What is Activation energy?
[[File:Activation energy.svg|thumb|right|250px|Activation energy with and without a catalyst]] The '''activation energy''' of a [[chemical reaction]] is the minimum energy that is needed to make the reaction happen. It usually has the symbol E<sub>a</sub> and it is measured in [[kilojoule per mole]]. It can be thought of as a barrier between the reagents and the products of a reaction. The activation energy is the difference in energy between the [[transition state]] and the starting reagents. == Arrhenius equation == This equation is a way of calculating the activation energy. Its form is: :<math>k = A e^{{-E_a}/{RT}}</math> Here, ''k'' is the rate constant of the reaction (which can be measured using [[chemical kinetics]]), ''A'' is a frequency factor, ''R'' is the universal [[gas constant]], ''T'' is the [[temperature]]. If we measure how fast the reaction goes at different temperatures, then we can make a graph and calculate the activation energy using this equation. == Catalysis == When using a [[catalyst]] the activation energy is made smaller. This means that the reaction can go faster at the same temperature. The catalyst does not change the energy of the reagents or the products. Also, the catalyst does not change in the reaction. Therefore, it can be used again and again and only a small amount of a catalyst is needed. [[Category:Chemical kinetics]] {{Commonscat}} == References ==