Based on their activation energies and energy changes and assuming that all collision factors are the same, rank the following reactions from slowest to fastest. (a) Ea = 45 kJ>mol; E = -25 kJ>mol (b) Ea = 35 kJ>mol; E = -10 kJ>mol (c) Ea = 55 kJ>mol; E = 10 kJ>mol

Key Concepts

Activation Energy (Ea)

Activation energy is the minimum energy required for a chemical reaction to occur. It represents the energy barrier that reactants must overcome to form products. A lower activation energy typically correlates with a faster reaction rate, as fewer energy collisions are needed to initiate the reaction.

Energy Change (E)

The energy change of a reaction, often denoted as ΔE, indicates the difference in energy between the reactants and products. A negative ΔE signifies an exothermic reaction, where energy is released, while a positive ΔE indicates an endothermic reaction, where energy is absorbed. This energy change can influence the reaction rate, but activation energy is the more critical factor.

Collision Theory

Collision theory states that for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation. While the frequency and energy of collisions are essential, the activation energy determines how many of these collisions will lead to a successful reaction. Thus, understanding collision factors helps in ranking reaction rates based on their activation energies.