Catalysis: Heterogeneous and Homogeneous

Catalysis: Heterogeneous and Homogeneous

Transition metals and their compounds are highly effective catalysts. Their catalytic ability arises from their ability to form variable oxidation states and (in heterogeneous catalysis) to adsorb reactants onto their surfaces.

What Is a Catalyst?

A catalyst increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy. The catalyst is regenerated at the end of the reaction — it is not consumed.

Important: a catalyst does not change the position of equilibrium or the value of K. It speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but at the same composition.

Heterogeneous Catalysis

A heterogeneous catalyst is in a different phase from the reactants — typically a solid catalyst with gaseous or liquid reactants.

How Heterogeneous Catalysis Works

1. Adsorption — reactant molecules diffuse to the catalyst surface and are adsorbed (they bond to active sites on the surface). This weakens the bonds within the reactant molecules.

2. Reaction — the weakened bonds break and new bonds form between the reactant molecules or fragments at the surface. The activation energy is lower because the surface holds the reactants in favourable orientations and weakens their internal bonds.

3. Desorption — the product molecules leave the surface, freeing the active sites for more reactant molecules.

Key Industrial Examples

The Haber process (iron catalyst):

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

• Iron (Fe) catalyst at 400–500 °C, 200 atm
• N₂ and H₂ molecules adsorb onto the iron surface
• The strong N≡N triple bond is weakened by interaction with the iron
• N and H atoms recombine on the surface to form NH₃
• NH₃ desorbs

The Contact process (vanadium(V) oxide catalyst):

2SO₂(g) + O₂(g) → 2SO₃(g)

• V₂O₅ catalyst at 400–450 °C
• V₂O₅ is also a homogeneous-like catalyst in that it operates via redox cycling:
• SO₂ reduces V₂O₅ to V₂O₄ (V⁵⁺ → V⁴⁺), and SO₂ is oxidised to SO₃
• O₂ reoxidises V₂O₄ back to V₂O₅ (V⁴⁺ → V⁵⁺)
• This demonstrates the role of variable oxidation states

Catalytic converters in cars (Pt, Pd, Rh):

• Convert CO + NO → CO₂ + ½N₂ (and unburnt hydrocarbons → CO₂ + H₂O)
• Platinum, palladium, and rhodium provide the catalytic surface
• The metals are finely divided on a honeycomb support to maximise surface area

Catalyst Poisoning

A catalyst poison is a substance that adsorbs more strongly than the reactants onto the active sites of a heterogeneous catalyst, blocking them permanently or temporarily.

• Lead poisons the platinum/rhodium in catalytic converters (why leaded petrol cannot be used with catalytic converters)
• Sulfur compounds poison iron catalysts in the Haber process (feedstock gases are desulfurised before entering the reactor)

Poisoning reduces catalyst effectiveness and requires costly replacement or regeneration.

Homogeneous Catalysis

A homogeneous catalyst is in the same phase as the reactants — typically all in aqueous solution.

Homogeneous catalysis works by the catalyst forming an intermediate compound with one of the reactants. The catalyst provides a pathway through an intermediate oxidation state, and is regenerated at the end.

Key Feature: Variable Oxidation States

Transition metals catalyse reactions homogeneously because they can change oxidation state, providing a lower-energy route:

Reactant A + Catalyst(state 1) → Intermediate + ...

Intermediate + Reactant B → Product + Catalyst(state 1) regenerated

Example 1: Reaction of S₂O₈²⁻ with I⁻ (Fe²⁺/Fe³⁺ catalyst)

The direct reaction between two negative ions is very slow (electrostatic repulsion):

S₂O₈²⁻(aq) + 2I⁻(aq) → 2SO₄²⁻(aq) + I₂(aq) (very slow uncatalysed)

With Fe²⁺ ions as catalyst:

Step 1: S₂O₈²⁻ + 2Fe²⁺ → 2SO₄²⁻ + 2Fe³⁺

(peroxodisulfate oxidises Fe²⁺ to Fe³⁺)

Step 2: 2Fe³⁺ + 2I⁻ → 2Fe²⁺ + I₂

(Fe³⁺ oxidises iodide to iodine, and is itself reduced back to Fe²⁺)

Overall: S₂O₈²⁻ + 2I⁻ → 2SO₄²⁻ + I₂ (catalyst Fe²⁺ regenerated)

The catalyst works because each step involves a positive ion reacting with a negative ion — there is electrostatic attraction instead of repulsion, reducing the activation energy.

Example 2: Autocatalysis (Mn²⁺ in MnO₄⁻/C₂O₄²⁻)

In the titration of potassium manganate(VII) with oxalic acid:

2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 8H₂O + 10CO₂

The reaction starts slowly (no catalyst). As Mn²⁺ ions are produced, they catalyse the reaction — the rate increases as the reaction proceeds. This is autocatalysis.

The Mn²⁺ acts as an intermediate oxidation state between MnO₄⁻ (Mn⁷⁺) and the final product.

Comparing Heterogeneous and Homogeneous Catalysis

Feature	Heterogeneous	Homogeneous
Phase	Different from reactants	Same as reactants
Mechanism	Surface adsorption	Intermediate compound formation
Recovery	Easy (different phase, filter)	Difficult (same phase, needs separation)
Selectivity	Can be less selective	Often more selective
Poisoning	Susceptible	Less susceptible
Examples	Fe (Haber), Pt/Rh (cat converter)	Fe²⁺/Fe³⁺, Mn²⁺ (autocatalysis)

Enzymes as Biological Catalysts

Enzymes are biological homogeneous catalysts (protein molecules). They operate via the lock-and-key or induced fit model, where the substrate fits into the active site, lowering the activation energy. Transition metal ions are often found at the active sites of enzymes (e.g. Fe²⁺/Fe³⁺ in cytochrome oxidase, Zn²⁺ in carbonic anhydrase).

Exam Tips

• For heterogeneous catalysis, always describe the three stages: adsorption, reaction, desorption
• For homogeneous catalysis, write the two (or more) steps showing the catalyst being used then regenerated, with oxidation states
• When explaining WHY transition metals are good catalysts, cite variable oxidation states (homogeneous) and ability to adsorb reactants (heterogeneous)
• Catalyst poisoning is a common exam question — explain in terms of strong adsorption blocking active sites