Catalysts (OCR A-Level Chemistry A): Revision Notes

Catalysts

What does a catalyst do?

A catalyst is a substance that increases the rate of a chemical reaction while remaining chemically unchanged at the end of the reaction. This means the catalyst is not permanently consumed or altered during the reaction process.

Key characteristics of catalysts:

• The catalyst is not used up during the reaction - it is regenerated at the end
• The catalyst may interact with a reactant to form a temporary intermediate species, or it may provide a surface where the reaction occurs
• Only small amounts of catalyst are needed because it can be used repeatedly

How catalysts work

Catalysts speed up reactions by providing an alternative reaction pathway that has a lower activation energy than the uncatalysed reaction. The activation energy ($E_a$) is the minimum energy required for reactant particles to successfully collide and form products.

By lowering the activation energy, more reactant particles have sufficient energy to overcome this barrier at a given temperature. This means more successful collisions occur per second, increasing the reaction rate.

Enthalpy profile diagrams

The effect of a catalyst can be shown using enthalpy profile diagrams, which plot energy against the progress of reaction.

In these diagrams:

• The uncatalysed reaction pathway (shown by the higher curve) has a higher activation energy ($E_a$)
• The catalysed reaction pathway (shown by the lower curve) has a lower activation energy ($E_c$)
• The enthalpy change ($\Delta H$) remains the same for both pathways
• For exothermic reactions, products are at a lower energy level than reactants ($\Delta H$ is negative)
• For endothermic reactions, products are at a higher energy level than reactants ($\Delta H$ is positive)

Types of catalyst

Catalysts are classified into two main types based on their physical state relative to the reactants: homogeneous and heterogeneous.

Homogeneous catalysts

A homogeneous catalyst exists in the same physical state as the reactants. For example, if the reactants are gases, the catalyst is also a gas; if the reactants are in solution, the catalyst is also dissolved in the same solution.

Examples of homogeneous catalysis:

Heterogeneous catalysts

A heterogeneous catalyst has a different physical state from the reactants. Typically, heterogeneous catalysts are solids that catalyse reactions involving gaseous or liquid reactants.

The catalyst provides a large surface area where reactant molecules can come together and react more easily. This lowers the activation energy for the reaction.

Industrial processes using heterogeneous catalysts

Many important industrial chemical processes rely on heterogeneous catalysis. These catalysts are typically transition metals or their compounds.

Key industrial processes:

1. Haber process (ammonia production):

$\text{N}_2\text{(g)} + 3\text{H}_2\text{(g)} \rightleftharpoons 2\text{NH}_3\text{(g)}$

Catalyst: Iron, Fe(s)

Ammonia is essential for manufacturing fertilizers. The solid iron catalyst allows the gaseous nitrogen and hydrogen to react more readily.

2. Reforming (petroleum refining):

$\text{C}_5\text{H}_{12}\text{(g)} \rightarrow \text{C}_5\text{H}_{12}\text{(g)} + \text{H}_2\text{(g)}$

Catalyst: Platinum, Pt(s), or Rhodium, Rh(s)

This process converts straight-chain alkanes into branched isomers and cyclic compounds to improve petrol octane ratings.

3. Hydrogenation of alkenes:

$\text{C}_2\text{H}_4\text{(g)} + \text{H}_2\text{(g)} \rightarrow \text{C}_2\text{H}_6\text{(g)}$

Catalyst: Nickel, Ni(s)

Used in the food industry to convert unsaturated fats (alkenes) into saturated fats (alkanes).

4. Contact process (sulfuric acid production):

$2\text{SO}_2\text{(g)} + \text{O}_2\text{(g)} \rightleftharpoons 2\text{SO}_3\text{(g)}$

Catalyst: Vanadium(V) oxide, $\text{V}_2\text{O}_5\text{(s)}$

This produces sulfur trioxide, which is then converted to sulfuric acid - one of the most important industrial chemicals.

Catalytic converters and atmospheric pollution

Since 1992, all petrol vehicles manufactured in the UK must be fitted with a catalytic converter to reduce harmful exhaust emissions and meet MOT test standards.

Structure and function:

Catalytic converters contain a ceramic honeycomb structure coated with precious metal catalysts (platinum, rhodium, and palladium). This honeycomb design provides a very large surface area for the heterogeneous catalysis to occur.

As hot exhaust gases pass over the catalyst surface, harmful pollutants are converted into less harmful products:

Catalysis - sustainability and economic importance

Catalysts play a crucial role in modern industry and sustainable chemistry. It is estimated that approximately 90% of all chemical materials are manufactured using catalysts.

Environmental and sustainability benefits:

Using catalysts supports sustainable chemistry by:

• Requiring less energy input, which reduces fossil fuel consumption
• Producing fewer carbon dioxide emissions from energy generation
• Reducing pollutant emissions through processes like catalytic converters
• Operating at lower temperatures, improving safety and reducing equipment wear

The focus on sustainability requires industries to use processes with high atom economy and minimal waste. Catalysts enable many reactions to proceed efficiently at lower temperatures, supporting these environmental goals and helping reduce contributions to global warming.

Autocatalysis

Autocatalysis occurs when a product of a reaction acts as a catalyst for that same reaction. The reaction typically starts slowly, then speeds up progressively as the catalytic product accumulates.