Search results
Results from the WOW.Com Content Network
In chemistry and physics, activation energy is the minimum amount of energy needed to start a chemical reaction. Reactants often get activation energy from heat, but sometimes energy comes from light or energy released by other chemical reactions.
The activation energy (Ea), labeled ΔG ‡ in Figure 2, is the energy difference between the reactants and the activated complex, also known as transition state. In a chemical reaction, the transition state is defined as the highest-energy state of the system.
In chemistry, activation energy is the minimum amount of energy required for a chemical reaction. The activation energy can be thought of as the magnitude of a potential barrier that the reacting molecules need to overcome to initiate a reaction and convert into products.
In the Arrhenius model of reaction rates, activation energy is the minimum amount of energy that must be available to reactants for a chemical reaction to occur. [1] The activation energy (E a) of a reaction is measured in kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol). [2]
A minimum energy (activation energy,v\(E_a\)) is required for a collision between molecules to result in a chemical reaction. Plots of potential energy for a system versus the reaction coordinate show an energy barrier that must be overcome for the reaction to occur.
Activation energy, in chemistry, the minimum amount of energy that is required to activate atoms or molecules to a condition in which they can undergo chemical transformation or physical transport. Activation energies are determined from experimental rate constants or diffusion coefficients.
Activation energy is the minimum amount of energy required to initiate a reaction. It is the height of the potential energy barrier between the potential energy minima of the reactants and products. Activation energy is denoted by E a and typically has units of kilojoules per mole (kJ/mol) or kilocalories per mole (kcal/mol).