Catalysis: Homogeneous and Heterogeneous

A catalyst is a substance that changes the rate of a chemical reaction -- in almost

every practically useful case, increases it -- without itself undergoing any net, permanent chemical

change over the course of the reaction, and without altering the overall thermodynamics of the

reaction (its $\Delta H$, or the position of equilibrium it eventually reaches -- a catalyst speeds up

the forward and reverse reactions equally, so it shortens the time needed to reach equilibrium but

never shifts where that equilibrium lies, and never changes the equilibrium constant $K$). Mechanistically,

a catalyst works by opening up an alternative reaction pathway, via a different activated complex, whose

activation energy is lower than that of the uncatalyzed pathway -- as explained in the section on

activation energy, a lower $E_a$ allows a far larger fraction of reactant molecules to react at the same

temperature.

Catalyzed reactions are classified into two broad types, according to whether the catalyst occupies the

same phase as the reactants or a different one.

Homogeneous catalysis. Here the catalyst is present in the same phase (solid, liquid, or gas) as

the reactants, typically dissolved together in a single solution. A classic example is the

acid-catalyzed hydrolysis (or esterification) of an organic ester, where dissolved $\text{H}^+(aq)$ ions

catalyze the reaction while remaining in the same aqueous liquid phase as the ester, water, acid and

alcohol/carboxylic-acid products throughout. Historically, the oxidation of $\text{SO}_2$ to

$\text{SO}_3$ by nitric oxide gas in the old lead-chamber process for sulfuric acid manufacture is

another example, with reactants and catalyst all in the gas phase.

Heterogeneous catalysis. Here the catalyst exists in a different phase from the reactants --

almost always a solid catalyst acting on gaseous or liquid reactants, with the reaction taking place at

the catalyst's surface (reactant molecules are first adsorbed onto active sites on that surface, react

there, and the products then desorb). Major industrial examples include the Haber process for

ammonia synthesis, $\text{N}_2(g) + 3\text{H}_2(g) \xrightarrow{\text{Fe}} 2\text{NH}_3(g)$, using finely

divided iron (with small amounts of added molybdenum as a "promoter" that enhances the iron's catalytic

activity); the Contact process for sulfuric acid, which uses solid vanadium pentoxide,

$\text{V}_2\text{O}_5$, to catalyze the oxidation of $\text{SO}_2$ to $\text{SO}_3$; and the catalytic

hydrogenation of vegetable oils to solid fats using finely divided nickel. Because a heterogeneous

catalyst acts through a fixed number of active surface sites, its own rate can itself display zero order …