Introduction
Many important reactions — including most industrial processes and biological ones — happen not in bulk solution but at a surface or interface, where the chemistry is governed by adsorption and catalysis rather than ordinary solution kinetics.
1. Adsorption and its types
Adsorption is the accumulation of a substance (adsorbate) at the surface of another (adsorbent), distinct from absorption (uptake throughout the bulk). Physisorption involves weak van der Waals forces, is reversible, and typically has low enthalpy of adsorption; chemisorption involves actual chemical bond formation between adsorbate and surface, is far stronger, usually irreversible, and involves much higher enthalpy — chemisorption is the type relevant to catalysis, since it activates the adsorbed molecule for reaction.
2. Adsorption isotherms
An adsorption isotherm describes how the amount adsorbed varies with pressure (for a gas) or concentration (for a solution) at constant temperature. The Langmuir isotherm models adsorption as forming a single layer (monolayer) on a fixed number of equivalent surface sites, predicting that surface coverage rises with pressure and eventually saturates once all sites are filled — a simple, widely used approximation for chemisorption on catalytic surfaces.
3. Heterogeneous and homogeneous catalysis
A catalyst speeds up a reaction without being permanently consumed, by providing an alternative reaction pathway with lower activation energy. In homogeneous catalysis, the catalyst is in the same phase as the reactants (e.g., an acid catalyst dissolved in the same liquid reaction mixture). In heterogeneous catalysis, the catalyst is in a different phase (typically a solid catalysing a gas- or liquid-phase reaction) — reactants adsorb onto the catalyst surface (chemisorption), react while adsorbed, and the products then desorb, freeing the site for further use. Industrial examples include the iron catalyst in the Haber process (ammonia synthesis) and platinum in catalytic converters.
4. Catalyst action and activation energy
A catalyst does not change a reaction’s thermodynamics (ΔG, and hence the equilibrium position, are unaffected); it only lowers the activation energy barrier, increasing the rate at which equilibrium is reached — visualised as a reaction-coordinate diagram where the catalysed path has a lower peak than the uncatalysed path.
5. Colloids and related surface phenomena
Colloids are mixtures where one substance is dispersed as very fine particles (larger than individual molecules but too small to settle out) throughout another — examples include milk, fog, and many paints. Their stability depends heavily on surface phenomena: charged particle surfaces repel each other (preventing aggregation), and this can be disrupted (causing coagulation) by adding ions of opposite charge, which is the principle behind water purification using alum.
Key Takeaways
- Physisorption (weak, reversible, van der Waals forces) and chemisorption (strong, often irreversible, true chemical bonding) differ fundamentally in strength and mechanism; catalysis relies on chemisorption.
- The Langmuir isotherm models monolayer adsorption on fixed sites, predicting saturation at high pressure/concentration.
- Catalysts lower activation energy without changing overall thermodynamics or the equilibrium position, only the rate of approach to it; heterogeneous catalysis proceeds via adsorption, surface reaction, and desorption.
- Colloidal stability depends on surface charge, and its disruption (coagulation) underlies practical applications such as water treatment.