Catalysts are one of the most important ideas in IB Chemistry, appearing in kinetics, equilibrium, industrial chemistry, and HL transition metal chemistry. Students often know that catalysts “speed up reactions,” but IB exams require a more precise explanation. This guide breaks down exactly what a catalyst does, how it works, and how to describe it in clear IB-friendly language.
Quick Start Checklist
A catalyst:
- Increases reaction rate
- Lowers activation energy
- Provides an alternative pathway for the reaction
- Is not consumed in the reaction
- Does not change the equilibrium position
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What Does a Catalyst Actually Do?
The key function of a catalyst is to lower the activation energy (Ea) of a reaction.
It does this by providing a new reaction mechanism or pathway that requires less energy for reactants to form products.
This increases the number of successful collisions per unit time, which increases reaction rate.
In IB terms:
“A catalyst increases the rate of reaction by providing an alternative pathway with lower activation energy.”
This exact phrasing earns marks consistently.
How a Catalyst Affects Collisions
1. More particles have sufficient energy
With a lower Ea, a greater fraction of molecules have energy ≥ activation energy.
This dramatically increases the rate of successful collisions.
2. Orientation requirements may be improved
Some catalysts help position reactants in the correct orientation, increasing the chance that a collision produces products.
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Catalysts in Energy Profile Diagrams
On a potential energy diagram:
- The peak representing the transition state is lower with a catalyst.
- The reactants and products stay at the same energy levels.
- Only the activation energy changes.
This diagram helps students understand why catalysts affect rate but not the thermodynamic favorability of a reaction.
Catalysts and Equilibrium
A catalyst does not change the position of equilibrium.
It only allows the system to reach equilibrium faster.
This is because the catalyst increases both forward and reverse reaction rates equally.
IB examiners frequently include this in multiple-choice and short-answer questions.
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Types of Catalysts
1. Homogeneous catalysts
- Same phase as reactants
- Often form intermediates
- Common in acid–base catalysis
2. Heterogeneous catalysts
- Different phase from reactants
- Often solid catalysts with gaseous or aqueous reactants
- Widely used in industry (e.g., Fe in the Haber process)
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Frequently Asked Questions
1. Does a catalyst get used up in a reaction?
No. A catalyst is chemically unchanged at the end of a reaction. It may temporarily form intermediates, but it is regenerated overall. IB always tests this idea.
2. Does a catalyst make a reaction more exothermic or endothermic?
No. Enthalpy change (ΔH) is unchanged. Catalysts affect kinetics, not thermodynamics. This distinction is essential for SL and HL students.
3. Can enzymes be considered catalysts?
Yes—enzymes are biological catalysts. They lower activation energy by binding to substrates and stabilizing transition states. They are an excellent example of homogeneous catalysis.
Conclusion
A catalyst speeds up a reaction by lowering its activation energy and providing an alternative pathway, while remaining chemically unchanged at the end. It increases the number of successful collisions and helps the reaction reach equilibrium more quickly without altering the equilibrium position. Mastering catalyst theory is essential for IB Chemistry kinetics, equilibria, and industrial applications. With RevisionDojo’s chemistry-focused support, mastering these ideas becomes clearer and more manageable.
