Page 199 - Introduction to chemical reaction engineering and kinetics

8.1 Catalysis and Catalysts 181

8.1.3.2 Catalytic Effect on Reaction Rate
Catalysts increase the reaction rate by lowering the energy requirements for the re-
action. This, in turn, results from the ability of the catalyst to form bonds to reaction
intermediates to offset the energy required to break reactant bonds. An example of
a catalyst providing energetically easier routes to products is illustrated in the multi-
step reaction coordinate diagram in Figure 8.3, for the methanol-synthesis reaction,
CO + 2H2 + CHsOH. The energies of the intermediate stages and the activation en-
ergies for each step are indicated schematically.
For this reaction to proceed by itself in the gas-phase, a high-energy step such as the
breakage of H-H bonds is required, and this has not been observed. Even with H, dis-
sociation, the partially hydrogenated intermediates are not energetically favored. Also,
even if an efficient radical-chain mechanism existed, the energetic cost to accomplish
some of the steps make this reaction too slow to measure in the absence of a catalyst.
The catalytic palladium metal surface also breaks the H-H bonds, but since this reaction
is exoergic (Pd-H bonds are formed), it occurs at room temperature. The exact details
of the catalytic reaction mechanism are unclear, but a plausible sequence is indicated in
Figure 8.3. The energy scale is consistent with published values of the energies, where
available. Notice how the bonding to the palladium balances the bonding changes in
the organic intermediates. A good catalyst must ensure that all steps along the way
are energetically possible. Very strongly bonded intermediates are to be avoided. Al-
though their formation would be energetically favorable, they would be too stable to
react further.
In general, the reaction rate is proportional to the amount of catalyst. This is true
if the catalytic sites function independently. The number of turnovers per catalytic
site per unit time is called the turnover frequewy;--The reactivity of a catalyst is
the product of the number of sites per unit mass or volume and the turnover fre-
quency.

8.1.3.3 Catalytic Control of Reaction Selectivity
In addition to accelerating the rates of reactions, catalysts control reaction selectivity
by accelerating the rate of one (desired) reaction much more than others. Figure 8.4
shows schematically how different catalysts can have markedly different selectivities.
Nickel surfaces catalyze the formation of methane from CO and HZ but methanol
is the major product on palladium surfaces. The difference in selectivity occurs be-
cause CO dissociation is relatively easy on nickel surfaces, and the resulting carbon
and oxygen atoms are hydrogenated to form methane and water. On palladium,
CO dissociation is difficult (indicated by a high activation energy and unfavorable
energetics caused by weaker bonds to oxygen and carbon), and this pathway is not
possible.

8.1.3.4 Catalyst Effect on Extent of Reaction

A catalyst increases only the rate of a reaction, not the thermodynamic affinity. Since
the presence of the catalyst does not affect the Gibbs energy of reactants or prod-
ucts, it does not therefore affect the equilibrium constant for the reaction. It follows
from this that a catalyst must accelerate the rates of both the forward and reverse re-
actions, since the rates of the two reactions must be equal once equilibrium is reached.
From the energy diagram in Figure 8.4, if a catalyst lowers the energy requirement
for the reaction in one direction, it must lower the energy requirement for the reverse
reaction.