Activation energy

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The activation energy of a reaction is the minimum energy requirement for a successful collision to occur. This energy can be lowered through the use of a [[Catalysts|catalyst]] or an [http://bms.ncl.ac.uk/wiki/index.php/Enzyme enzyme]. 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, therefore there are more successful collisions per second.
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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 [http://bms.ncl.ac.uk/wiki/index.php/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 14:36, 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.

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