Key Takeaways
- A catalyst increases the rate of a chemical reaction without being consumed; its mass and chemical properties are unchanged before and after the reaction.
- It works by providing an alternative reaction path with a lower activation energy; it does not change the equilibrium or the Gibbs free-energy change of the reaction.
- Catalysts are classified as homogeneous or heterogeneous by the phase of the reaction system, and as liquid or solid by their state.
- Key industrial examples: iron catalyst for ammonia synthesis, vanadium catalyst for sulfuric acid, and zeolite or metal catalysts in petroleum refining and automotive exhaust treatment.
1. Definition
A catalyst generally refers to a substance that increases the rate of a reaction without changing the total standard Gibbs free-energy change of the reaction. Equivalently, it increases the rate of a chemical reaction without changing the chemical equilibrium, and its mass and chemical properties remain unchanged before and after the reaction. The catalyst works by offering an alternative pathway with a lower activation energy: more reactant molecules have enough energy to react per unit time, so the reaction is faster while the thermodynamics of the overall reaction stay the same.
2. Classification
| Catalyst Type | Description | Examples |
|---|---|---|
| Homogeneous Catalysts | The catalyst is in the same phase as the reactants, usually in the liquid phase. | Acids, bases, soluble transition-metal compounds, peroxide catalysts |
| Heterogeneous Catalysts | The catalyst is in a different phase from the reactants, typically a solid catalyst interacting with liquid or gas reactants. | Iron for ammonia synthesis, vanadium pentoxide for sulfuric acid production, zeolites and metal oxides for petroleum cracking, three-way automotive catalysts |
Classification by Physical State: Catalysts can also be classified as liquid catalysts or solid catalysts based on their physical state. However, the distinction between homogeneous and heterogeneous catalysts is determined by whether the catalyst and the reaction system are in the same or different phases.
3. Key Properties
A useful catalyst is judged by three properties: activity (how fast it accelerates the reaction), selectivity (how much it favours the desired product over by-products) and stability (how long it keeps its activity in service). Industrial catalysts are often supported on high-surface-area carriers to maximise the active surface per unit volume.
4. Industrial Importance
It is commonly estimated that more than 90% of industrial chemical processes use catalysts - in chemical, petrochemical, biochemical and environmental applications. Well-known examples: iron-based catalysts in ammonia production (Haber–Bosch process), vanadium catalysts in sulfuric acid production (contact process), and various catalysts for ethylene polymerisation and the production of rubber from butadiene. Catalysts also underpin petroleum refining (fluid catalytic cracking) and automotive exhaust emission control (three-way catalysts).
Frequently Asked Questions
Q1. Is a catalyst consumed in the reaction?
No - its mass and chemical properties are unchanged before and after the reaction; it is regenerated as part of the catalytic cycle.
Q2. Does a catalyst change the equilibrium?
No. It speeds up both directions equally and only shortens the time to equilibrium.
Q3. What is the difference between homogeneous and heterogeneous catalysis?
Homogeneous catalysts are in the same phase as the reactants; heterogeneous catalysts are in a different phase, usually a solid with liquid or gas reactants.
Q4. What is the role of vanadium in sulfuric acid production?
Vanadium pentoxide catalyses the oxidation of sulfur dioxide to sulfur trioxide in the contact process.