All Flashcards
Define 'elementary reaction'.
A single step in a reaction mechanism that represents the change of reactants to products.
Define 'overall reaction'.
The balanced chemical equation representing the entire reaction process, combining all elementary steps.
Define 'reaction energy profile'.
A graph showing the potential energy of reactants and products during a reaction, illustrating activation energy and overall energy change.
Define 'transition state'.
The highest energy point on a reaction energy profile, representing the unstable intermediate complex during a reaction.
Define 'activation energy (Ea)'.
The minimum energy required for a reaction to occur, represented by the energy difference between reactants and the transition state.
Define 'overall energy change (ΔE)'.
The difference in potential energy between the reactants and products, determining if a reaction is exothermic (ΔE < 0) or endothermic (ΔE > 0).
Define 'intermediate'.
A species formed during a reaction mechanism that is consumed in a subsequent step and does not appear in the overall balanced equation.
What is the effect of a higher activation energy (Ea) on the reaction rate?
A higher activation energy results in a slower reaction rate because more energy is required for the reaction to occur.
What is the effect of adding a catalyst to a reaction?
A catalyst increases the rate of reaction by lowering the activation energy.
What are the differences between 'intermediates' and 'transition states'?
Intermediates: Relatively stable, exist for a short time, appear between elementary steps. | Transition States: Unstable, exist momentarily, represent the highest energy point.
Differentiate between an 'exothermic' and 'endothermic' reaction in terms of energy change.
Exothermic: Releases heat, ΔE < 0, products have lower energy than reactants. | Endothermic: Absorbs heat, ΔE > 0, products have higher energy than reactants.