Theories of Catalysis

For any chemical reaction to occur, the reactant molecules must first be activated into an 'activated complex', and the minimum energy required to bring the reactants to this activated state is the activation energy. Activation energy can, in principle, always be supplied by simply raising the reaction temperature; a catalyst instead achieves the same activation at a LOWER temperature — in other words, a catalyst works by lowering the activation energy needed. Since lowering the activation energy allows a much larger fraction of reactant molecules to possess enough energy to react at a given temperature, the overall rate of reaction increases. Two complementary theories explain exactly how a catalyst manages to lower this barrier: the intermediate compound formation theory, which applies mainly to homogeneous catalysis, and the adsorption theory, which applies to heterogeneous (contact) catalysis.

The intermediate compound formation theory. Here, a catalyst C works by combining with one of the reactants to form a short-lived intermediate compound, which then reacts further with the remaining reactant(s) to give the final product, regenerating the catalyst unchanged at the end: for the overall reaction $A + B \rightarrow AB$, the catalysed pathway instead runs as $A + C \rightarrow AC$ (the intermediate) followed by $AC + B \rightarrow AB + C$. Crucially, the activation energies of these two catalysed steps are each LOWER than the activation energy of the original, uncatalysed single step $A + B \rightarrow AB$ — it is exactly this lower activation-energy pathway, made available by the intermediate's formation and decomposition, that accelerates the overall reaction. Four worked illustrations from real chemistry: (1) in the Friedel–Crafts reaction of benzene with methyl chloride over anhydrous $\text{AlCl}_3$, the catalyst first forms the intermediate $[\text{CH}_3^+][\text{AlCl}_4^-]$ from $\text{CH}_3\text{Cl} + \text{AlCl}_3$, which then reacts with benzene to give toluene, HCl, and regenerated $\text{AlCl}_3$; (2) the thermal decomposition of $\text{KClO}_3$ catalysed by $\text{MnO}_2$ proceeds via the intermediate $\text{KMnO}_4$ (formed from $\text{KClO}_3$ and $\text{MnO}_2$), which decomposes to release $\text{O}_2$ and regenerate $\text{MnO}_2$; (3) the formation of water from $\text{H}_2$ and $\text{O}_2$ catalysed by copper proceeds via the intermediate $\text{Cu}_2\text{O}$; and (4) the oxidation of HCl by air catalysed by $\text{CuCl}_2$ proceeds via the intermediates $\text{Cu}_2\text{Cl}_2$ and $\text{Cu}_2\text{OCl}_2$. This theory successfully explains both the SPECIFICITY of a given catalyst (since only certain reactants can chemically combine with it to form the necessary intermediate) and the observed increase in reaction rate with increasing catalyst concentration (more catalyst means more intermediate can form). Its limitations, however, are real: it fails to explain the action of catalytic poisons and promoters/activators, and it cannot account for the mechanism of HETEROGENEOUS catalysed reactions, where the catalyst and reactants are not even in the same phase to form a chemical intermediate in the first place.

The adsorption theory. Langmuir explained heterogeneous catalysis — also called contact catalysis for exactly this reason — on the basis of adsorption: the reactant molecules are first adsorbed onto the catalyst's surface, forming an activated complex there, which then decomposes to release the product. The full mechanism runs through five steps: (1) reactant molecules diffuse from the bulk gas or solution phase to the catalyst surface; (2) the reactant molecules are adsorbed onto the catalyst surface; (3) the adsorbed molecules are activated and form an activated complex on the surface, which decomposes to give the products; (4) the product molecules are desorbed from the surface; and (5) the product molecules then diffuse away from the surface back into the bulk phase. …