Heterogeneous catalysts are usually solids, while the reactants are gases or liquids. The reaction occurs at specific locations on the catalyst surface known as active sites.
The process follows the Surface Adsorption Theory, which involves three main steps: 1. Adsorption (reactants form temporary bonds with the surface), 2. Reaction (bonds in reactants weaken and new bonds form), and 3. Desorption (products break away from the surface).
The strength of adsorption is a critical variable. If adsorption is too weak (e.g., Silver), reactants do not stay on the surface long enough to react; if it is too strong (e.g., Tungsten), products cannot leave the surface, blocking active sites.
To maximize efficiency and minimize cost, catalysts are often spread as a thin layer over a support medium, such as a ceramic honeycomb, to provide a massive surface area.
In homogeneous catalysis, the catalyst and reactants exist in the same phase, typically an aqueous solution. This often involves the formation of a reactive intermediate species.
A classic application is overcoming electrostatic repulsion. For example, a reaction between two negatively charged ions is naturally slow due to repulsion. A positive transition metal ion can react with one negative ion, changing its own oxidation state, and then react with the second, facilitating the process.
The catalyst is regenerated at the end of the cycle. The reaction pathway is characterized by multiple steps, each with a lower activation energy than the single-step uncatalyzed version.
Autocatalysis occurs when one of the products of a reaction acts as a catalyst for that same reaction. This leads to a unique kinetic profile where the reaction rate is initially slow but increases as the catalyst-product accumulates.
On a concentration-time graph, the gradient (rate) becomes steeper over time before eventually leveling off as reactants are depleted. This is the opposite of standard reactions where the rate is highest at the start.
Transition metal ions produced during a redox reaction often serve as autocatalysts by participating in a redox cycle that speeds up the remaining reactants.
| Feature | Heterogeneous Catalysis | Homogeneous Catalysis |
|---|---|---|
| Phase | Different from reactants (usually solid) | Same as reactants (usually liquid/gas) |
| Mechanism | Surface adsorption/desorption | Formation of intermediate species |
| Active Sites | Reaction limited to surface sites | Reaction occurs throughout the bulk phase |
| Separation | Easy to filter or separate from products | Difficult to separate from products |
| Industrial Use | Preferred for continuous flow processes | Less common in large-scale industry |
Terminology Precision: Always distinguish between adsorption (surface attachment) and absorption (penetration into the bulk). Using the wrong term is a frequent source of lost marks.
Oxidation State Cycles: When explaining a catalytic cycle, clearly state the initial oxidation state, the intermediate state after reacting with the first species, and the regeneration of the original state.
Graph Interpretation: If you see a rate graph where the slope increases after the start, immediately identify it as autocatalysis. Be prepared to explain why the rate eventually slows down (reactant depletion).
Surface Area Logic: If asked how to improve a heterogeneous catalyst, always mention increasing the surface area or using a support medium to maximize the number of available active sites.
