Catalysis

A catalyst is a substance that increases the rate of a chemical reaction without being consumed in the overall process. It works by providing an alternative reaction pathway with a lower activation energy (Ea). This is one of the most practically important concepts in chemistry -- from industrial manufacturing to biological systems, catalysts are everywhere.

How Catalysts Work

Without a catalyst, reactant molecules must overcome the full activation energy barrier. A catalyst creates a different pathway -- typically involving the formation of an intermediate species -- that has a lower energy barrier. Since more particles now have energy ≥ Ea, the rate increases.

Key characteristics of all catalysts:

Not consumed -- they are regenerated at the end of the reaction and can be reused.
Lower Ea -- they provide an alternative pathway, not extra energy to the reactants.
Do not change ΔH -- the enthalpy change is the same with or without a catalyst (it depends only on the energy of reactants and products).
Do not change the equilibrium position -- they speed up both the forward and reverse reactions equally, so the equilibrium is reached faster but the position is unchanged.
Do not appear in the overall equation -- though they appear in the mechanism.

Quantifying the Catalytic Effect

Using the Arrhenius equation, we can see why lowering Ea has such a dramatic effect:

At 300 K, for a reaction with Ea = 100 kJ mol⁻¹:

e^(−Ea/RT) = e^(−100000/(8.314 × 300)) = e^(−40.1) = 3.9 × 10⁻¹⁸

With a catalyst reducing Ea to 60 kJ mol⁻¹:

e^(−Ea/RT) = e^(−60000/(8.314 × 300)) = e^(−24.1) = 3.4 × 10⁻¹¹

The Boltzmann factor has increased by a factor of about 10⁷ -- ten million times more particles can now react! This illustrates why catalysts are so effective.

Heterogeneous Catalysis

A heterogeneous catalyst is in a different phase from the reactants. Most commonly, the catalyst is a solid and the reactants are gases or liquids.

The Mechanism: Adsorption

Heterogeneous catalysis typically involves the following steps:

Diffusion -- reactant molecules diffuse to the catalyst surface.
Adsorption -- reactant molecules bond to the catalyst surface at active sites. This weakens the bonds within the reactant molecules.
Reaction -- with weakened bonds and correct orientation (enforced by the surface geometry), the reaction occurs at the surface with lower Ea.
Desorption -- product molecules leave the catalyst surface.
Diffusion -- product molecules diffuse away from the surface.

The adsorption step is critical: by bonding to the surface, the intramolecular bonds of the reactants are weakened, effectively lowering the energy needed to break them and form new bonds.

Examples of Heterogeneous Catalysis

Process	Reaction	Catalyst	Key detail
Haber Process	N₂ + 3H₂ ⇌ 2NH₃	Iron (+ Al₂O₃, K₂O promoters)	Weakens N≡N triple bond
Contact Process	2SO₂ + O₂ ⇌ 2SO₃	V₂O₅	Redox cycle: V₂O₅ → V₂O₄ → V₂O₅
Catalytic converter	2CO + O₂ → 2CO₂	Pt, Pd, Rh on honeycomb	Large surface area critical
Hydrogenation	R—CH=CH—R + H₂ → alkane	Nickel (Ni)	Used in margarine production
Ostwald Process	4NH₃ + 5O₂ → 4NO + 6H₂O	Platinum-rhodium gauze	Step in HNO₃ manufacture

Homogeneous Catalysis

A homogeneous catalyst is in the same phase as the reactants. The catalyst typically reacts with one reactant to form an intermediate, which then reacts with another reactant to regenerate the catalyst.

The Mechanism: Intermediate Formation

Step 1: Catalyst + Reactant A → Intermediate

Step 2: Intermediate + Reactant B → Product + Catalyst

The catalyst is consumed in Step 1 but regenerated in Step 2. The intermediate provides the alternative pathway with lower Ea. Both steps have lower activation energies than the uncatalysed single-step reaction.

Energy Profile Comparison

For an uncatalysed reaction, there is one large energy barrier. For a homogeneously catalysed reaction, the energy profile shows two smaller "humps" with an intermediate energy well between them:

graph LR
    A["Reactants"] --> B["Transition state 1<br/>(lower Ea)"]
    B --> C["Intermediate<br/>(energy well)"]
    C --> D["Transition state 2<br/>(lower Ea)"]
    D --> E["Products"]

The maximum energy reached in the catalysed pathway is lower than the single transition state in the uncatalysed pathway.

Example: Acid-Catalysed Ester Hydrolysis

The hydrolysis of an ester is catalysed by H⁺ ions (acid catalyst). The H⁺ is in aqueous solution, the same phase as the ester and water -- homogeneous catalysis. The H⁺ protonates the carbonyl oxygen, making the ester carbon more susceptible to nucleophilic attack by water.

Example: Persulfate-Iodide Reaction Catalysed by Fe²⁺/Fe³⁺

S₂O₂²⁻(aq) + 2I⁻(aq) → 2SO₄²⁻(aq) + I₂(aq)

Without catalyst: two negatively charged ions must collide (electrostatic repulsion makes this slow).

With Fe²⁺ catalyst:

Step 1: S₂O₂²⁻ + 2Fe²⁺ → 2SO₄²⁻ + 2Fe³⁺ (Fe²⁺ oxidised)

Step 2: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂ (Fe²⁺ regenerated)

Each step involves oppositely charged ions attracting each other, making collision much more likely.

Example: Atmospheric Ozone Destruction

Cl• (from CFCs) catalyses the destruction of ozone:

Step 1: Cl• + O₃ → ClO• + O₂

Step 2: ClO• + O₃ → Cl• + 2O₂

Overall: 2O₃ → 3O₂

Cl• is regenerated -- it is a homogeneous catalyst (gas phase, same as the ozone). A single chlorine radical can destroy thousands of ozone molecules before being removed.

Autocatalysis

In autocatalysis, a product of the reaction acts as a catalyst for the reaction itself. This means the reaction starts slowly (little catalyst present), speeds up as more product/catalyst is formed, and then slows again as reactants are consumed.

Example: MnO₄⁻ with C₂O₄²⁻

The reaction between potassium manganate(VII) and ethanedioic acid in acidic solution:

2MnO₄⁻(aq) + 5C₂O₄²⁻(aq) + 16H⁺(aq) → 2Mn²⁺(aq) + 10CO₂(g) + 8H₂O(l)

The Mn²⁺ ions produced catalyse the reaction. Initially, the reaction is slow (the purple MnO₄⁻ decolourises slowly). As Mn²⁺ accumulates, the rate increases, and the decolourisation becomes noticeably faster.

On a rate-time graph, autocatalysis shows an initial increase in rate (unlike normal reactions where rate decreases from the start), followed by a peak, then the usual decline as reactants are consumed.

Enzymes as Biological Catalysts

Enzymes are protein catalysts found in living systems. They are highly specific -- each enzyme catalyses one particular reaction or type of reaction. Enzymes operate by a lock-and-key or induced-fit mechanism:

The substrate binds to the enzyme's active site.
The active site has a specific shape complementary to the substrate.
The enzyme-substrate complex lowers the activation energy.
Products are released and the enzyme is regenerated.

Enzymes are homogeneous catalysts (both enzyme and substrate are typically dissolved in aqueous solution in cells).

Property	Chemical catalyst	Enzyme
Specificity	Low to moderate	Very high
Operating temperature	Often 200--500 °C	25--40 °C
Ea reduction	Moderate	Large (often Ea < 40 kJ mol⁻¹)
Denaturation	Generally stable	Denatured at high T or extreme pH
Availability	Mined/synthesised	Produced by living cells

Catalyst Poisoning

A catalyst poison is a substance that permanently bonds to the active sites of a heterogeneous catalyst, blocking them from adsorbing reactants. This reduces catalytic activity.

Common examples:

Lead poisons the Pt/Pd/Rh in catalytic converters (which is why leaded petrol cannot be used with catalytic converters).
Sulfur compounds poison the iron catalyst in the Haber process (so sulfur must be removed from the N₂/H₂ feedstock).
Carbon monoxide poisons platinum catalysts used in fuel cells.

Catalyst poisoning is distinct from a reversible inhibitor -- a poison binds irreversibly to the surface.

Common Mistakes to Avoid

"A catalyst gives particles more energy." Wrong. A catalyst lowers the energy barrier, not raises the energy of particles.
"A catalyst shifts the equilibrium to the right." Wrong. A catalyst speeds up both forward and reverse reactions equally; the equilibrium position is unchanged.
"A catalyst is used up in the reaction." Wrong. A catalyst is regenerated. It may be consumed in one step but is reformed in a later step.
Confusing heterogeneous and homogeneous. Heterogeneous = different phase; homogeneous = same phase. An aqueous catalyst with aqueous reactants is homogeneous, even if it looks like a different substance.
"Enzymes are not catalysts." They are biological catalysts. All the rules of catalysis apply.

Summary

Catalysts increase reaction rates by providing alternative pathways with lower Ea. Heterogeneous catalysts (different phase) work via adsorption at active sites on a surface. Homogeneous catalysts (same phase) form intermediates that provide a lower-energy route. Autocatalysis occurs when a product catalyses its own formation. Key industrial examples include iron in the Haber process, V₂O₅ in the Contact process, and Pt/Pd/Rh in catalytic converters. Catalysts do not change ΔH, do not shift equilibrium position, and are not consumed overall.

A-Level Deep Dive: Catalysis

Spec mapping

Edexcel 9CH0 specification, Topic 9 — Kinetics I and Topic 16 — Kinetics II require candidates to define a catalyst as a substance that increases the rate of reaction without being consumed, to distinguish homogeneous from heterogeneous catalysis, and to explain that catalysts work by providing an alternative pathway with a lower activation energy (refer to the official specification document for exact wording). Topic 11 — Equilibrium II — explicitly states that catalysts have no effect on the position of equilibrium, only on the time taken to reach it. Topic 19 (transition metals) covers heterogeneous catalysts including Fe in Haber, V₂O₅ in Contact, Ni in hydrogenation, and Pt-Rh in catalytic converters. Examined on Papers 1, 2 and 3. The data booklet does not list standard catalyst examples — students must memorise the major industrial catalysts.

Worked example with full mark scheme

Question (8 marks):

Hydrogen peroxide decomposes very slowly at room temperature. Adding solid manganese(IV) oxide $\text{MnO}_2$ produces vigorous gas evolution; adding aqueous $\text{Mn}^{2+}$ to acidic peroxide also accelerates decomposition.

(a) Classify each of these as homogeneous or heterogeneous catalysis. (2)

(b) Sketch a reaction profile showing the uncatalysed and catalysed pathways for the decomposition. (2)

(c) Explain why a catalyst does not shift the equilibrium position of a reversible reaction, even though it increases the forward rate. (4)

Solution with mark scheme:

(a) M1 — $\text{MnO}_2(\text{s})$ + $\text{H}_2\text{O}_2(\text{aq})$ : heterogeneous (different phases, solid catalyst with aqueous reactant). A1 — $\text{Mn}^{2+}(\text{aq})$ + $\text{H}_2\text{O}_2(\text{aq})$ : homogeneous (same phase, both aqueous).

(b) M1 — sketch with energy on y-axis, reaction coordinate on x-axis; reactants at higher energy than products (decomposition is exothermic, $\Delta H < 0$ ). A1 — uncatalysed pathway with high $E_a$ peak; catalysed pathway with lower $E_a$ peak (potentially with intermediate(s) shown as small minima for the catalysed route — the heterogeneous mechanism involves adsorption + reaction + desorption). $\Delta H$ is the same in both pathways; only the barrier height differs.

(c) M1 — at equilibrium, forward and reverse rates are equal: $k_f [A] = k_r [B]$ . M1 — a catalyst lowers the activation energy of the forward path and the reverse path equally — the same transition state is shared. M1 — therefore both $k_f$ and $k_r$ increase by the same factor, leaving $k_f / k_r = K_c$ unchanged. A1 — the system reaches equilibrium faster, but the position of equilibrium (concentration ratio) is determined by $\Delta G^\circ = -RT \ln K$ , which depends only on thermodynamics — not on the kinetics of either pathway.

Total: 8 marks (M5 A3, split as shown).

Specimen question modelled on the Edexcel 9CH0 Paper 2 format

Question (6 marks): Catalytic converters use platinum, palladium and rhodium to convert exhaust gases (CO, NO, unburned hydrocarbons) into less harmful products (CO₂, N₂, H₂O).

(a) Write a balanced equation for the reaction between CO and NO catalysed in a catalytic converter. (2)

(b) Describe the four-step heterogeneous catalysis mechanism on the metal surface (adsorption, reaction, desorption, etc.). (3)

Mark scheme decomposition by AO: