A Short Note on Activation Energy

Consider waking up on a day when you have a lot of exciting things scheduled for the day. It happens to everyone from time to time that, despite the great day that lies ahead, it takes a little additional effort to get out of bed. It is true that once you are awake, you can glide through the remainder of the day, but there is a small hurdle you must overcome in order to reach that point.

The activation energy of a chemical reaction is analogous to the “hump” that you have to go over in order to get out of bed in the morning. Even energy-releasing (exergonic) processes require a small amount of energy to get started before they can progress to the energy-releasing steps of the reaction. This initial energy input, which is later repaid as the reaction progresses, is referred to as the activation energy and is abbreviated EA in the scientific literature.

Activation Energy 

The activation energy is the extra energy given to get useful work done.

k=Ae−Ea/RT

We refer to it as the lowest quantity of energy (also known as the threshold energy) required to activate or energise molecules or atoms in order for them to conduct a chemical reaction or transformation in chemistry.

The activation energy units are measured in LCal/mo, KJ/mol, and J/mol, respectively.

Factors Affecting Activation Energy 

Activation energy depends on two factors.

1. Nature of Reactants

In the case of an ionic reactant, the value of (Ea) will be low due to the attraction that exists between the reactant and the reacting species. As opposed to this, in the case of a covalent reactant, the value of Ea will be high because it takes a lot of energy to break apart the older bonds.

2. Effect of Catalyst

The positive catalyst creates an alternate path in which the value of Ea will be low, whereas the negative catalyst creates an alternate way in which the value of Ea will be high, and vice versa.

Activation Energy Formula

A reaction’s activation energy is the difference between the threshold energy required for the reaction and the combined average kinetic energy of all of the reacting molecules. i.e.,                       

Each reaction has a certain value of Ea and this determines the fraction of total collisions that are effective.

This means if the activation energy for a reaction is low, numerous molecules have this energy, and the fractions of effective collisions are large. The reaction proceeds at a pace. If the activation energy is high, a fraction of effective collisions is small, and the reaction takes place slowly. 

Thus,

1. Endothermic Chemical Reaction

We already know that the energy released during an endothermic process is positive.

∴ Δ H = + ve (enthalpy of a chemical reaction is positive).

We discovered that the reactant already had 20 KJ of energy, and that it required an additional 40 KJ of energy to break the bonds. The energy produced during the bond formation, on the other hand, is smaller than Er.

So, Δ H = Enthalpy of the product (HP) – Enthalpy of reactant (HR)

2. Exothermic Chemical Reaction

Let us suppose that the reactant contains 10 KJ of energy and that it turns into the product when it receives an additional 50 KJ of energy. However, 60 KJ of energy is released throughout the product creation process. This translates to an additional 10 KJ of energy being released, which is greater than Er’s total.

As a result, this is an exothermic reaction.

Here, Δ H = – ve, i.e., HP < HR, i.e., – (HP – HR)

The relationship between activation energy and rate constant is seen in the graph below.

The accurate relationship between the rate constant ‘K’ and the temperature T is explained by the Arrhenius equation. This was accomplished through the use of an equation, which is to say,

K=Ae–EaRT

Where K denotes the rate constant.

A = Arrhenius factor/pre-exponential factor/frequency factor

E = Activation energy in J/mol or KJ/mol

R denotes the universal gas constant.

T is the temperature in Kelvin degrees Celsius. In response to a rise in the activation energy Ea, the rate constant K lowers, and as a result, the rate of reaction decreases as well.

The value of m = – Ea/R, where R = 8.314 J/K-mol

Catalysts

A catalyst is a chemical substance that has the ability to either raise or reduce the rate of a chemical reaction in either direction. In the case of activation energy, a catalyst lowers the level of activation energy. The energy of the initial reactants, on the other hand, remain unchanged. The activation energy is the only thing that a catalyst changes.

Conclusion

The rate at which a certain reaction will progress is determined by the activation energy of that process. When the activation energy is high, the chemical reaction will proceed at a slower rate.

Activation energy is required for all chemical reactions, even exothermic reactions, in order to get them started. It is necessary to have activation energy in order for reactants to move together, overcome forces of repulsion, and begin to dissolve bonds.