Catalysts play a central role in chemical kinetics by offering alternative reaction pathways that require less energy, significantly influencing reaction rates and industrial processes.

Understanding Catalysts

Catalysts are substances that alter the rate of a chemical reaction while remaining chemically unchanged at the end. When a catalyst is first introduced in this context, it is essential to define it precisely.

A catalyst does not affect the overall enthalpy change of a reaction, nor does it alter the position of equilibrium. Instead, it influences the kinetics of the process, making reactions proceed faster under the same conditions.

Alternative Reaction Pathways

A key feature of catalytic action is the provision of an alternative reaction pathway. This pathway has a lower activation energy, enabling a greater proportion of reacting particles to reach the energy threshold needed to form products.

• Catalysts reduce the height of the activation energy barrier, as shown clearly on enthalpy profile diagrams.

• Lower activation energy means a higher fraction of molecules possess sufficient energy to react at a given temperature.

• The overall enthalpy change (ΔH) remains unchanged because reactants and products are unaffected thermodynamically.

Catalysts therefore change the rate rather than the energetics of a reaction.

On an enthalpy profile diagram, a catalyst provides an alternative route with a lower activation energy while the overall enthalpy change (ΔH) between reactants and products stays the same.

Energy profile for a reaction comparing uncatalysed and catalysed pathways. The catalysed route has a noticeably lower activation energy peak, but the reactants and products lie at the same energy, so ΔH is unchanged. This visual emphasises that catalysts speed up reactions by lowering Ea, not by altering the overall enthalpy change. Source

Enthalpy Profile Diagrams and Catalysts

Enthalpy profile diagrams illustrate the change in enthalpy as a reaction progresses from reactants to products. When a catalyst is present:

• The diagram includes a lower peak, representing the reduced activation energy.

• The relative enthalpies of reactants and products remain the same.

• This visual comparison helps clarify why catalysed reactions proceed more rapidly.

Such diagrams are a crucial part of OCR A-Level Chemistry and are widely used to explain catalytic behaviour.

Types of Catalysts

Catalysts are commonly classified based on the physical state of the catalyst relative to the reactants. The specification emphasises the distinction between homogeneous and heterogeneous catalysts.

Homogeneous Catalysts

A homogeneous catalyst is in the same physical state as the reactants, usually aqueous or gaseous.

Homogeneous catalysis typically involves the catalyst forming an intermediate species with the reactants. This intermediate then decomposes to regenerate the catalyst and produce the reaction products.

Important features include:

• Close molecular interaction with reactants

• Often used in biochemical and solution-based reactions

• Catalyst and reactants mix uniformly, promoting effective collisions

Homogeneous catalysts are advantageous where precise control over reaction mechanisms is required.

Heterogeneous Catalysts

In heterogeneous catalysis, the catalyst is in a different physical state from the reactants, usually a solid with gases or liquids reacting at its surface.

Schematic of a heterogeneous catalytic system showing reactant molecules interacting with the surface of a solid catalyst in a separate phase. The diagram highlights how reactions occur at the interface, with adsorption, reaction, and product release. Any additional industrial context shown goes beyond the OCR syllabus but still illustrates the same key idea: heterogeneous catalysts operate at a surface between phases. Source

Heterogeneous catalysis involves several essential steps:

• Adsorption of reactant molecules onto the catalyst surface

• Weakening of bonds within adsorbed molecules

• Reaction on the catalyst surface

• Desorption of product molecules, leaving the catalyst available for reuse

Key features include:

• Catalyst surfaces contain active sites crucial for reaction

• Common in industrial processes such as hydrogenation and catalytic converters

• Easy separation of catalyst and products, improving practicality

Heterogeneous catalysts are widely used because they offer durability, reusability, and straightforward recovery.

How Catalysts Affect Reaction Rate

Catalysts influence reaction rate by increasing the number of successful collisions per unit time. They do so by:

• Lowering activation energy, increasing the fraction of molecules that exceed this threshold

• Providing specific binding sites that stabilise transition states

• Enhancing the orientation or local concentration of reactants

These effects are directly linked to the ideas of collision theory and energy distribution, although those topics belong to separate subsubtopics.

Although this relationship does not change with the presence of a catalyst, the observed rate increases because the catalyst allows more successful collisions.

A catalyst’s effectiveness depends on factors such as surface area (for heterogeneous systems), catalyst purity, and temperature.

Catalyst Behaviour in Reactions

Catalysts must be regenerated at the end of a reaction. Their role is cyclical:

• Participate in the reaction mechanism

• Form intermediates

• Re-emerge unchanged after product formation

Important considerations:

• The catalyst must be chemically stable

• Poisons can block active sites in heterogeneous catalysts, reducing efficiency

• Reaction conditions must be compatible with catalyst integrity

Catalysts are essential for modern chemical synthesis and environmental technologies, aligning with the specification’s emphasis on alternative pathways and activation energy reduction.