Activation Energy

The minimum amount of energy required to bring atoms or molecules to a state where they can execute a chemical transformation or physical transport is known as activation energy in chemistry. In transition-state theory, the activation energy is the difference in energy content between molecules or atoms in an activated or transition-state configuration and their starting configuration. In mathematical formulations for quantities like the reaction rate constant, k = Ae^(-Ea/RT), and the diffusion coefficient, D = D_oe^(−Ea/RT) , the activation energy is often indicated by the sign Ea.
k = Ae ^(−Ea/RT)

A= Arrhenius constant

Ea = Activation energy
K=Rate constant
R=Gas Constant
T=Temperature

Formula

The difference seen between the threshold energy required for the reaction to proceed and the average kinetic energy of all reacting molecules in the reactant species is the activation energy that is
E_a= E_threshold – E_average
This means that if a reaction’s activation energy is low, the proportions of effective collisions are big, the threshold energy is high, and the reaction rate is high. The number of effective collisions is minimal and the reaction rate is slow when the activation energy is high.

Chemical Reaction activation energy

Why would a negative G reaction, which releases energy, necessitate the use of energy to proceed? The reason for this is that a chemical reaction involves multiple steps. During chemical reactions, chemical bonds are broken and new ones are formed. When a glucose molecule, for example, is broken down, the bonds between the carbon atoms are broken. Because they are energy-storing connections, when they are broken, they release energy. However, the molecule must be twisted in order to allow the bonds to be broken. The transition state is a high-energy, unstable state that can be reached with only a little amount of energy.As a result, reactant molecules do not spend much time in their transition state and instead travel through it fast.
Exergonic reactions (G0) are sometimes coupled with endergonic reactions (G>0), allowing them to progress. Energy coupling is the term for this spontaneous switch from one reaction to another. The endergonic reaction absorbs the free energy generated by the exergonic reaction. A transmembrane ion pump, which is critical for cellular function, is one example of energy coupling using ATP.

Free energy diagram

The energy profiles for a specific reaction are depicted in free energy diagrams. The products in the diagram will exist in a lower or higher energy state than the reactants, depending on whether the reaction is exergonic (G0) or endergonic (G>0). The activation energy, on the other hand, is independent of the reaction’s G. To put it another way, the activation energy at a given temperature is determined by the nature of the chemical transformation taking place, not by the relative energy states of the reactants and products.
Although the activation energy notion is discussed in the context of an exergonic forward reaction in the graphic above, the same concepts apply to the reverse reaction, which must be endergonic. It’s worth noting that the reverse reaction’s activation energy is Greater than the forward reaction.

Heat energy from the surroundings

The activation energy required to move processes forward is frequently provided by heat energy from the environment. Heat energy (the total bond energy of reactants or products in a chemical reaction) accelerates molecular motion, resulting in more collisions with greater frequency and force. By gradually shifting atoms and bonds inside the molecule, it also aids them in reaching their transition state. As a result, heating a system causes chemical reactants inside that system to react more frequently. Increasing the pressure in a system can provide a similar effect. Once the reactants have absorbed enough thermal energy from their surroundings to reach the transition state, the reaction will proceed.
The rate at which a reaction will progress is determined by its activation energy. The slower the chemical reaction, the higher the activation energy. The example of iron rusting exemplifies a response that is intrinsically sluggish. Because of the high EA, this reaction takes a long time to complete. Furthermore, until the activation energy of several fuels is overcome by sufficient heat from a spark, the burning of numerous fuels, which is strongly exergonic, will take place at a low pace. However, once they start to burn, the chemical processes generate enough heat to keep the fire going, providing the activation energy for the surrounding fuel molecules.

Temperature and activation energy

When two billiard balls collide, they bounce off each other. When two molecules, X and Y, collide, this is also the most common outcome: they bounce off each other, fully unmodified and undamaged. X and Y must collide with enough energy to break chemical bonds in order for a chemical reaction to occur. Because chemical bonds in the reactants are broken and new bonds in the products are established in any chemical reaction, this is the case. To properly launch a reaction, the reactants must be traveling rapidly enough (with enough kinetic energy) to collide with enough force to break bonds. This is the bare minimum of energy with which molecules must move. The activation energy is the minimal energy with which molecules must move in order for a collision to result in a chemical reaction.

Conclusion

The activation energy is defined as the difference in energy content between molecules or atoms in an activated or transition-state configuration and their initial configuration in transition-state theory.
Chemical bonds are destroyed and new ones are generated during chemical reactions. The bonds between the carbon atoms in a glucose molecule, for example, are destroyed as it is broken down. When these bonds are broken, they release energy since they are energy-storing bonds. The molecule must be distorted, though, to get them into a state where the bonds can be broken