**Activation Energy**

Activation energy is a concept used in chemistry to describe the energy barrier that must be overcome for a chemical reaction to occur. It is the minimum amount of energy required for reactant molecules to undergo a chemical reaction and form products.

**Definition:**
Activation energy can be defined as the minimum amount of energy required for a chemical reaction to occur by breaking the bonds of reactant molecules and initiating the formation of new bonds to produce products.

**Explanation:**
1. Energy Barrier: Activation energy represents the energy barrier that reactant molecules must overcome in order to reach the activated state, where the reaction can proceed. This energy barrier exists because reactant molecules must reach a specific energy level, known as the transition state, to break the existing bonds and form new ones.

2. Collision Theory: According to the collision theory, for a chemical reaction to occur, reactant molecules must collide with sufficient energy and with the proper orientation. The activation energy corresponds to the energy required for reactant molecules to collide with enough energy to surpass the energy barrier and initiate the reaction.

3. Reaction Progress: As the reaction progresses, the reactant molecules gradually transition from their initial state to the activated state. Once the activation energy is surpassed, the reaction proceeds spontaneously as the products are formed.

**Threshold Energy**

Threshold energy is another term used to describe the energy required for a reaction to occur. It is often used interchangeably with activation energy, but there is a subtle difference in their definitions.

**Definition:**
Threshold energy is the minimum energy required for a reactant molecule to undergo a chemical reaction. It represents the energy level at which the reactant molecule possesses enough energy to surpass the energy barrier and initiate the reaction.

**Explanation:**
1. Individual Reactant Molecules: Threshold energy is defined at the level of individual reactant molecules. It represents the energy level at which a specific reactant molecule has enough energy to proceed with the reaction.

2. Activation Energy for Ensemble: Activation energy, on the other hand, is defined as the average energy required for a collection or ensemble of reactant molecules to undergo the reaction. It takes into account the energy distribution of the reactant molecules and represents the average energy barrier that needs to be overcome.

3. Energy Distribution: In a system of reactant molecules, each molecule has a different energy level due to their thermal motion. Some molecules may possess energy higher than the threshold energy, while others may have energy lower than the threshold. Activation energy considers the energy distribution and provides an average value that characterizes the energy barrier for the reaction.

In summary, activation energy and threshold energy both describe the minimum energy required for a chemical reaction to occur. However, activation energy is an average value for a collection of reactant molecules, while threshold energy is defined at the level of individual molecules.