Part 5: Surface catalysis, catalytic cycles and autocatalysis

Explain how a catalyst provides a lower-activation-energy route, show how it is regenerated, and interpret why an autocatalytic reaction first speeds up then slows down.

A catalyst changes the pathway, not the equilibrium

A catalyst increases reaction rate by providing an alternative pathway with lower activation energy and is regenerated overall. It may react during an individual step, so ‘does not take part in the reaction’ is misleading. Transition metals can bind reactants at surfaces or change oxidation state in solution, making alternative pathways available.

A heterogeneous catalyst is in a different phase from the reactants and reaction occurs at its surface. A homogeneous catalyst is in the same phase as the reactants and the pathway proceeds through intermediate species. Phase means a physically distinct region: two immiscible liquids are different phases even though both are liquids.

A catalyst does not change ΔH, ΔG°, E°cell, K or the equilibrium composition at a fixed temperature. It speeds approach to equilibrium by lowering barriers in both directions. It can permit an industrial process to run usefully at a lower temperature, but that changed operating temperature, not catalysis itself, changes equilibrium.

Adsorb, react, then desorb

CO and NO molecules adsorb onto active sites of a catalytic converter’s metal surface: bonds form between reactants and surface atoms. This changes electron distribution and weakens bonds within reactants, while holding reactants close in suitable orientations. At the surface, bonds rearrange to form CO₂ and N₂. These products desorb, freeing active sites for more molecules.

Adsorption is attachment at a surface; absorption means entry into the bulk of a material. A catalyst must bind reactants strongly enough to activate them but allow products to leave. A poison occupies active sites strongly and blocks their use. A high-area honeycomb support coated with small quantities of platinum-group metals exposes many sites without requiring a large mass of expensive metal.

The converter lowers carbon monoxide and nitrogen monoxide emissions through 2CO + 2NO → 2CO₂ + N₂. Carbon monoxide is oxidised and nitrogen monoxide reduced. Carbon dioxide is less acutely toxic than CO but remains a greenhouse gas, so ‘all pollutants become harmless’ is an overstatement. A cold converter works less effectively because the catalytic reaction still has an activation barrier.

Vanadium(V) oxide is reduced and regenerated

In the Contact process, V₂O₅ catalyses oxidation of SO₂ to SO₃. A simplified redox cycle shows V(V) in V₂O₅ reduced to V(IV) in V₂O₄ while sulfur(IV) becomes sulfur(VI). Oxygen then reoxidises V₂O₄ to V₂O₅. The catalyst’s oxidation state changes during the cycle and returns to its starting value.

Double the first equation before adding to the second. Two V₂O₅ consumed and two regenerated cancel, as do two V₂O₄ formed and consumed. The net equation is 2SO₂ + O₂ → 2SO₃. This cancellation is the evidence for regeneration; a catalytic scheme that leaves the catalyst permanently consumed is incomplete.

This is the required school-level oxide cycle. It links catalytic behaviour to accessible vanadium oxidation states; it is not a complete mechanistic description of every species in a working industrial catalyst. Use the model to account for atoms and electron transfer.

Iron(II) carries electrons between two reacting anions

The overall reaction S₂O₈²⁻ + 2I⁻ → 2SO₄²⁻ + I₂ is slow without a catalyst partly because the reacting ions both have negative charge and repel on approach. Dissolved Fe²⁺ provides a different route: peroxodisulfate oxidises Fe²⁺ to Fe³⁺, then Fe³⁺ oxidises iodide and regenerates Fe²⁺. All reactants and catalyst are in the aqueous phase, so this is homogeneous catalysis.

Each stage involves attraction between oppositely charged ions and a different electron-transfer pathway. Avoid treating electrostatic attraction alone as a universal proof of a lower barrier; the observed catalysis and viable intermediate cycle support the mechanism. Fe³⁺ can also initiate the same cycle because the steps can begin in the other order when the necessary reactants are present.

Using E°(S₂O₈²⁻/SO₄²⁻) = +2.01 V, E°(Fe³⁺/Fe²⁺) = +0.77 V and E°(I₂/I⁻) = +0.54 V gives +1.24 V for the first step and +0.23 V for the second. The iron potential lies between the two reactant couples, making both redox stages thermodynamically feasible. This does not by itself prove a fast catalytic rate: kinetics must still permit the steps.

A reaction product can accelerate its own formation

Acidified manganate(VII) oxidises ethanedioate to carbon dioxide and is reduced to Mn²⁺. Initially the uncatalysed reaction is slow. As Mn²⁺ accumulates, it provides a faster route through Mn³⁺ intermediates, so the rate increases. Because the catalyst is also a product of the overall reaction, this is autocatalysis.

One simplified sequence is 4Mn²⁺ + MnO₄⁻ + 8H⁺ → 5Mn³⁺ + 4H₂O, followed by 2Mn³⁺ + C₂O₄²⁻ → 2Mn²⁺ + 2CO₂. Multiply the first by two and the second by five before adding. Ten Mn³⁺ cancel, and ten Mn²⁺ are regenerated after eight were used: two Mn²⁺ remain as net products, matching the overall reaction.

The rate eventually falls because reactants are depleted, despite the high catalyst concentration. For a plot of [MnO₄⁻] against time, expect an initially shallow downward gradient, then a steeper descent, then a shallow approach to completion. The magnitude of the gradient is the rate of manganate consumption. Adding a small known amount of Mn²⁺ at the start shortens or removes the initial slow period.

Follow the purple ion’s absorbance at a suitable wavelength using a blank and a calibration where needed. Keep temperature, starting concentrations, total volume and optical path length controlled when comparing runs. Absorbance can follow concentration without withdrawing samples, but bubbles, dirty cuvettes and excessively high absorbance can distort results. Label a drawn curve as schematic unless it is backed by actual measurements.

Build a connected comparison

To compare platinum in a converter with aqueous Mn²⁺, start with the shared lower-activation-energy route and regeneration. Then explain the physical distinction: gases interact at a solid surface, whereas dissolved ions react through soluble intermediates. Link adsorption, weakened bonds and desorption in the first case; link product formation, rising catalyst concentration and acceleration in the second.

June 2023 9CH0/01 Q8 assessed this comparison through linked reasoning as well as content. Its report highlights confusion between adsorption and absorption and omitted explanations of the initial slow anion–anion reaction. A list of unrelated catalyst facts does not show how the contrasting mechanisms produce the observed behaviour.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources