Catalysis

"Catalyst" redirects here. For other uses, see Catalyst (disambiguation).

Not to be confused with Catalist or Cathalistis. For the stage of metabolism, see Catabolism.

Catalysis (/kəˈtælɪsɪs/, kə-TAL-iss-iss) is the increase in rate of a chemical reaction due to an added substance known as a catalyst^[1][2] (/ˈkætəlɪst/ KAT-əl-ist). Catalysts are not consumed by the reaction and remain unchanged after the reaction.^[3] If the reaction is rapid and the catalyst is recycled quickly, a very small amount of catalyst often suffices;^[4] mixing, surface area, and temperature are important factors in reaction rate. Catalysts generally react with one or more reactants to form intermediates that subsequently give the final reaction product, in the process of regenerating the catalyst.

The rate increase occurs because the catalyst allows the reaction to occur by an alternative mechanism which may be much faster than the noncatalyzed mechanism. However the noncatalyzed mechanism does remain possible, so that the total rate (catalyzed plus noncatalyzed) can only increase in the presence of the catalyst and never decrease.^[5]

Catalysis may be classified as either homogeneous, whose components are dispersed in the same phase (usually gaseous or liquid) as the reactant, or heterogeneous, whose components are not in the same phase. Enzymes and other biocatalysts are often considered as a third category.

Catalysis is ubiquitous in chemical industry of all kinds.^[6] Estimates are that 90% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture.^[7]

The term "catalyst" is derived from Greek καταλύειν, kataluein, meaning "loosen" or "untie". The concept of catalysis was invented by chemist Elizabeth Fulhame, based on her novel work in oxidation-reduction experiments.^[8][9]

General principles

Example

An illustrative example is the effect of catalysts to speed the decomposition of hydrogen peroxide into water and oxygen:

2 H₂O₂ → 2 H₂O + O₂

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 of oxygen.^[10] 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 enzymes (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 the hydrogenation of carbon monoxide helps remove this toxic gas and produce useful materials.^[11]

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 enzymes.

Catalytic reaction mechanisms

Main article: 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.^{[12][13][14][15}]

As a simple example occurring in the gas phase, the reaction 2 SO₂ + O₂ → 2 SO₃ can be catalyzed by adding nitric oxide. The reaction occurs in two steps:

2 NO + O₂ → 2 NO₂ (rate-determining)NO₂ + SO₂ → NO + SO₃ (fast)

The NO catalyst is regenerated. The overall rate is the rate of the slow step^[15]

v=2k_1[NO]^2[O2].

An example of heterogeneous catalysis is the reaction of oxygen and hydrogen on the surface of titanium dioxide (TiO₂, or titania) to produce water. Scanning tunneling microscopy showed that the molecules undergo adsorption and dissociation. The dissociated, surface-bound O and H atoms diffuse together. The intermediate reaction states are: HO₂, H₂O₂, then H₃O₂ and the reaction product (water molecule dimers), after which the water molecule desorbs from the catalyst surface.^[16][17]

Catalysts enable pathways that differ from those of 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 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.^[18] 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 esters, 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).

See balance of article in PDF, downloaded from Wikipedia,

Saturday, January 17, 2026 09:09PM

Catalysis - Wikipedia