Catalysts and activation energy: how they work
The rate and extent of chemical change • Rate of reaction
Flashcards
Test your knowledge with interactive flashcards
Key concepts
What you'll likely be quizzed about
Definition of a catalyst
A catalyst is a substance that increases the rate of a chemical reaction without undergoing permanent chemical change. The catalyst participates in intermediate steps but regenerates by the end of the process. The catalyst therefore does not appear in the overall balanced chemical equation for the reaction. A catalyst changes the mechanism of the reaction rather than altering the position of chemical equilibrium for reversible reactions. The catalyst affects the rate at which equilibrium is reached but does not change the equilibrium concentrations.
Catalysts are not included in chemical equations
Catalysts do not appear as reactants or products in the overall chemical equation because they are not consumed. Laboratory notes or reaction descriptions often list catalysts above the reaction arrow or in the reaction conditions. Identification of a catalyst requires observation of its persistence after the reaction rather than presence in the equation. Catalyst omission from the equation requires careful attention when interpreting chemical equations and reaction conditions. The absence of a substance from the equation does not imply it played no role in determining the observed reaction rate.
Activation energy and energy profiles
Activation energy is the minimum energy that reacting particles require to form the activated complex and convert into products. Energy profile diagrams show reactants, products, and the energy barrier (activation energy) between them. The peak of the diagram represents the activated complex; the height of that peak above the reactants gives the activation energy. A greater activation energy corresponds to fewer particles having enough energy to react at a given temperature, and therefore to a slower reaction rate. The area under a Boltzmann distribution curve above the activation energy represents the fraction of particles able to react.
How catalysts lower activation energy
Catalysts provide an alternative reaction pathway that has a lower activation energy than the uncatalysed route. The catalyst stabilises transition states or forms temporary intermediates, reducing the energy required to reach the activated complex. Lowering the activation energy increases the fraction of collisions that result in reaction, thereby increasing the rate. Heterogeneous catalysts provide active surfaces where reactant molecules adsorb, react, and desorb as products. Homogeneous catalysts operate in the same phase as the reactants and often form short-lived intermediates that lead to products with lower energy barriers.
Effect of catalysts on reaction rate and identification
The presence of a catalyst increases the reaction rate by lowering activation energy and providing an alternative mechanism. Catalysts change the measured rate without altering the overall enthalpy change of the reaction. A catalyst speeds up both the forward and reverse reactions equally, so it does not shift equilibrium position but allows equilibrium to be reached faster. Identification of a catalyst can occur by observing changes in reaction rate when a substance is present or removed. Catalytic action often appears as faster product formation or reduced time to reach equilibrium while the catalyst itself remains chemically unchanged at the end of the reaction.
Key notes
Important points to keep in mind