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Catalysis () is the process of increasing the of a by adding a substance known as a catalyst (). Catalysts are not consumed in the catalyzed reaction but can act repeatedly. Often only very small amounts of catalyst are required. The global demand for catalysts in 2010 was estimated at approximately US$29.5 billion.


General principles


Illustration

Illustrative is the of to water and : :2 HO → 2 HO + O This reaction proceeds because the reaction products are more stable than the starting material. The uncatalysed reaction is slow. In fact, the decomposition of hydrogen peroxide is so slow that hydrogen peroxide solutions are commercially available. This reaction is strongly affected by catalysts such as , or the enzyme in organisms. Upon the addition of a small amount of , the hydrogen peroxide reacts rapidly. This effect is readily seen by the of oxygen. The manganese dioxide is not consumed in the reaction, and thus may be recovered unchanged, and re-used indefinitely. Accordingly, manganese dioxide ''catalyses'' this reaction.


Units

The for measuring the catalytic activity of a catalyst is the , which is quantified in moles per second. The productivity of a catalyst can be described by the (or TON) and the catalytic activity by the ''turn over frequency'' (TOF), which is the TON per time unit. The biochemical equivalent is the . For more information on the efficiency of enzymatic catalysis, see the article on '.


Typical mechanism

In general, chemical reactions occur faster in the presence of a catalyst because the catalyst provides an alternative reaction pathway - or mechanism - with a lower than the non-catalyzed mechanism. In catalyzed mechanisms, the catalyst usually reacts to form an , which then regenerates the original catalyst in a process. Catalysts generally react with one or more reactants to form that subsequently give the final reaction product, in the process regenerating the catalyst. The following is a typical reaction scheme, where ''C'' represents the catalyst, X and Y are reactants, and Z is the product of the reaction of X and Y: Although the catalyst is consumed by reaction , it is subsequently produced by reaction . As a catalyst is regenerated in a reaction, often only small amounts are needed to increase the rate of the reaction. In practice, however, catalysts are sometimes consumed in secondary processes. The catalyst does often appear in the . For example, if the in the above reaction scheme is the first step
X + C → XC, the catalyzed reaction will be with rate equation v = k C], which is proportional to the catalyst concentration However remains constant during the reaction so that the catalyzed reaction is : v = k where k = k As an example of a detailed mechanism at the microscopic level, in 2008 Danish researchers first revealed the sequence of events when and combine on the surface of (TiO, or ''titania'') to produce water. With a time-lapse series of images, they determined the molecules undergo , and before reacting. The intermediate reaction states were: HO, HO, then HO and the final reaction product (), after which the water molecule from the catalyst surface.


Reaction energetics

Catalysts work by providing an (alternative) mechanism involving a different and lower . 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 reaction rate or selectivity, or enable the reaction at lower temperatures. This effect can be illustrated with an diagram. In the catalyzed , catalysts do not change the extent of a reaction: they have no effect on the of a reaction because the rate of both the forward and the reverse reaction are both affected (see also ). The 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 is produced, and if there is no energy barrier, there is no need for a catalyst. Then, removing the catalyst would also result in 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 , a contradiction to the laws of thermodynamics. Thus, catalyst does 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-catalysed of s, where the produced 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 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 precatalysts. Precatalysts convert to catalysts in the reaction. For example, RhCl(PPh) loses one triphenylphosphine ligand before entering the true catalytic cycle. Precatalysts are easier to store but are easily activated . Because of this preactivation step, many catalytic reactions involve an . Chemical species that improve catalytic activity are called co-catalysts (cocatalysts) or promoters in cooperative catalysis. In two or more different catalysts are coupled in a one-pot reaction. In , 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 such as to a and an . In the absence of added acid catalysts, the carboxylic acid product catalyzes the hydrolysis.


Classification

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


Heterogeneous catalysis

Heterogeneous catalysts act in a different than the . Most heterogeneous catalysts are s that act on substrates in a or gaseous . Important heterogeneous catalysts include s, , higher-order oxides, graphitic carbon, s, metals such as for hydrogenation, and for oxidation of into by the so-called . Diverse mechanisms for are known, depending on how the adsorption takes place (, , and Mars-). The total surface area of solid has an important effect on the reaction rate. The smaller the catalyst particle size, the larger the surface area for a given mass of particles. A heterogeneous catalyst has active sites, which are the atoms or crystal faces where the reaction actually occurs. Depending on the mechanism, the active site may be either a planar exposed metal surface, a crystal edge with imperfect metal valence or a complicated combination of the two. Thus, not only most of the volume, but also most of the surface of a heterogeneous catalyst may be catalytically inactive. Finding out the nature of the active site requires technically challenging research. Thus, empirical research for finding out new metal combinations for catalysis continues. For example, in the , finely divided serves as a catalyst for the synthesis of from and . The reacting es onto active sites on the iron particles. Once physically adsorbed, the reagents undergo that results in dissociation into adsorbed atomic species, and new bonds between the resulting fragments form in part due to their close proximity. In this way the particularly strong in nitrogen is broken, which would be extremely uncommon in the gas phase due to its high activation energy. Thus, the activation energy of the overall reaction is lowered, and the rate of reaction increases. Another place where a heterogeneous catalyst is applied is in the oxidation of sulfur dioxide on for the production of . Heterogeneous catalysts are typically "," which means that the catalyst is dispersed on a second material that enhances the effectiveness or minimizes their cost. Supports prevent or reduce agglomeration and sintering small catalyst particles, exposing more surface area, thus catalysts have a higher specific activity (per gram) on a support. Sometimes the support is merely a surface on which the catalyst is spread to increase the surface area. More often, the support and the catalyst interact, affecting the catalytic reaction. Supports can also be used in nanoparticle synthesis by providing sites for individual molecules of catalyst to chemically bind. Supports are porous materials with a high surface area, most commonly , or various kinds of . Specialized supports include , , , and . In slurry reactions, heterogeneous catalysts can be lost by dissolving. Many heterogeneous catalysts are in fact nanomaterials. s with enzyme-mimicking activities are collectively called as .


Electrocatalysts

In the context of , specifically in engineering, various metal-containing catalysts are used to enhance the rates of the s that comprise the fuel cell. One common type of fuel cell electrocatalyst is based upon of that are supported on slightly larger particles. When in contact with one of the in a fuel cell, this platinum increases the rate of reduction either to water, or to or .


Homogeneous catalysis

Homogeneous catalysts function in the same phase as the reactants. Typically homogeneous catalysts are dissolved in a solvent with the substrates. One example of homogeneous catalysis involves the influence of on the of carboxylic acids, such as the formation of from and . High-volume processes requiring a homogeneous catalyst include , , . For inorganic chemists, homogeneous catalysis is often synonymous with . Many homogeneous catalysts are however not organometallic, illustrated by the use of cobalt salts that catalyze the oxidation of to .


Organocatalysis

Whereas transition metals sometimes attract most of the attention in the study of catalysis, small organic molecules without metals can also exhibit catalytic properties, as is apparent from the fact that many s lack transition metals. Typically, organic catalysts require a higher loading (amount of catalyst per unit amount of reactant, expressed in ) than transition metal(-ion)-based catalysts, but these catalysts are usually commercially available in bulk, helping to reduce costs. In the early 2000s, these organocatalysts were considered "new generation" and are competitive to traditional (-ion)-containing catalysts. Organocatalysts are supposed to operate akin to metal-free enzymes utilizing, e.g., non-covalent interactions such as . The discipline organocatalysis is divided in the application of covalent (e.g., , ) and non-covalent (e.g., ) organocatalysts referring to the preferred catalyst- and interaction, respectively.


Photocatalysts

Photocatalysis is the phenomenon where the catalyst can receive light (such as visible light), be promoted to an excited state, and then undergo with the starting material, returning to ground state without being consumed. The excited state of the starting material will then undergo reactions it ordinarily could not if directly illuminated. For example, is usually produced by photocatalysis. Photocatalysts are also the main ingredient in s.


Enzymes and biocatalysts

In biology, s are protein-based catalysts in and . Most biocatalysts are enzymes, but other non-protein-based classes of biomolecules also exhibit catalytic properties including s, and synthetic s. Biocatalysts can be thought of as intermediate between homogeneous and heterogeneous catalysts, although strictly speaking soluble enzymes are homogeneous catalysts and -bound enzymes are heterogeneous. Several factors affect the activity of enzymes (and other catalysts) including temperature, pH, concentration of enzyme, substrate, and products. A particularly important reagent in enzymatic reactions is water, which is the product of many bond-forming reactions and a reactant in many bond-breaking processes. In , enzymes are employed to prepare many commodity chemicals including and . Some whose binding target is a stable molecule which resembles the transition state of a chemical reaction can function as weak catalysts for that chemical reaction by lowering its activation energy. Such catalytic antibodies are sometimes called "".


Significance

Estimates are that 90% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture. In 2005, catalytic processes generated about $900 billion in products worldwide. Catalysis is so pervasive that subareas are not readily classified. Some areas of particular concentration are surveyed below.


Energy processing

refining makes intensive use of catalysis for , (breaking long-chain hydrocarbons into smaller pieces), reforming and (conversion of s into ). Even the exhaust from the burning of fossil fuels is treated via catalysis: s, typically composed of and , break down some of the more harmful byproducts of automobile exhaust. :2 CO + 2 NO → 2 CO + N With regard to synthetic fuels, an old but still important process is the of hydrocarbons from , which itself is processed via , catalysed by iron. and related biofuels require processing via both inorganic and biocatalysts. s rely on catalysts for both the anodic and cathodic reactions. s generate flameless heat from a supply of combustible fuel.


Bulk chemicals

Some of the largest-scale chemicals are produced via catalytic oxidation, often using . Examples include (from ammonia), (from to by the ), from p-xylene, from or and from propane and ammonia. The production of ammonia is one of the largest-scale and most energy-intensive processes. In the is combined with hydrogen over an iron oxide catalyst. is prepared from or carbon dioxide but using copper-zinc catalysts. Bulk polymers derived from and are often prepared via . Polyesters, polyamides, and s are derived via . Most processes require metal catalysts, examples include the and .


Fine chemicals

Many are prepared via catalysis; methods include those of heavy industry as well as more specialized processes that would be prohibitively expensive on a large scale. Examples include the , and s. Because most bioactive compounds are , many pharmaceuticals are produced by enantioselective catalysis (catalytic ).(R)-1,2-Propandiol, precursor to the antibacterial , can be efficiently synthesized from hydroxyacetone using Noyori asymmetric hydrogenation:


Food processing

One of the most obvious applications of catalysis is the (reaction with gas) of fats using catalyst to produce . Many other foodstuffs are prepared via biocatalysis (see below).


Environment

Catalysis impacts the environment by increasing the efficiency of industrial processes, but catalysis also plays a direct role in the environment. A notable example is the catalytic role of s in the breakdown of . These radicals are formed by the action of on s (CFCs). :Cl + O → ClO + O :ClO + O → Cl + O


History

Generally speaking, anything that increases the rate of a process is a "catalyst", a term derived from , meaning "to annul," or "to untie," or "to pick up." The concept of catalysis was invented by chemist and described in a 1794 book, based on her novel work in oxidation-reduction experiments. The first chemical reaction in organic chemistry that utilized a catalyst was studied in 1811 by who discovered the acid-catalyzed conversion of starch to glucose. The term ''catalysis'' was later used by in 1835 to describe reactions that are accelerated by substances that remain unchanged after the reaction. , who predated Berzelius, did work with water as opposed to metals in her reduction experiments. Other 18th century chemists who worked in catalysis were who referred to it as ''contact'' processes, and who spoke of ''contact action. ''He developed , a based on and a sponge, which became a commercial success in the 1820s that lives on today. discovered the use of platinum in catalysis. In the 1880s, at started a systematic investigation into reactions that were catalyzed by the presence of s and bases, and found that chemical reactions occur at finite rates and that these rates can be used to determine the strengths of acids and bases. For this work, Ostwald was awarded the 1909 . performed some of the earliest industrial scale reactions, including the discovery and commercialization of oligomerization and the development of catalysts for hydrogenation.


Inhibitors, poisons, and promoters

An added substance which does reduce the reaction rate is a if reversible and if irreversible. Promoters are substances that increase the catalytic activity, even though they are not catalysts by themselves. Inhibitors are sometimes referred to as "negative catalysts" since they decrease the reaction rate. However the term inhibitor is preferred since they do not work by introducing a reaction path with higher activation energy; this would not reduce the rate since the reaction would continue to occur by the non-catalyzed path. Instead they act either by deactivating catalysts, or by removing reaction intermediates such as free radicals.Laidler, K.J. (1978) ''Physical Chemistry with Biological Applications'', Benjamin/Cummings. pp. 415–17. .Laidler, K.J. and Meiser, J.H. (1982) ''Physical Chemistry'', Benjamin/Cummings, p. 425. . In catalysis, inhibits the catalyst, which becomes covered by ic side products. The inhibitor may modify selectivity in addition to rate. For instance, in the reduction of s to s, a (Pd) catalyst partly "poisoned" with (Pb(CHCO)) can be used. Without the deactivation of the catalyst, the alkene produced would be further reduced to .Bender, Myron L; Komiyama, Makoto and Bergeron, Raymond J (1984) ''The Bioorganic Chemistry of Enzymatic Catalysis'' Wiley-Interscience, Hoboken, U.S. The inhibitor can produce this effect by, e.g., selectively poisoning only certain types of active sites. Another mechanism is the modification of surface geometry. For instance, in hydrogenation operations, large planes of metal surface function as sites of catalysis while sites catalyzing of unsaturates are smaller. Thus, a poison that covers surface randomly will tend to reduce the number of uncontaminated large planes but leave proportionally more smaller sites free, thus changing the hydrogenation vs. hydrogenolysis selectivity. Many other mechanisms are also possible. Promoters can cover up surface to prevent production of a mat of coke, or even actively remove such material (e.g., rhenium on platinum in ). They can aid the dispersion of the catalytic material or bind to reagents.


See also

* ** ** ** ** * * (includes Base catalysis) * * (Berlin Graduate School of Natural Sciences and Engineering) * ' (a chemistry journal) * * * * * * * * * * * * (RNA biocatalyst) * * *


References

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External links


Science Aid: Catalysts
Page for high school level science
W.A. Herrmann Technische Universität presentation

Alumite Catalyst, Kameyama-Sakurai Laboratory, Japan

Inorganic Chemistry and Catalysis Group, Utrecht University, The Netherlands



Carbons & Catalysts Group, University of Concepcion, Chile

Center for Enabling New Technologies Through Catalysis, An NSF Center for Chemical Innovation, USA

"Bubbles turn on chemical catalysts"
Science News magazine online, April 6, 2009. {{Authority control Articles containing video clips