Kinetics hydrogen activation energy

Supported metal catalysts are much easier to employ and have obvious attractions for industrial use from their ease of handling and economic considerations of obtaining maximum utilisation of the catalytically active metal, by using very small particles with a high surface-to-volume ratio, which are stable on the support and not susceptible to sintering. In spite of the inherent difficulties of variable activity, kinetics and activation energies [11] associated with their use, supported metals have been extensively used as hydrogenation catalysts. 

The hydrogenation of ethylene has been extensively studied over a wide variety of metal catalysts. In this section we review some of the results obtained for the kinetics and activation energies and from the use of deuterium as a tracer. 

Kinetics and activation energies for ethylene hydrogenation Rate = k Ph2 po1ua... 

Initial rate kinetics and activation energies for acetylene hydrogenation... 

Kinetics and activation energies observed in the hydrogenation of buta-1 3-diene over various metals... 

It is evident that the equation for Ref. 13 has broken down completely for CO hydrogenation. The other equations (II, 14) for CO hydrogenation gave correlations similar to those obtained by the simple kinetics. These equations are all, however, of relatively simple form. They use low activation energies and in general show an activity dependence on the square root of the pressure, similar to that of the simple kinetics. 

The NO reduction over Cu-Ni-Fe alloys has been studied recently by Lamb and Tollefson. They tested copper wires, stainless steel turnings, and metal alloys from 378 to 500°C, at space velocities of 42,000-54,000 hr-1. The kinetics is found to be first order with respect to hydrogen between 400 and 55,000 ppm, and zero order with respect to NO between 600 and 6800 ppm 104). The activation energies of these reactions are found to be 12.0-18.2 kcal/mole. Hydrogen will reduce both oxygen and NO when they are simultaneously present. CO reduction kinetics were also studied over monel metals by Lunt et al. 43) and by Fedor et al. 105). Lunt speculated that the mechanism begins by oxidant attack on the metal surface... 

Table III lists the kinetic equations for the reactions studied by Scholten and Konvalinka when the hydride was the catalyst involved. Uncracked samples of the hydride exhibit far greater activation energy than does the a-phase, i.e. 12.5 kcal/mole, in good accord with 11 kcal/mole obtained by Couper and Eley for a wire preexposed to the atomic hydrogen. The exponent of the power at p amounts to 0.64 no matter which one of the reactions was studied and under what conditions of p and T the kinetic experiments were carried out. According to Scholten and Konvalinka this is a unique quantitative factor common to the reactions studied on palladium hydride as catalyst. It constitutes a point of departure for the authors proposal for the mechanism of the para-hydrogen conversion reaction catalyzed by the hydride phase.

The reaction was followed by means of the strong absorption of the Os(II) complex at 480 m/i. Unlike the Tl(riI) + Fe(II) system, there is a slight increase in rate as the hydrogen-ion concentration is increased. The kinetic data were interpreted on the basis that both Tl and TIOH react with Os(bipy)3 (with rate coefficients and respectively). At 24.5 °C and ju = 2.99 M, kj = 36.0 l.mole. see and= 14.7 l.mole sec corresponding activation energies are 6.90 and 11.5 kcal.mole" The latter values are considerably smaller than those for the T1(III) + T1(I) exchange and for the Tl(III)- -Fe(II) reaction . On the other hand, all three reactions are subject to retardation by Cl ions. 

The kinetics of the Ce(IV) sulphate oxidation of oxalic acid are simple second order although the rate coefficient is inversely proportional to both hydrogen and bisulphate-ion concentrations, and it is also reduced at very high oxalic acid concentrations. Values of the activation energy from 13.4+1.5 (ref. 411) to 16.5+0.4 (ref. 409) kcal.mole have been reported. An intermediate has been detected spectroscopically " this decays in first-order fashion with E = 10.5+0.5 kcal.mole" and with a rate independent of acidity. However, the extent of formation of this complex is reduced as the acidity is increased ", and it appears that a less reactive dioxalato complex is formed at higher substrate concentrations ". 

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