The activation energy for a simple chemical reaction \(\mathrm{A} \longrightarrow \mathrm{B}\) is \(\mathrm{E}\) in forward direction. The activation energy for reverse reaction: (a) Can be less than or more than \(\mathrm{E}_{\mathrm{a}}\) (b) Is always double of \(E_{a}\) (c) Is negative of \(\mathrm{E}_{\mathrm{a}}\) (d) Is always less than \(\mathrm{E}_{\mathrm{a}}\)

When we talk about a chemical reaction, we're referring to the process in which substances, called reactants, are transformed into different substances known as products. This change is not just a mixing of entities but is a result of breaking bonds between atoms in the reactants and forming new bonds to create the products. The beauty of chemical reactions lies in their diversity; they can be simple, like the rusting of iron, or complex, involving many steps, like in photosynthesis.

Every chemical reaction involves an initial push to get started, which is where the concept of activation energy, denoted by \(E_a\), comes into play. Think of it as the effort needed to push a boulder over a hill before it can roll down on the other side. Without this energy, the reactants may remain as they are indefinitely, even if the product formation is favorable.

In the realm of chemical processes, we often find that reactions are reversible, meaning that the transformation can proceed in both 'forward' and 'reverse' directions. In the forward reaction, reactants are converted into products, while in the reverse reaction, the products can revert back to the original reactants.

The important thing to note here is that the forward and reverse reactions involve different activation energies. The activation energy required for the forward reaction \(E_a\) is not necessarily the same as that for the reverse process. Because reactions can differ in how they release or absorb energy, this dictates whether the reverse reaction will require more, less, or equal energy to proceed when compared to the forward reaction.

Understanding the concept of energy in chemical reactions is pivotal. Each chemical species involved in a reaction stores a certain amount of energy. During a reaction, this energy will either be released to the surroundings, which characterizes an 'exothermic' reaction, or absorbed from the surroundings, known as an 'endothermic' reaction.

The difference in energy between products and reactants essentially determines the 'energy hill' that must be overcome by both the forward and reverse reactions. In an exothermic reaction, the products have less energy than the reactants, while the reverse is true for endothermic reactions. This contrast in energy levels significantly influences what we know as the activation energy for the reverse reaction. It is not fixed and can be higher or lower than the activation energy for the forward reaction, depending on these energy dynamics.