22.6 Coverage and Catalysis
In the 1830s, the Swedish chemist Jöns Jakob Berzelius (Figure 22.20) coined
the term “catalysis” to describe the effect of a substance that increases the speed
of a chemical reaction but is not itself consumed by the reaction. The sub-
stance itself is called a catalyst.(By contrast, an inhibitoris a substance that de-
creases the speed of a reaction.) In modern terms, a catalyst provides a differ-
ent pathway for a chemical reaction to take place, one that has a lower activation
energy, thereby increasing the rate of reaction. Ideally, a catalyst does not ap-
pear in the overall stoichiometry of the reaction and is not used up during the
course of the reaction. In reality, most catalysts eventually lose their effective-
ness due to various mechanisms including poisoning.
Catalysis can be separated into two types,homogeneous catalysisand hetero-
geneous catalysis.The difference depends on the phases of the chemical reaction
and the catalyst. If all substances including the catalyst are in the same phase,
it is considered homogeneous catalysis. Examples include aqueous-solution
reactions that are catalyzed by acid (Hions) or base (OHions), and gas-
phase reactions like the breakdown of ozone, O 3 , by chlorine atoms in the
upper atmosphere.
If catalysis occurs at a phase boundary because reactants and catalyst are
in different phases, it is heterogeneous catalysis. Examples include the de-
composition of NOxpollutants by catalytic converters in cars and the forma-
tion of H 2 O from H 2 and O 2 gases in the presence of finely divided metal
powders.
In both of the examples given for heterogeneous catalysis, gaseous reactants
are interacting with a solid catalyst and making products. The interaction must
be occurring at the surface of the solid. In order to understand how a surface
catalyzes a reaction even in a simplistic sense, we need to understand how to
model interactions between gases and surfaces.
It is a good assumption to think that the rate of the catalyzed reaction is re-
lated to the rate at which the gas reactant(s) interact with that surface. That is,
the rate of the catalyzed reaction must be related to the rate at which reactant
molecules are adsorbed on the surface:gas reactant →catalyst gas reactant (adsorbed) Rate
ads
Let us consider processes that involve the adsorption of a single gaseous species
on a solid surface. We will make two simplifying assumptions. First, we assume
that the gas molecules that are adsorbed directly onto the surface are the ones
that react faster, that is, are catalyzed. From this, we conclude immediately that
the maximum amount of gas that can be adsorbed and catalyzed would be a
complete monolayer of gas molecules. The variable coverageis defined as the
decimal fraction of possible positions on the surface that have an adsorbed gas
molecule on them. Coverage is symbolized by the Greek letter and varies be-
tween 0 (for no coverage) to 1 (for a complete monolayer of coverage).
Molecules that might be adsorbed on top of a monolayer are assumed to not
experience any catalysis effects of the surface.
Second, we assume that the adsorption of gas molecules is an elemen-
tary process so that the rate of adsorption, Rateads, can be determined di-
rectly from the stoichiometry of the reaction. The rate of adsorption is
therefore directly proportional to the concentration of the gas reactant,
which we will designate [gas]. But the rate of adsorption is also propor-
tional to the amount of surface positions available to adsorb onto. These
surface positions are called sites of adsorption.If the coverage is , then the22.6 Coverage and Catalysis 783Figure 22.20 The Swedish chemist Jöns Jakob
Berzelius (1779–1848) was considered a world
authority on chemistry in his time. In 1813, he
suggested the use of alphabetical symbols to stand
for elements in chemical formulas, thereby get-
ting away from alchemical symbols. He also in-
vented the term “catalysis” to describe the speed-
ing up of chemical reactions by the presence of
nonreactive components.© Bettmann/CORBIS