Catalyst, catalysis; adsorption theory

Catalyst and catalysis

A catalyst is a substance that changes the rate of a chemical reaction without being consumed in the reaction (its mass and chemical nature are the same at the end). The phenomenon is called catalysis. Example: MnO₂ speeds up the decomposition of KClO₃ to O₂.

Adsorption theory of heterogeneous catalysis

According to the adsorption theory (for heterogeneous catalysis), the catalyst works by adsorbing the reactant molecules on its surface. The steps are:

1. Diffusion and adsorption: reactant molecules diffuse to the catalyst surface and are adsorbed on the active sites, so their concentration at the surface becomes high.
2. Activation: adsorption weakens the bonds of the reactants (e.g. H–H breaks), forming activated, reactive species. Adsorption is exothermic, and the heat released helps to supply activation energy.
3. Surface reaction: the activated adsorbed molecules, now close together, react to form the product.
4. Desorption: the product molecules have little affinity for the surface; they leave the surface, which frees the active sites.
5. The freed sites adsorb new reactant molecules and the cycle repeats.

 H₂    C₂H₄          H–H  C=C         H H
  ↓     ↓             ↓    ↓           | |
 ▒▒▒▒▒▒▒▒▒▒▒  →  ▒▒▒▒▒▒▒▒▒▒▒  →   H–C–C–H   (leaves)
 Ni surface      adsorbed, H–H      | |
 (adsorption)    bond broken        H H

Example: hydrogenation of ethene on nickel: C₂H₄ + H₂ → C₂H₆ (Ni). H₂ is adsorbed on Ni and splits into H atoms, ethene is also adsorbed, the H atoms add to the double bond, and ethane is desorbed.

The theory explains why a finely divided catalyst is more active (larger surface), why poisons stop catalysis (they occupy the active sites) and why promoters help (they increase the active surface).