CatalystsEdexcel A-Level Chemistry: Revision notes

Section 1

What a catalyst does

A catalyst is a substance that increases the rate of a reaction by providing an alternative reaction route with a lower activation energy. It is not used up: it is chemically unchanged at the end, although it may be physically changed.

Because the activation energy is lower, a greater proportion of collisions are successful at the same temperature. A catalyst does not change the enthalpy change of the reaction, the position of the products' energy, or the amount of product formed.

Key termscatalystalternative route

Common mistake

Saying a catalyst lowers the activation energy 'of the reaction' without mentioning an alternative route, or saying it makes the reaction more exothermic. It does neither: ΔH is unchanged.

Section 2

Reaction profiles

A reaction profile plots energy against progress of reaction. The reactants start at one energy level and the products finish at another. The difference is the enthalpy change ΔH (products lower for exothermic, higher for endothermic). The peak above the reactants is the activation energy.

For a catalysed reaction, draw a second curve with a lower peak between the same start and finish levels. If a catalyst forms an intermediate, the catalysed profile may show two small peaks, each lower than the uncatalysed peak.

For the reverse reaction, $E_a(\text{reverse}) = E_a(\text{forward}) - \Delta H$ . A catalyst lowers the forward and reverse activation energies by the same amount.

Example: ΔH = −60, forward Ea = 120, reverse Ea = 120 + 60 = 180 kJ mol⁻¹.

Key termsreaction profileenthalpy change

Exam tip

Label the axes (energy, progress of reaction), the reactant and product levels, ΔH, and Ea for each pathway. Start and end points for both curves must be the same.

Section 3

Solid (heterogeneous) catalysts

A heterogeneous catalyst is in a different phase from the reactants. Many industrial gas-phase reactions use a solid catalyst: iron in the Haber process, vanadium(V) oxide in the Contact process, and platinum and rhodium in catalytic converters.

The reaction takes place on the surface of the solid:

Reactant molecules are adsorbed onto active sites on the surface.
Bonds in the adsorbed molecules are weakened, and they react by the alternative route.
Product molecules desorb, freeing the active sites for more reactants.

The catalyst is often made as small pieces, a porous solid or a thin coating on a support, giving a large surface area and many active sites.

Key termsheterogeneous catalystadsorptionactive site

Common mistake

Saying the reactants are 'absorbed' into the catalyst. They are adsorbed onto the surface.

Section 4

Economic benefits in industry

Catalysts matter because many reactions would be far too slow at a sensible temperature.

A lower activation energy means the reaction is fast enough at a lower temperature, which saves energy and fuel costs.
Lower temperatures (and pressures) mean cheaper plant and less demanding equipment.
Faster rates produce more product per hour, so the cost per tonne falls.
Catalysts are not used up, so a small amount can be used for a long time before replacing.
Less fuel burned also means less carbon dioxide released.

The benefits are weighed against the cost of the catalyst (some contain platinum or rhodium) and the need to replace or recycle it eventually.

Key termseconomic benefits

Exam tip

In 'explain the economic benefits', link each point to a reason: lower Ea, so lower temperature, so less energy cost.

Must Know

A catalyst gives an alternative route with a lower activation energy and is not used up
It does not change ΔH or the products formed
Reaction profile: same start and finish levels, lower peak for catalysed route
Heterogeneous catalyst: different phase; reaction on the surface by adsorption and desorption
Large surface area gives more active sites
Economic benefits: lower temperature, less energy, more product per hour, catalyst reused

That's the notes covered.

Carry on to the next subtopic.

Exam questions on Catalysts

Ammonia is made industrially by the Haber process, in which nitrogen and hydrogen gases pass over a solid iron catalyst: N₂(g) + 3H₂(g) → 2NH₃(g).
Describe how a solid catalyst such as iron provides a surface for the reaction of gaseous reactants.2 marks
A reaction A → B has an enthalpy change of −60 kJ mol⁻¹ and, uncatalysed, an activation energy of 120 kJ mol⁻¹. When a solid catalyst is used the activation energy falls to 70 kJ mol⁻¹.
Describe how the reaction profile for the catalysed reaction differs from the uncatalysed reaction, and state what is the same in the two.2 marks
In the Contact process, sulfur dioxide and oxygen gases pass over a solid vanadium(V) oxide catalyst at about 700 K: 2SO₂(g) + O₂(g) → 2SO₃(g), ΔH = −196 kJ mol⁻¹. In a simplified model the uncatalysed reaction has an activation energy of 250 kJ mol⁻¹ and the catalysed reaction has an activation energy of 100 kJ mol⁻¹.
Explain how a solid catalyst speeds up the reaction between two gases, and why the catalyst is used as small pellets or spread over a support.3 marks

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Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).