Final answer:
The energy difference for a chemical reaction A to B is represented by the enthalpy change (ΔH). Exergonic reactions are more likely to occur due to their spontaneous nature. Activation energy is the minimum energy needed to initiate a reaction, not the energy change of the reaction itself.
Step-by-step explanation:
The difference in energy between two points that equals the energy change for reaction A to B can be represented by the enthalpy change (ΔH). In the context of the Gibbs free energy equation, this is denoted by the letter 'b' or ΔΔH in the given options. Furthermore, when discussing the type of reactions, exergonic reactions are more likely to occur because they release energy and typically occur spontaneously, contrasting with endergonic reactions that require an input of energy.
Regarding the activation energy, it is represented in a potential energy diagram by the height of the hill between the reactants and the products. This energy barrier must be overcome for a reaction to proceed. Thus, the activation energy is not equivalent to the energy change for the reaction but instead represents the minimum energy required to initiate the reaction.