Editing Activation energy (section)

== Catalysts ==
{{Main|Catalysis}}
[[File:Activation2 updated.svg|left|thumb|Example of an enzyme-catalysed [[Exothermic reaction|exothermic]] reaction]]
[[Image:Activation energy.svg|thumb|240x240px|The relationship between activation energy (<math>E_\textrm{a}</math>) and [[Standard enthalpy of reaction|enthalpy of reaction]] (Δ''H'') with and without a catalyst, plotted against the [[reaction coordinate]]. The highest energy position (peak position) represents the transition state. With the catalyst, the energy required to enter transition state decreases, thereby decreasing the energy required to initiate the reaction.|alt=]]A substance that modifies the transition state to lower the activation energy is termed a [[catalyst]]; a catalyst composed only of protein and (if applicable) small molecule cofactors is termed an [[enzyme]]. A catalyst increases the rate of reaction without being consumed in the reaction.<ref>{{Cite web|url=http://antoine.frostburg.edu/chem/senese/101/reactions/faq/examples-of-catalysts.shtml|title=General Chemistry Online: FAQ: Chemical change: What are some examples of reactions that involve catalysts?|website=antoine.frostburg.edu|access-date=2017-01-13|archive-date=2016-11-29|archive-url=https://web.archive.org/web/20161129120052/http://antoine.frostburg.edu/chem/senese/101/reactions/faq/examples-of-catalysts.shtml|url-status=live}}</ref> In addition, the catalyst lowers the activation energy, but it does not change the energies of the original reactants or products, and so does not change equilibrium.<ref>{{cite web|last1=Bui|first1=Matthew|title=The Arrhenius Law: Activation Energies|url=https://chem.libretexts.org/Core/Physical_and_Theoretical_Chemistry/Kinetics/Modeling_Reaction_Kinetics/Temperature_Dependence_of_Reaction_Rates/The_Arrhenius_Law/The_Arrhenius_Law%3A_Activation_Energies|website=Chemistry LibreTexts|date=2 October 2013 |publisher=UC Davis|access-date=February 17, 2017|archive-date=February 18, 2017|archive-url=https://web.archive.org/web/20170218143413/https://chem.libretexts.org/Core/Physical_and_Theoretical_Chemistry/Kinetics/Modeling_Reaction_Kinetics/Temperature_Dependence_of_Reaction_Rates/The_Arrhenius_Law/The_Arrhenius_Law%3A_Activation_Energies|url-status=live}}</ref> Rather, the reactant energy and the product energy remain the same and only the ''activation energy'' is altered (lowered).

A catalyst is able to reduce the activation energy by forming a transition state in a more favorable manner. Catalysts, by nature, create a more "comfortable" fit for the [[Substrate (chemistry)#Biochemistry|substrate]] of a reaction to progress to a transition state. This is possible due to a release of energy that occurs when the substrate binds to the [[active site]] of a catalyst. This energy is known as Binding Energy. Upon binding to a catalyst, substrates partake in numerous stabilizing forces while within the active site (e.g. [[hydrogen bond]]ing or [[van der Waals force]]s). Specific and favorable bonding occurs within the active site until the substrate forms to become the high-energy transition state. Forming the transition state is more favorable with the catalyst because the favorable stabilizing interactions within the active site ''release'' energy. A chemical reaction is able to manufacture a high-energy transition state molecule more readily when there is a stabilizing fit within the active site of a catalyst. The binding energy of a reaction is this energy released when favorable interactions between substrate and catalyst occur. The binding energy released assists in achieving the unstable transition state. Reactions without catalysts need a higher input of energy to achieve the transition state. Non-catalyzed reactions do not have free energy available from active site stabilizing interactions, such as catalytic enzyme reactions.<ref>{{Cite book|title=Biochemistry - Ninth Edition|last=Berg|first=Jeremy |publisher=WH Freeman and Company|year=2019|isbn=978-1-319-11467-1|location=New York, NY|pages=240–244}}</ref>