3.5.2 Catalysis

3.5.2 Catalysis: Explain how a catalyst speeds up a reaction by providing an alternative pathway with a lower activation energy, and distinguish homogeneous from heterogeneous catalysis.

A catalyst changes how fast equilibrium is reached, not where the energy of the reaction ends up.

Alternative route

A catalyst offers a different reaction pathway that has a lower activation energy than the uncatalysed route.

Lower Eₐ

With a smaller energy barrier, a larger fraction of colliding particles has enough energy to react at the same temperature.

Faster rate

More successful collisions happen every second, so the rate of reaction increases without any rise in temperature.

Not used up

The catalyst is regenerated at the end of the reaction, so it is chemically unchanged and appears in no overall equation.

No change to ΔH

Reactants and products keep the same energies, so the enthalpy change of the reaction is identical with or without the catalyst.

Homogeneous and Heterogeneous

Homogeneous

The catalyst is in the same physical state as the reactants, for example aqueous H⁺ ions catalysing the formation of an ester in solution.

Heterogeneous

The catalyst is in a different state from the reactants, usually a solid with gaseous or liquid reactants flowing over it.

Adsorption

In heterogeneous catalysis, reactant molecules bond weakly to the surface, which weakens bonds and brings them together.

Desorption

After reacting on the surface the product molecules leave, freeing the active sites for the next reactant molecules.

Process	Catalyst	Type
Haber process, N₂ + 3H₂ ⇌ 2NH₃	Iron	Heterogeneous
Contact process, 2SO₂ + O₂ ⇌ 2SO₃	Vanadium(V) oxide, V₂O₅	Heterogeneous
Car exhaust clean-up	Platinum, palladium, rhodium	Heterogeneous
Esterification of an acid with an alcohol	Concentrated H₂SO₄ (H⁺)	Homogeneous

Draw the energy distribution of the molecules at one temperature, with the uncatalysed activation energy marked on the energy axis.

Mark a second, smaller activation energy further left to represent the catalysed pathway on the same graph.

The curve itself does not move, because the temperature and the number of molecules are unchanged.

The area to the right of the catalysed Eₐ is much larger, so a greater proportion of molecules can react.

A greater proportion of effective collisions gives a faster rate, which is why catalysts are so useful in industry.

Lower temperature

Reactions run quickly at lower temperatures, which cuts the fuel needed and the carbon dioxide released from burning it.

Lower cost

Less energy and shorter reaction times reduce the running cost and make the process more sustainable.

Reusable

A catalyst can be used repeatedly, although impurities can poison its surface and reduce its effectiveness over time.

Surface area

Heterogeneous catalysts are often spread on a support or made into a fine mesh to expose as much surface as possible.

3.5.2 – Key Points for the Exam

Catalyst: A substance that increases the rate of a reaction by providing an alternative route of lower activation energy and is not used up.
Activation energy: The minimum energy that colliding particles must have for a reaction to occur.
Homogeneous catalyst: A catalyst in the same physical state as the reactants.
Heterogeneous catalyst: A catalyst in a different physical state from the reactants, commonly a solid.
Adsorption: The attachment of molecules to the surface of a solid catalyst.

Common Mistakes

Saying a catalyst increases the energy of the particles; it lowers the energy barrier instead.
Stating that a catalyst is used up or changes the equilibrium yield; it speeds up forward and reverse reactions equally.
Moving the peak of the Boltzmann curve when a catalyst is added; only the activation energy marker moves.
Forgetting to say 'alternative pathway' when answering a question that asks how a catalyst works.