK
92
“130
O
I
c
N
(Y
5l
I
0
0 c L
0 P 0
CD P 0
Y
Y
I-
I-
70
-
!k
LL r
I
a
a
60-
0
32o0
2
6
0
3I 0
t
1
0
IO
40
I 20
1
I
30
40
I 50
*
I 60
i
DIMETHYL ETHER COMPLEXES Figure 1. Variation of the chemical shift for dimethyl ether-boron halide solutions in dichloromethane a t 23’ : , calculated shifts. 0, 0 , measured shifts;
-
2. Diethyl Ether Addition Compounds. A single triplet and quartet, characteristic of the methyl and methylene protons in an ethyl group with J = 6.9 f 0.10 c.P.s., was observed for all of the diethyl etherBF3 solutions. No change of the coupling constant was noted with increasing concentration of BF3. Chemical shifts for the methylene protons are shown in Figure 2. Corresponding changes in the methyl group shifts were considerably less and have not been included. Analysis and K values of the chemical shift data yields the 6,,, listed in Table I. At high BF3 concentrations a discernible broadening of the niultiplets was observed, indicating ‘a change in the exchange rate for the BF3. All of the diethyl ether-BF3 solutions were stable over a period of several days. Freshly prepared solutions of BC13 and diethyl ether in dichloroniethane also showed a characteristic ethyl group pattern. The chemical shift changes are shown in Figure 2 and the resultant K and Bcomp values are listed in Table 1. Several differences in behavior from other ether--boron halide solutions were noted, however.
0.i
Physical Chemistry
20
30
40
SO
60
‘
0
I
CONCENTRATION (mole ‘10 BX,)
The Journal
IO
CONCENTRATION (mole Y o EX,) DIETHYL ETHER COMPLEXEB Figure 2. Variation of the chemical shift for diethyl ether-boron halide solutions in dichloromethane at 23 O : 0, 0 , measured shifts; , calculated shifts.
-
A pronounced broadening of the multiplet signals was observed with increasing BC13 concentration. For example, line widths for the methylene protons increased steadily from 0.70 C.P.S.for pure diethyl ether to 3.5 C.P.S.for 27 mole % BC13 and then decreased again to 1.0 C.P.S. for a 1: 1 mixture. The spectra for solutions with mixing ratios greater than 1 undergo a pronounced change with time when allowed to stand at room temperature. A similar time dependence has been noted previously16 for the B” spectrum of (C, H6)20-BC13. A decrease of intensity was observed for the broad signals corresponding to ( C Z H , ).BC13 ~~ and two new sets of quartets and triplets appeared with the quartets centered at 63 f 1 C.P.S. and 106 f 1 C.P.S. relative to dichloroniethane. After a period of several weeks, no signals were observable for the addition compound while intensities for theother signals had increased proportionately. A check of the chemical shift and coupling constant for the quartet a t 106 C.P.S. showed that it was identical with that obtained for dilute solu(16) T. P. Onak, H. Landesman, R. E. Williams, and I. Shapiro, J . Phys. Chem., 63, 1533 (1959).
ETHER-BORON HALIDE ADDITIONCOMPOUNDS IN DICHLOROMETHANE
‘110
*
O
O
93
from 6.2 to 6.6 C.P.S.with increasing BF3 concentration; no change in line width was observed over the concentration ranges studied.
Discussion I . Proton Chemical Shifts. The proton resonances for all of the ethers are shifted to low field on complexing with boron halide; protons a to the oxygen are shifted by amounts approximately five times as large as p-protons. This deshielding results because of a decrease of electron density on the donor niolecule (largely a t the donor site) upon coordination with the acceptor. The 70 decrease in shielding for a-protons in a given ether is significantly greater for the more stable BC13 compounds than for BF, conipounds, in agreement with the eo trends observed for other addition compounds. l 2 However, no correlation is noted between (BR,o - 6comp) for 60 different ethers and the equilibrium constants of the corresponding addition conipounds, indicating that factors such as steric hindrance and bond anisotropies are influencing the shifts in addition to electron density changes. 30 A comparison of the chemical shifts for (CH3)20.BF3 and (CzHs)zO.BF3 reported by Craig and Richards13 with shifts for equimolar solutions of the components in dichloromethane shows that the latter are a t higher field CONCENTRATION (mole YOBFJ in each case. I n addition, a linear plot of mixing ratio OIISOPROPYL ETHER COMPLEX us. 6 was only obtained for (CzHs),O-BCI, solutions in Figure 3. Variation of the chemical shift for diisopropyl dichloroniethane; S-shaped curves resulted for all of ether-boron trifluoride solutions in dichloromethane a t 23 ’: the other systems with the largest deviation noted for 0, measured shifts; , calculated shifts. (CH3)z0-BF3 solutions. Both of these differences can be attributed to a partial dissociation of the addition compounds in dichloromethane a t 23”. A sigtions of CzHsCl in dichloromethane. This multiplet is nificant difference (not explainable by susceptibility accordingly attributed to CzHsC1formed in the irreverchanges) is also noted between the calculated shift, sible reaction of BC13 with diethyl ether. The second , , , , 6 , for (CH3)20.BF3 in dichloroniethane and the set of signals with considerably broader line widths shift for the pure compound.13 I n this instance it is -2.0 C.P.S.is assigned to CzH50BC12. This compound likely that the discrepancy is due to hydrogen-bond inhas been postulated as an intermediate in the over-all teraction between (CH3)20 and dichloromethane (secreaction of diethyl ether and BC13 (cj. Discussion). The 3). tion CH3-CH2 coupling constants for the two new quartets Of the ethers studied, the largest shift to low field on were different from the (C2H6)20.BC13quartet. For the coordination with boron halide was observed in the quartet at 63 c.P.s., J = 7.2 f. 0.2 c.P.s., and for the diethyl ether-boron trichloride solutions. The niethylquartet a t 106 c.P.s., J = 6.3 f 0.2 C.P.S. ene proton quartet of the addition compound, T 3. Diisopropyl Ether Addition Compound. The 5.32 p.p.ni., lies outside the range norlnally observed proton spectra for diisopropyl ether-boron fluoride for ethyl derivatives, z.e., T 5.64 to 6.80 p.p.ni., and is solutions showed a septet and doublet characteristic of in fact 0.32 p.p.m. to low field from ethyl fluoride.I7 the isopropyl group. The variation of chemical shift This rather surprising deshielding cannot be interpreted for the tertiary proton (midpoint of the septet) with in ternis of any reasonable combination of electronegBF3concentration is shown in Figure 3. A summary of ativity and diamagnetic anisotropy changes associated parameters derived from Figure 3 is given in Table I. Solutions of diisopropyl ether-BF3 were stable over a period of several days with no evidence of other signals (17) “ N M R Spectra Catalogue,” Vol. 1, Varian Associates Limited, apparent. The coupling constant JcH~--Hincreases Palo Alto, Calif., 1962.
-
-
401
Volume 69,Number 1
January 1.966
94
with formation of the B-0 coordinate bond. A more likely possibility is the presence of the addition compound as an ethyl ethoxychloroborate salt, [CzHs]+[CzH50BC13]-, in solution with the positive charge localized on the methylene carbon of the C2Ha+ion and the negative charge delocalized along the 0-B-C1 bonds. The existence of such an ionic compound is supported by other experimental data. A study of the electrical conductivities of diethyl ether-BC13 solutions1° has shown that the addition compound is in fact ionized to a significant extent in the molten state and in dichloromethane. For example, the equivalent conductance of a 1: 1 mixture of (CzH5)2Oand BC13 is 0.277 ohm-' in dichloromethane a t 25'. This is somewhat higher than the value reported for the molten (C2H6)20.BF3 addition compound where the existence of ethyl -, has been ethoxyfluoroborate, [C2H5]+ [CzH50BF3] confirmed by electrolysis measurements. 9,1* I n addition, conductivity changes with time observed for freshly prepared diethyl ether-BC13 solutions in dichloromethane can be correlated with changes observed in the proton resonance spectra for these solutions. It is likely that the ethyl ethoxychloroborate is formed as an intermediate in the irreversible reaction occurring a t room temperature to form ethyl chloride and ethylchloroborinate. l9 2 , Relative Acceptor Strengths of the Boron Halides. The stabilities of the addition compounds in dichloromethane vary in the order R20.BC13 > Rz0.BF3 for a given ether. Boron trichloride is therefore a stronger electron acceptor than boron trifluoride. This is in agreement with the relative acceptor strengths established in studies of dipole moments and heats of reaction for pyridine,20-22piperidine,z1and acetonitrilez3addition compounds. The order of acceptor strengths is consistent with a higher x-bond character in the B-F bond as compared with the B-Cl b ~ n d . ~ vThus, ~ * although it might be expected on the basis of electronegativities and steric requirements that BF3 should be a stronger acceptor than BC13,the reorganization energy for the change in hybridization of the B atom from spZ to sp3 is much higher for BF3and accounts for its weaker acceptor properties. 3. Relative Donor Strengths of Alkyl Ethers. It has been shown in earlier n.m.r.,l 3 gas phase dissociation,6 and distribution studies" that dimethyl ether is a stronger donor toward BF3than diethyl ether, while diisopropyl ether is a somewhat weaker donor than diethyl ether. Although the opposite order of base strengths would be expected on the basis of simple inductive effects, it is generally assumed that a strong steric effect predominates in the case of the diethyl and diisopropyl ethers. However, a recent study of inThe Journal of Physical Chemistry
ERNEST GOREAND STEVENS. DANYLUK
frared association shifts for hydroxy compounds25 has shown that a larger shift, A V O Hoccurs , for methanol and phenol solutions containing diethyl ether as donor than for solutions containing dimethyl ether. Since the steric requirements for hydrogen-bond formation are undoubtedly less stringent than for addition compound formation with boron halides, the order of base strengths for the ethers is as expected on inductive grounds. I n the present investigation, the relative donor strengths decrease in the order diethyl ether > diisopropyl ether > dimethyl ether for a given boron halide in dichloromethane a t 23". The dimethyl ether addition compounds are clearly Iess stabIe than the other compounds with (CH3)20.BF3 having the lowest equilibrium constant of the five compounds studied. The relatively low stabilities of the dimethyl ether compounds in dichloromethane are also indicated by electrical conductivity and vapor pressure measurements on these systems.'O A quantitative comparison of the present equilibrium constants with earlier results is not feasible because of wide differences in experimental conditions. However, the disagreement in the order of donor strengths and, in particular, the low stability observed for (CH3)20.BF3 is surprising. It seems unlikely that the dichloromethane exerts any influence on the steric requirements of the higher ethers and the observed differences must be due to other complicating factors. Although dichloromethane has been regarded as an inert solvent in this work, the possibility of a weak hydrogen-bond interaction with ether molecules cannot be ruled out completely and would in fact be favored for dimethyl ether. A small concentration-dependent shift (3-5 c.p.8.) of dimethyl and diethyl ether protons was noted over the range 50 mole yo ether to 1 mole % ether in CH2Cl2, A competing interaction of this type would alter the relative amounts of free and complexed ether calculated in (3) and would affect the calculated chemical shift-concentration curves in Figures 1-3. The
~
(18) N. N. Greenwood, R. L. Martin, and H. J. Emelbus, J. Chem. Soc., 3030 (1950). (19) H.Ramser and E. Wiberg, Ber., 63, 1136 (1930). (20) C. M. Box, A. R. Katritzky, and L. E. Sutton, J. Chem. Soc., 1248 (1958). (21) N. N. Greenwood and K. Wade, ibid., 1141 (1960). (22) H . C. Brown and R. R. Holmes, J. Am. Chem. Soc., 78, 2173 (1956). (23) A. W. Laubengayer and D. S. Sears, ibid., 67, 146 (1945). (24) F.A. Cotton and J. R. Leto, J. Chem. Phys., 30, 993 (1959). (25) C. H. Van Dyke and A. G. MacDiarmid, J. Phys. Chem., 67, 1930 (1963).
95
CATALYSIS OVER SUPPORTED METALS
rather poor fit of the calculated and experimental chemical shift-concentration curves for the (CH3)20-BC13 and (CH3)2-0-BF3systems could also be explained on this basis. An additional experimental difficulty arising from the high volatillty of (CH3)20.BF3 in dichloro-
Catalysis over Supported Metals.
111.
methane would also affect the accuracy of the d-concentration plot. Acknowledgments. The financial assistance of the National Research Council and the Sational Cancer Institute of Canada are gratefully acknowledged.
Comparison of
Metals of Known Surface Area for Ethane Hydrogenolysis
by J. H. Sinfelt, W. F. Taylor, and D. J. C. Yates Process Research Division, Esso Research and Engineering Co., Linden, New Jersey
(Received J u n e 3, 1964)
The kinetics of hydrogenolysis of ethane to methane have been investigated over a series of silica-supported metal catalysts containing 10 wt. % metal. The metals studied were nickel, cobalt, platinum, and copper, the surface areas of the metals being determined by hydrogen chemisorption. Over the range of temperatures studied (175-385"), the specific catalytic activities of the nickel, cobalt, and platinum for ethane hydrogenolysis vary in the order: Ni > Co > Pt. The position of copper in this sequence is far below that of nickel or cobalt, but its position relative to platinum changes over the range 173-385'. The rate of hydrogenolysis was found to be essentially first order in ethane pressure and to decrease with increasing hydrogen pressure over all the metals. However, the magnitude of the hydrogen pressure effect varied markedly for the different metals, the effect being greatest for nickel and platinum and least for copper. Apparent activation energies ranged from a maximum of 54 kcal./mole for platinum to a minimum of 21 kcal./mole for copper.
I. Introduction Much of the fundamental work on catalysis over nietals has been done using evaporated metal films as catalysts. I n the classical work of Beeck and coworkers,' the catalytic activities of various metal films were determined for ethylene hydrogenation, and it was shown how the activities could be related to the lattice spacings. Boudart,2 and subsequently also Beeck,' showed how the activities could be equally well explained in terms of an electronic picture. I n any case, the activities of various metal films for ethylene hydrogenation were clearly established by the work of Beeck and co-workers. The activities of supported metals for ethylene
hydrogenation were studied by Schuit and van Reijen3 in an effort to determine whether the earlier findings of Beeck and co-workers with metal filins applied to supported metals. These workers reported data on a series of silica-supported metals arid concluded that the results were generally in agreement with the results obtained over metal films. In this work the authors used hydrogen chemisorption measurements to enable them to account for differences in the surface areas of the various metals. This type of study is extremely (1) 0. Beeck, Discussions Faraday SOC.,8 , 118 (1950) (2) M.Boudart, J . Am. Chem. SOC, 72, 1040 (1950) (3) G. C. A. Schuit and L. L. van Reijen, AdLan Catalysis, 10, 242 (1950).
Volume 69,Number 1
January 1966
