Activation energy: Difference between revisions
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The activation energy of a reaction is the minimum energy requirement for a successful collision to occur. All chemical reactions have an intermediate transition state that is at a higher energy than the reactants. This transition state must be reached and overcome before the product can be reached. This energy can be lowered through the use of a [[Catalysts|catalyst]] or an [ | The activation energy of a reaction is the minimum energy requirement for a successful collision to occur. All chemical reactions have an intermediate transition state that is at a higher energy than the reactants. This transition state must be reached and overcome before the product can be reached. This energy can be lowered through the use of a [[Catalysts|catalyst]] or an [[enzyme|enzyme]] via the creation of a lower second transition state. This means that a [[Catalysts|catalyst]] can make a reaction appear to occur faster, as more [[Molecule|molecules]] will have energy greater than activation energy (second transition state), therefore there are more successful collisions per second. | ||
Revision as of 15:28, 2 December 2011
The activation energy of a reaction is the minimum energy requirement for a successful collision to occur. All chemical reactions have an intermediate transition state that is at a higher energy than the reactants. This transition state must be reached and overcome before the product can be reached. This energy can be lowered through the use of a catalyst or an enzyme via the creation of a lower second transition state. This means that a catalyst can make a reaction appear to occur faster, as more molecules will have energy greater than activation energy (second transition state), therefore there are more successful collisions per second.