Activation energy

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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.
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The [[Activation_energy|activation energy]] of a [[Reaction|reaction]] is the minimum amount of energy required to break [[Chemical bond|chemical bonds]].  All chemical reactions have an [[Intermediate transition state|intermediate transition state]] that is at a higher energy than the reactants. This [[Transition State|transition state]] must be overcomed before the product can form.  This energy can be lowered by the use of a [[Catalysts|catalyst]] known as [[Enzyme|enzymes]] in biochemistry. A catalyst lowers the activation energy of a reaction without being used up in the process. This means that a [[Catalysts|catalyst]] can increase the speed of a reaction, as more [[Molecule|molecules]] will have an energy greater than activation energy ([[Second transition state|second transition state]]). Therefore, there are more successful collisions per second. In enzyme catalysed reactions, [[Enzymes|enzymes react]] with [[Substrate|substrate]] to form an [[Enzyme-substrate complex|enzyme-substrate complex]]. This enzyme substrate complex lowers the activation energy of the reaction and thus, leads to the formation of products. In some cases, the activation energy is too high so it is not possible for the reaction to even happen without the presence of [[Enzyme|enzyme]]. An enzyme makes the reaction more likely to happen by lowering the activation energy. This helps the [[molecule|molecule]] to get to the transition state which is the state where old bonds are broken and form new bonds which leads to the formation of products.

Latest revision as of 03:13, 30 November 2013

The activation energy of a reaction is the minimum amount of energy required to break chemical bonds.  All chemical reactions have an intermediate transition state that is at a higher energy than the reactants. This transition state must be overcomed before the product can form.  This energy can be lowered by the use of a catalyst known as enzymes in biochemistry. A catalyst lowers the activation energy of a reaction without being used up in the process. This means that a catalyst can increase the speed of a reaction, as more molecules will have an energy greater than activation energy (second transition state). Therefore, there are more successful collisions per second. In enzyme catalysed reactions, enzymes react with substrate to form an enzyme-substrate complex. This enzyme substrate complex lowers the activation energy of the reaction and thus, leads to the formation of products. In some cases, the activation energy is too high so it is not possible for the reaction to even happen without the presence of enzyme. An enzyme makes the reaction more likely to happen by lowering the activation energy. This helps the molecule to get to the transition state which is the state where old bonds are broken and form new bonds which leads to the formation of products.

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