Catalysis On Surfaces Explained For JEE And NEET

Catalysis on surfaces, also called heterogeneous catalysis, happens when a solid catalyst speeds up a reaction between gases or liquids on its surface. Reactant molecules first stick (adsorb) onto the solid, react while held there, and then the products detach (desorb), leaving the surface free for the next round. Examples: iron in the Haber process, nickel in oil hydrogenation, platinum in exhaust converters.

Nearly all large-scale chemical manufacturing depends on surface catalysis. For JEE and NEET, it explains how catalysts lower activation energy, what active sites and poisoning mean, and why finely divided catalysts work best.

Key Concepts

An active site is a specific spot on the catalyst surface, often a step, kink, or defect, where reactants bind and react. The Langmuir-Hinshelwood mechanism describes reactions where both reactants adsorb before reacting on the surface, while the Eley-Rideal mechanism has one adsorbed reactant struck by another arriving directly from the gas phase. The Sabatier principle states that the surface-reactant bond must be neither too weak (reactants will not stick) nor too strong (products will never leave), the best catalysts sit in the middle. Catalyst poisoning occurs when an impurity like sulphur or carbon monoxide binds irreversibly to active sites and kills activity. Turnover frequency (TOF), measured in s-1, counts how many reaction cycles one active site completes per second.

Formulas

TOF = (moles of reactant converted per second) / (moles of active sites), with units of s-1. The reaction rate on a surface is proportional to the fraction of surface covered by reactants, θ, so rate ∝ θ, which leads to the Langmuir isotherm at equilibrium. Effectiveness scales with surface area, which is why catalysts are fine powders or on porous supports. These relations assume the surface reaction step is rate-determining and the catalyst is not deactivating.

Worked Example

A platinum catalyst contains 2.0 × 10-3 mol of active sites and converts 0.60 mol of reactant per second. What is the turnover frequency, and how many cycles does one site complete in an hour? TOF = 0.60 / (2.0 × 10-3) = 300 s-1. Each active site therefore finishes 300 reaction cycles every second. In one hour there are 3600 seconds, so cycles per site = 300 × 3600 = 1,080,000, about 1.1 × 106 turnovers. This shows why tiny amounts of catalyst process huge quantities of reactants, while the surface stays clean.

Common Misconceptions

• The catalyst gets used up in the reaction. It does not. The catalyst is regenerated at the end of every cycle when products desorb, which is why small quantities suffice.
• A catalyst changes the equilibrium of a reaction. It cannot. A catalyst speeds up the forward and backward reactions equally, so equilibrium position and yield stay the same; only the time to reach equilibrium shrinks.
• Adding more catalyst always speeds things up proportionally. Only up to a point. Once the surface is saturated with reactants or poisoned by impurities, extra catalyst gives diminishing returns.

Real-World Applications

The Haber process uses iron to make ammonia for fertilisers; the Contact process uses vanadium(V) oxide to make sulphuric acid. Catalytic converters use platinum, palladium, and rhodium to turn CO, unburnt fuel, and nitrogen oxides into CO2, water, and nitrogen. Ziegler-Natta catalysts polymerise ethylene and propylene into plastics on solid surfaces. Fuel cells depend on platinum-catalysed surface reactions.

Surface catalysis is a cycle of adsorb, react, and desorb on active sites, governed by the Sabatier principle and measured by turnover frequency.