Activation Energy

If I ask you how you start your car, your answer to this will be very simple. Firstly, turn on the key. This turning on gives a spark which activates the burning of gasoline in the engine and the car starts.

Here, the combustion of the gas acts as a source of energy to activate the engine.

So, this source of energy is the activation energy in chemistry.


What is the Activation Energy?

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

In chemistry, we call it the minimum amount of energy (or threshold energy) needed to activate or energize molecules or atoms to undergo a chemical reaction or transformation. 

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


Concept of Activation Energy

We are familiar with chemical reactions such as the burning of gas in air or the combustion of hydrogen or oxygen gases. This does not occur unless we give some energy in some form to the reacting system. Thus, extra energy should be given to the reactants to bring their energy equal to the threshold energy.

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Let’s consider the two products A and B undergo a reaction to form the product C.

A + B → C

We know that the reactant already has some energy, i.e., Er. Now, some amount of energy is given to the reactant. As they gain energy, the molecules of A and B collide and after that, they stick together to form AB at the transition state. This transition state is the energy barrier.

Now, to cross this barrier, extra energy is given and this energy is the activation energy.

So, the excess energy (the energy above the average energy of the reactants) is supplied to reactants to undergo chemical reactions is called activation energy.

Let’s say the reactants (A + B) have 20 KJ of energy, and for crossing the transition state, it needs 60 KJ of energy,  and this energy is the threshold energy (ET). This means 40 KJ of extra energy is added to cross the barrier. So, this extra energy is the activation energy or Ea. After that, they transform into product C.

Here, the energy of product (Ep) > Energy of reactant (Er). However, the reverse is also true.

This means the reactants have to absorb the energy to transform into a product. Therefore, this process is endothermic.

If the  Ep < Er,  the process is exothermic because more energy is released.

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Hence, for any transformation to occur, two things are necessary:

  1. Effective collision, and

  2. Enough energy to cross the energy barrier.


Activation Energy Formula

The activation energy is equal to the difference between the threshold energy needed for the reaction and the average kinetic energy of all the reacting molecules. i.e.,

                        

Ea = Threshold energy (EThreshold) - Average kinetic energy of the reacting molecules (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,


Low activation energies

Fast reactions

High activation energies

Slow reactions


1. Endothermic Chemical Reaction

We know that in an endothermic reaction, the energy released is positive.

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

We saw that the reactant already had 20 KJ of energy, and it needed 40 KJ of extra energy for breaking the bonds. However, the energy released during the bond formation is less than Er.

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


2. Exothermic Chemical Reaction

Let’s say, the reactant has 10 KJ of energy, on getting 50 KJ of extra energy, it transforms into the product. However, during product formation, 60 KJ of energy is released. This means an extra 10 KJ of energy is released, which is more than Er.

Hence, this process is exothermic.

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

The relation between activation energy and rate constant

Arrhenius equation explains the accurate dependency of the rate constant ‘K’ and the temperature T. For this; he gave an equation, i.e.,

K = Ae\[^{-\frac{Ea}{RT}}\]

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Where,

K = Rate constant

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

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

R = Universal gas constant

T = Temperature in Kelvin

∴ With the increase in the activation energy Ea, the rate constant K decreases, and therefore the rate of reaction decreases.

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This is a graph of ln (K) versus 1/T. A straight line drawn has a slope, i.e., m.

The value of m = - Ea/R, where R = 8.314 J/K\[^{-mol}\].

Through this formula, we can calculate the activation energy. 

FAQ (Frequently Asked Questions)

Q1: What is the Role of Activation Energy?

Ans: In a chemical reaction, some bonds break while some bonds form. So, the excess energy, i.e., the energy above the average kinetic energy of the reacting molecules are supplied to them so they can move together and overcome forces of repulsion and break bonds to transform. The energy that aids in this process is the activation energy.

Q2: What are the Five Factors that Affect the Rate of Reaction?

Ans: Following are the five factors:

  1. The surface area of a solid reactant

  2. Concentration/pressure of a reactant

  3. Temperature

  4. Nature of reactants

  5. Presence or absence of a catalyst

Q3:  How is Activation Energy Related to Temperature?

Ans: As the temperature increases, the kinetic energy of molecules increases. We know that 10²⁷ collisions occur in 1cm³ of a reacting molecule just in a second. So, on increasing the temperature, the collision between reacting molecules increases, and therefore, the activation energy increases.

Q4: Can Activation Energy be Negative?

Ans: No. The activation energy is extra energy supplied for any reaction to occur. Therefore, it is always positive. However, at lower temperatures, when the activation energy is lower, the rate constant K  approaches the pre-exponential factor.