Some chemical reactions start as soon as the reactants come into contact, whereas for many others, the chemicals fail to react without the supply of external energy. There are many reasons due to which reactants in close proximity may not immediately involve in a chemical reaction. Activation Energy is the energy which must be provided to potential reactants.
Moreover, in a chemical or a nuclear system so that a spark can be given to a chemical reaction or a nuclear reaction. We generally denote this energy by E. But it is very important to know which types of reactions require activation energy, and how much. Then only can we initiate and control chemical reactions in a safe manner.
Meaning of Activation Energy
To understand its concept, we can visualize it as the magnitude of the energy barrier i.e. the potential barrier. This barrier separates the minima of potential energy surfaces. It involves the initial and the final thermodynamic state of the system.
To define it, we have to analyze the initiation of chemical reactions. These reactions occur when molecules exchange electrons or when we bring together ions with opposite charges. To exchange electrons in the molecules, the bonds keeping the electrons tied with the molecule, have to be broken.
An external energy source can give the energy which is required to dislodge the electrons and hence can allow the chemical reaction to proceed. We express activation energy in units such as kilojoules, kilocalories or kilowatt-hours.
Once the reaction is on the way, it releases energy and then it is self-sustaining. So, we need this activation energy only at the beginning, to let the chemical reaction start.
Concept of Activation Energy
Based on this analysis, we say that activation energy is the minimum energy that we need to start a chemical reaction. Actually, when we supply this energy to reactants from a suitable external source, the molecules speed up and consequently collide more violently.
This violent collision knocks the electrons free. Therefore, the resulting atoms or ions react with each other to release energy and keep the reaction going.
Examples of Chemical Reactions Requiring Activation Energy
The common type of reaction using activation energy involves many kinds of fire or combustion. Such reactions are combining oxygen with a material that contains carbon. Carbon has molecular bonds with other elements in the fuel, whereas oxygen gas exists as two oxygen atoms together.
Carbon and oxygen don’t normally react with each other due to very strong existing molecular bonds. Ordinary molecule collisions cannot break this as it is very hard. But when external energy such as a flame breaks some of the bonds, the resulting oxygen and carbon atoms react to release the energy. And the fire generated will continue until it runs out of fuel.
Negative Activation Energy
A negative value of activation energy exists in the Arrhenius equations of reactions. Because the rate of reaction slows down while there is an increase in temperature.
Actually, reactions which show a negative value for the activation energy are the barrier for fewer reactions. An increase in the temperature for these reactions results in a reduction in the possibility of the colliding molecules to capture each other.
Effect of Catalysts on Activation Energy
Basically, catalysts are substances which increase the rate of the reaction without being consumed in the reaction while accomplishing this. Such substances work to modify the transition states of reactions in order to lower the activation energies of the aforementioned reactions. An enzyme is a special type of catalyst which comprises of only proteins and small molecule cofactors.
Although the action of catalysts reduces the overall activation energy of the reaction, it has no effect on the energies of the reactants or products of the reaction.
Solved Question for You
Q: Who discovered Activation Energy and when?
Ans: Swedish scientist Svante Arrhenius defined the term “activation energy” in 1880. This term defines the minimum energy needed for a set of chemical reactants to interact and form products.