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Spec Mapping — OCR H432 Module 3.2.2 — Catalysts, covering the definition of a catalyst as a substance that increases the rate of a reaction by providing an alternative reaction pathway with lower activation energy and that is not consumed overall, the effect of a catalyst on the Maxwell–Boltzmann distribution (overlay diagram with two Ea lines on a single curve) and on the enthalpy-profile diagram (lower peak, same reactants/products, same ΔH), the distinction between heterogeneous (adsorption-mechanism, different phase) and homogeneous (intermediate-mechanism, same phase) catalysts with named industrial examples (refer to the official OCR H432 specification document for exact wording).
A catalyst is the most powerful kinetic lever in chemistry. Where temperature can typically deliver a factor of 2–10 rate increase per 10 K, a well-chosen catalyst routinely delivers 103 to 106. The catalytic converter in a modern car reduces NOx and CO emissions by over 95% at exhaust temperatures, using precious-metal honeycomb that catalyses oxidation and reduction simultaneously. The Haber process produces 150 million tonnes of ammonia per year over an iron catalyst at "only" 450 °C — without that catalyst the reaction would require red heat (~1000 °C) and a vastly more expensive plant. Enzymes — biological catalysts — speed metabolic reactions by factors of 1010 or more at body temperature, making life itself a catalytic phenomenon. Yet a catalyst is never consumed in the net reaction: any catalytic intermediate that forms is regenerated in a subsequent step. This lesson develops the conceptual framework — alternative pathway with lower Ea, Boltzmann-distribution overlay, two-peak enthalpy profile — and surveys the industrially important named catalysts that OCR exams routinely test (Fe for Haber, V2O5 for Contact, Pt/Pd/Rh for catalytic converters, Ni for hydrogenation, zeolites for cracking).
Key Equation: rate enhancement by a catalyst at fixed T is given by the ratio of Boltzmann factors: kuncatkcat=e(Eauncat−Eacat)/RT Lowering Ea by 30 kJ mol−1 at 298 K multiplies the rate by e30000/(8.314×298)≈1.8×105.
A catalyst is a substance that increases the rate of a chemical reaction without itself being used up in the process. It works by providing an alternative reaction pathway with a lower activation energy Ea than the uncatalysed route.
The four key facts that every OCR mark scheme tests:
The mechanism is the Boltzmann distribution from Lesson 7. At any temperature T, a fraction e−Ea/RT of molecules has kinetic energy at least Ea — they can react if they collide effectively. Lowering Ea by, say, 30 kJ mol−1 increases this fraction exponentially:
funcatfcat=e−Eauncat/RTe−Eacat/RT=e(Eauncat−Eacat)/RT
For a typical Eauncat=75 and Eacat=45 kJ mol−1 at 298 K:
funcatfcat=e30000/(8.314×298)=e12.11≈1.8×105
The catalyst makes the reaction proceed ∼180000× faster — five orders of magnitude. This is why catalysis is so dramatic.
This is the most commonly tested diagram in OCR Module 3.2.2. Draw the standard Maxwell–Boltzmann curve (Lesson 7) — origin, peak, asymptotic tail. Then mark two vertical lines on the x-axis:
Shade the area under the curve beyond each line. The catalysed-area is much larger than the uncatalysed-area, demonstrating that a much greater proportion of molecules has enough energy to react when the catalyst is present.
Note carefully: the curve shape and total area are unchanged — only the position of the Ea line on the x-axis has shifted. A common error is to redraw the entire Boltzmann curve; a catalyst affects Ea, not T or the population.
On the energy-profile diagram, a catalysed reaction has a lower peak than the uncatalysed route. Both curves connect the same reactants and products at the same enthalpies, so ΔH is unchanged — only the barrier height changes. Sometimes the catalysed route is split into multiple smaller peaks because the catalyst breaks the reaction into a series of easier elementary steps via catalyst-substrate intermediates.
Both paths start at the same reactant level and finish at the same product level: ΔH is identical. The catalysed pathway has a lower maximum — typically split into two or more smaller peaks corresponding to the formation and decay of catalyst-substrate intermediates.
A heterogeneous catalyst is in a different physical state from the reactants. Most commonly the catalyst is a solid and the reactants are gases or in solution. The classic industrial examples — Fe in the Haber process, V2O5 in the Contact process — are all heterogeneous solids.
The Sabatier principle (1911): the binding strength between catalyst and reactant must be neither too weak (insufficient activation) nor too strong (products cannot escape, "catalyst poisoning"). Optimal catalysts sit at the top of the volcano plot.
| Process | Reaction | Catalyst | Conditions |
|---|---|---|---|
| Haber | N2+3H2⇌2NH3 | Iron (Fe) with K2O promoter | 450 °C, 200 atm |
| Contact | 2SO2+O2⇌2SO3 | Vanadium(V) oxide V2O5 | 450 °C, 1–2 atm |
| Catalytic cracking | Long-chain alkanes → short alkanes + alkenes | Zeolite (alumino-silicate) | 500 °C, slight pressure |
| Catalytic reforming | Cyclohexane → benzene + 3H2 | Pt/Re on alumina | 500 °C, 5 atm |
| Hydrogenation of alkenes | C2H4+H2→C2H6 | Nickel (Ni); finely divided | 150 °C, 5 atm |
| Three-way catalytic converter | 2NO+2CO→N2+2CO2; HCs → CO2 + H2O | Pt, Pd, Rh (precious-metal washcoat on ceramic honeycomb) | Exhaust temperatures (~400–800 °C) |
| Ostwald (HNO3 manufacture) | 4NH3+5O2→4NO+6H2O | Platinum-rhodium gauze | 900 °C, 4–10 atm |
| Methanol synthesis | CO+2H2→CH3OH | Cu/ZnO/Al2O3 | 250 °C, 50–100 atm |
Impurities can bind irreversibly to active sites, poisoning the catalyst. Sulfur compounds poison iron catalysts in the Haber process — which is why the H2 feed must be desulfurised first; lead poisons the precious-metal catalyst in early catalytic converters — which is why all modern petrol is unleaded; carbon monoxide can poison platinum at low temperatures by binding too strongly.
A homogeneous catalyst is in the same physical state as the reactants — typically both dissolved in solution, or both in the gas phase. The mechanism is fundamentally different from heterogeneous catalysis.
The catalyst reacts with one of the reactants to form a catalyst–substrate intermediate, which then reacts further to give products and regenerate the catalyst. Each elementary step has a lower Ea than the uncatalysed pathway:
A+cat→[A-cat]‡→A-cat (intermediate) A-cat+B→P+cat regenerated
The catalyst is chemically identical at the start and end of the cycle, even though it has formally taken part in the mechanism.
Example: acid catalysis of ester hydrolysis. The reaction CH3COOC2H5+H2O→CH3COOH+C2H5OH is catalysed by H+. The protonation of the carbonyl makes the ester more electrophilic; nucleophilic attack by water proceeds at a lower Ea. H+ is regenerated at the end.
| Reaction | Catalyst | Mechanism notes |
|---|---|---|
| Esterification + ester hydrolysis | H2SO4 or HCl (provides H+) | Catalyses both directions equally; Kc unchanged |
| Catalytic decomposition of H2O2(aq) | I−, Fe2+/Fe3+, or MnO2 (heterogeneous variant) | Transition-metal redox cycling |
| Ozone depletion in the stratosphere | Cl• radicals (from photolysis of CFCs) | Each Cl• destroys ~105 O3 before being trapped |
| Autocatalysis of MnO4− + C2O42− | Mn2+ (a product of the reaction) | Reaction self-accelerates as more catalyst is generated |
| Aldol condensation | OH− or H+ | Either acid or base can catalyse different mechanisms |
The Cl-catalysed depletion of stratospheric ozone is the textbook gas-phase homogeneous catalysis example:
Step 1:Cl⋅+O3→ClO⋅+O2 Step 2:ClO⋅+O3→Cl⋅+2O2 Overall:2O3→3O2(Cl⋅ regenerated)
A single chlorine radical can catalyse the destruction of ∼105 ozone molecules before being trapped (typically by reaction with NO2 to form ClONO2, a reservoir species). This is why CFCs — the source of stratospheric Cl• via UV photolysis — were banned under the Montreal Protocol (1987).
| Feature | Heterogeneous | Homogeneous |
|---|---|---|
| Phase of catalyst | Different from reactants | Same as reactants |
| Physical example | Solid + gas/liquid; "Fe + N2/H2" | Aqueous + aqueous; "H+ in ester hydrolysis" |
| Separation from products | Easy — filter, decant, or simply leave in reactor | Difficult — distillation, extraction, ion exchange |
| Mechanism | Adsorption at active sites | Intermediate complex |
| Selectivity | Often shape-selective (zeolites discriminate by pore size) | Often chemistry-selective (acid/base specific) |
| Sensitivity to poisoning | High; surface sites can be permanently blocked | Lower; cation/anion competition is reversible |
| Industrial dominance | Most bulk processes (Haber, Contact, cracking) | Fine chemicals, pharmaceuticals, biotech |
| Examples | Fe (Haber), V2O5 (Contact), Pt/Pd/Rh (catalytic converter) | H2SO4 (esterification), Cl• (stratospheric ozone) |
Heterogeneous catalysts dominate high-volume industrial chemistry because they are easy to separate and reuse, reducing costs and waste. Homogeneous catalysts are preferred when high selectivity (single product, defined stereochemistry) matters more than recovery — typical of pharmaceutical synthesis.
Catalysts provide enormous economic and environmental benefits:
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