Editing Catalysis (section)

==General principles==

===Example===
An illustrative example is the effect of catalysts to speed the decomposition of [[hydrogen peroxide]] into water and [[oxygen]]:
:2 H{{sub|2}}O{{sub|2}} → 2 H{{sub|2}}O + O{{sub|2}}
This reaction proceeds because the reaction products are more stable than the starting compound, but this decomposition is so slow that hydrogen peroxide solutions are commercially available. In the presence of a catalyst such as [[manganese dioxide]] this reaction proceeds much more rapidly. This effect is readily seen by the [[effervescence (chemistry)|effervescence]] of oxygen.<ref>{{cite web |publisher=[[University of Minnesota]] |title=Genie in a Bottle |date=2005-03-02 |url=http://www.chem.umn.edu/services/lecturedemo/info/genie.htm |url-status=dead |archive-url=https://web.archive.org/web/20080405195443/http://www.chem.umn.edu/services/lecturedemo/info/genie.htm |archive-date=2008-04-05}}</ref> The catalyst is not consumed in the reaction, and may be recovered unchanged and re-used indefinitely. Accordingly, manganese dioxide is said to ''catalyze'' this reaction. In living organisms, this reaction is catalyzed by [[enzyme]]s (proteins that serve as catalysts) such as [[catalase]].

Another example is the effect of catalysts on air pollution and reducing the amount of carbon monoxide. Development of active and selective catalysts for the conversion of carbon monoxide into desirable products is one of the most important roles of catalysts. Using catalysts for hydrogenation of carbon monoxide helps to remove this toxic gas and also attain useful materials.<ref>{{cite journal |last1=Torkashvand |first1=Mostafa |last2=Sarabadani Tafreshi |first2=Saeedeh |last3=de Leeuw |first3=Nora H. |title=Density Functional Theory Study of the Hydrogenation of Carbon Monoxide over the Co (001) Surface: Implications for the Fischer–Tropsch Process |date=May 2023 |journal=Catalysts |language=en |volume=13 |issue=5 |pages=837 |issn=2073-4344 |doi=10.3390/catal13050837 |doi-access=free}}</ref>

===Units===
The [[SI derived unit]] for measuring the '''catalytic activity''' of a catalyst is the [[katal]], which is quantified in moles per second. The productivity of a catalyst can be described by the [[turnover number]] (TON) and the catalytic activity by the ''turn over frequency'' (TOF), which is the TON per time unit. The biochemical equivalent is the [[enzyme unit]]. For more information on the efficiency of enzymatic catalysis, see the article on ''[[enzyme#Kinetics|enzymes]]''.

===Catalytic reaction mechanisms===
{{Main|catalytic cycle}}
In general, chemical reactions occur faster in the presence of a catalyst because the catalyst provides an alternative [[reaction mechanism]] (reaction pathway) having a lower [[activation energy]] than the noncatalyzed mechanism. In catalyzed mechanisms, the catalyst is regenerated.<ref name=LM82/><ref>{{cite book |last1=Laidler |first1=Keith J. |last2=Meiser |first2=John H. |title=Physical Chemistry |date=1982 |publisher=Benjamin/Cummings |isbn=0-8053-5682-7 |pages=424–425}}</ref><ref>{{cite book |last1=Atkins |first1=Peter |last2=de Paula |first2=Julio |title=Atkins' Physical Chemistry |date=2006 |publisher=W.H.Freeman |isbn=0-7167-8759-8 |page=839 |edition=8th |quote=The catalyst lowers the activation energy of the reaction by providing an alternative path that avoids the slow, rate-determining step of the uncatalyzed reaction}}</ref><ref name=Steinfeld>{{cite book |last1=Steinfeld |first1=Jeffrey I. |last2=Francisco |first2=Joseph S. |last3=Hase |first3=William L. |title=Chemical Kinetics and Dynamics |date=1999 |publisher=Prentice Hall |isbn=0-13-737123-3 |pages=147–150 |edition=2nd |quote=The catalyst concentration [C] appears in the rate expression, but not in the equilibrium ratio.}}</ref>

As a simple example occurring in the gas phase, the reaction 2 SO<sub>2</sub> + O<sub>2</sub> → 2 SO<sub>3</sub> can be catalyzed by adding [[nitric oxide]]. The reaction occurs in two steps: 
: 2{{nbsp}}NO + O<sub>2</sub> → 2{{nbsp}}NO<sub>2</sub> (rate-determining)
: NO<sub>2</sub> + SO<sub>2</sub> → NO + SO<sub>3</sub> (fast)
The NO catalyst is regenerated. The overall rate is the rate of the slow step<ref name=Steinfeld/> 
:v=2k<sub>1</sub>[NO]<sup>2</sup>[O<sub>2</sub>].

An example of [[heterogeneous catalysis]] is the reaction of [[oxygen]] and [[hydrogen]] on the surface of [[titanium dioxide]] (TiO{{sub|2}}, or ''titania'') to produce water. [[Scanning tunneling microscopy]] showed that the molecules undergo [[adsorption]] and [[dissociation (chemistry)|dissociation]]. The dissociated, surface-bound O and H atoms [[diffusion|diffuse]] together. The intermediate reaction states are: HO{{sub|2}}, H{{sub|2}}O{{sub|2}}, then H{{sub|3}}O{{sub|2}} and the reaction product ([[water dimer|water molecule dimers]]), after which the water molecule [[desorption|desorbs]] from the catalyst surface.<ref>{{cite news
 |last=Jacoby
 |first=Mitch
 |title=Making Water Step by Step
 |date=16 February 2009
 |newspaper=[[Chemical & Engineering News]]
 |page=10
 |url=https://pubsapp.acs.org/cen/news/87/i07/8707notw6.html?
}}</ref><ref>{{cite journal
 |vauthors=Matthiesen J, Wendt S, Hansen JØ, Madsen GK, Lira E, Galliker P, Vestergaard EK, Schaub R, Laegsgaard E, Hammer B, Besenbacher F
 |year=2009
 |title=Observation of All the Intermediate Steps of a Chemical Reaction on an Oxide Surface by Scanning Tunneling Microscopy
 |journal=[[ACS Nano]]
 |volume=3
 |issue=3
 |pages=517–26
 |issn=1520-605X
 |citeseerx=10.1.1.711.974
 |pmid=19309169
 |doi=10.1021/nn8008245
}}</ref>

===Reaction energetics===
[[File:CatalysisScheme-en.svg|thumb|upright=1.25|Generic potential energy diagram showing the effect of a catalyst in a hypothetical exothermic chemical reaction X + Y to give Z. The presence of the catalyst opens a different reaction pathway (shown in red) with lower activation energy. The final result and the overall thermodynamics are the same.]]
Catalysts enable pathways that differ from the uncatalyzed reactions. These pathways have lower [[activation energy]]. Consequently, more molecular collisions have the energy needed to reach the [[transition state]]. Hence, catalysts can enable reactions that would otherwise be blocked or slowed by a kinetic barrier. The catalyst may increase the reaction rate or selectivity, or enable the reaction at lower temperatures. This effect can be illustrated with an [[energy profile (chemistry)|energy profile]] diagram.

In the catalyzed [[elementary reaction]], catalysts do '''not''' change the extent of a reaction: they have '''no''' effect on the [[chemical equilibrium]] of a reaction. The ratio of the forward and the reverse reaction rates is unaffected (see also [[thermodynamics]]). The [[second law of thermodynamics]] describes why a catalyst does not change the chemical equilibrium of a reaction. Suppose there was such a catalyst that shifted an equilibrium. Introducing the catalyst to the system would result in a reaction to move to the new equilibrium, producing energy. Production of energy is a necessary result since reactions are spontaneous only if [[Gibbs free energy]] is produced, and if there is no energy barrier, there is no need for a catalyst. Then, removing the catalyst would also result in a reaction, producing energy; i.e. the addition and its reverse process, removal, would both produce energy. Thus, a catalyst that could change the equilibrium would be a [[perpetual motion machine]], a contradiction to the laws of thermodynamics.<ref>Robertson, A.J.B. (1970) ''Catalysis of Gas Reactions by Metals''. Logos Press, London.</ref> Thus, catalysts '''do not''' alter the equilibrium constant. (A catalyst can however change the equilibrium concentrations by reacting in a subsequent step. It is then consumed as the reaction proceeds, and thus it is also a reactant. Illustrative is the base-catalyzed [[hydrolysis]] of [[ester]]s, where the produced [[carboxylic acid]] immediately reacts with the base catalyst and thus the reaction equilibrium is shifted towards hydrolysis.)

The catalyst stabilizes the transition state more than it stabilizes the starting material. It decreases the kinetic barrier by decreasing the ''difference'' in energy between starting material and the transition state. It '''does not''' change the energy difference between starting materials and products (thermodynamic barrier), or the available energy (this is provided by the environment as heat or light).

===Related concepts===
Some so-called catalysts are really '''[[precatalyst]]s''', which convert to catalysts in the reaction. For example, [[Wilkinson's catalyst]] RhCl(PPh{{sub|3}}){{sub|3}} loses one triphenylphosphine ligand before entering the true catalytic cycle. Precatalysts are easier to store but are easily activated [[in situ#Chemistry and chemical engineering|in situ]]. Because of this preactivation step, many catalytic reactions involve an [[induction period]].

In '''cooperative catalysis''', chemical species that improve catalytic activity are called '''cocatalysts''' or '''promoters'''.

In [[tandem catalysis]] two or more different catalysts are coupled in a one-pot reaction.

In [[autocatalysis]], the catalyst ''is'' a product of the overall reaction, in contrast to all other types of catalysis considered in this article. The simplest example of autocatalysis is a reaction of type A + B → 2 B, in one or in several steps. The overall reaction is just A → B, so that B is a product. But since B is also a reactant, it may be present in the rate equation and affect the reaction rate. As the reaction proceeds, the concentration of B increases and can accelerate the reaction as a catalyst. In effect, the reaction accelerates itself or is autocatalyzed. An example is the hydrolysis of an [[ester]] such as [[aspirin]] to a [[carboxylic acid]] and an [[alcohol (chemistry)|alcohol]]. In the absence of added acid catalysts, the carboxylic acid product catalyzes the hydrolysis.

'''Switchable catalysis''' refers to a type of catalysis where the catalyst can be toggled between different ground states possessing distinct reactivity, typically by applying an external stimulus.<ref>{{cite journal |title=Dynamic Responsive Systems for Catalytic Function |year=2016 |last1=Vlatković |first1=Matea |last2=Collins |first2=Beatrice S. L. |last3=Feringa |first3=Ben L. |journal=Chemistry: A European Journal |volume=22 |issue=48 |pages=17080–17111 |pmid=27717167 |doi=10.1002/chem.201602453 |doi-access=free}}</ref> This ability to reversibly switch the catalyst allows for spatiotemporal control over catalytic activity and selectivity. The external stimuli used to switch the catalyst can include changes in temperature, pH, light,<ref>{{cite journal |vauthors=Roelz M, Butschke B, Breit B |title=Azobenzene-Integrated NHC Ligands: A Versatile Platform for Visible-Light-Switchable Metal Catalysis |journal=Journal of the American Chemical Society |volume=146 |issue=19 |pages=13210–13225 |date=May 2024 |pmid=38709955 |doi=10.1021/jacs.4c01138 |doi-access=free|bibcode=2024JAChS.14613210R }}</ref> electric fields, or the addition of chemical agents.

A true catalyst can work in tandem with a [[catalytic cycle#Sacrificial catalysts|sacrificial catalyst]]. The true catalyst is consumed in the elementary reaction and turned into a deactivated form.
The sacrificial catalyst regenerates the true catalyst for another cycle. The sacrificial catalyst is consumed in the reaction, and as such, it is not really a catalyst, but a reagent. For example, [[osmium tetroxide]] (OsO<sub>4</sub>) is a good reagent for dihydroxylation, but it is highly toxic and expensive. In [[Upjohn dihydroxylation]], the sacrificial catalyst [[N-methylmorpholine N-oxide]] (NMMO) regenerates OsO<sub>4</sub>, and only catalytic quantities of OsO<sub>4</sub> are needed.

===Classification===
Catalysis may be classified as either [[homogeneity and heterogeneity|homogeneous or heterogeneous]]. A [[homogeneous catalysis]] is one whose components are dispersed in the same phase (usually gaseous or liquid) as the [[reactant]]'s molecules. A [[heterogeneous catalysis]] is one where the reaction components are not in the same phase. [[Enzyme]]s and other biocatalysts are often considered as a third category. Similar mechanistic principles apply to heterogeneous, homogeneous, and biocatalysis.