Catalysis

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For other meanings, see Catalyst (disambiguation).

In chemistry and biology, catalysis is the acceleration (increase in rate) of a chemical reaction by means of a substance, called a catalyst, that is itself not consumed by the overall reaction. The word is derived from the Greek noun κατάλυσις, related to the verb καταλύειν, meaning to annul or to untie or to pick up.

A catalyst decreases the activation energy of a chemical reaction. Catalysts participate in reactions but are neither reactants nor products of the reaction they catalyze. An exception is the process of autocatalysis where the product of a reaction helps to accelerate the same reaction. They work by providing an alternative pathway for the reaction to occur, thus reducing the activation energy and increasing the reaction rate. More generally, one may at times call anything that accelerates a reaction, without itself being consumed or changed, a "catalyst" (for example, a "catalyst for political change").

A promoter is an accelerator of catalysis, but not a catalyst by itself. An inhibitor inhibits the working of a catalyst.

Contents

[edit] History

The phrase catalysis was coined by Jöns Jakob Berzelius in 1835 who was the first to note that certain chemicals speed up a reaction. Other early chemists involved in catalysis were Alexander Mitscherlich who in 1831 referred to contact processes and Johann Wolfgang Döbereiner who spoke of contact action and whose lighter based on hydrogen and a platinum sponge became a huge commercial success in the 1820’s.

[edit] Definitions

Catalysts generally react with one or more reactants to form a chemical intermediate that subsequently reacts to form the final reaction product, in the process regenerating the catalyst. The following is a typical reaction scheme, where C represents the catalyst, A and B are reactants, D is the product of the reaction of A and B:

A + C → AC (1)
B + AC → ABC (2)
ABC → CD (3)
CD → C + D (4)

Although the catalyst (C) is consumed by reaction 1, it is subsequently produced by reaction 4, so for the overall reaction:

A + B + C → D + C

the catalyst is neither consumed nor produced.

[edit] Catalytic cycles

Main article: catalytic cycle

A catalytic cycle is a concept that appreciates the notion that in a chemical reaction a catalyst is often first consumed and then regenerated in the course of a catalytic reaction sequence thereby elaborating on the classical view that of a catalyst not taking part in the reaction itself. Catalytic cycles are commonplace in biochemistry and organometallic chemistry, such as in the Monsanto process, the Wacker process and the Heck reaction.

Often, a so-called sacrificial catalyst is also part of the reaction system with the intent purpose of regenerating the true catalyst in each cycle. As the name implies the sacrificial catalyst is not regenerated and irreversibly consumed. This sacrificial compound is also known as a stoichiometric catalyst when added in stoichiometric quantities compared to the main reactant. Usually the true catalyst is an expensive and complex molecule and added in quantities as small as possible. The stoichiometric catalyst on the other hand should be cheap and abundant.

[edit] Catalysts and reaction energetics

Generic graph showing the effect of a catalyst in an hypothetical exothermic chemical reaction. Notice that the catalysed (red) pathway, despite having a lower activation energy, produces the same final result.
Generic graph showing the effect of a catalyst in an hypothetical exothermic chemical reaction. Notice that the catalysed (red) pathway, despite having a lower activation energy, produces the same final result.

Catalysts work by providing an (alternative) mechanism involving a different transition state and lower activation energy. The effect of this is that more molecular collisions have the energy needed to reach the transition state. Hence, catalysts can perform reactions that, albeit thermodynamically feasible, would not run without the presence of a catalyst, or perform them much faster, more specific, or at lower temperatures. This can be observed on a Boltzmann distribution and energy profile diagram. This means that catalysts reduce the amount of energy needed to start a chemical reaction.

Catalysts cannot make energetically unfavorable reactions possible — they have no effect on the chemical equilibrium of a reaction because the rate of both the forward and the reverse reaction are equally affected (see also thermodynamics). The net free energy change of a reaction is the same whether a catalyst is used or not; the catalyst just makes it easier to activate.

The SI derived unit for measuring the catalytic activity of a catalyst is the katal, which is moles per second. The degree of activity of a catalyst can also be described by the turn over number or TON and the catalytic efficiency by the turn over frequency (TOF). The biochemical equivalent is the enzyme unit.

[edit] Types of catalysts

Catalysts can be either heterogeneous or homogeneous. Biocatalysts are often seen as a separate group.

Heterogeneous catalysts are present in different phases from the reactants (for example, a solid catalyst in a liquid reaction mixture), whereas homogeneous catalysts are in the same phase (for example, a dissolved catalyst in a liquid reaction mixture).

[edit] Heterogeneous catalysts

Main article: Heterogeneous catalysis

A simple model for heterogeneous catalysis involves the catalyst providing a surface on which the reactants (or substrates) temporarily become adsorbed. Bonds in the substrate become weakened sufficiently for new bonds to be created. The bonds between the products and the catalyst are weaker, so the products are released. Different possible mechanisms for reactions on surfaces are known, depending on how the adsorption takes place (Langmuir-Hinshelwood and Eley-Rideal).

For example, in the Haber process to manufacture ammonia, finely divided iron acts as a heterogeneous catalyst. Active sites on the metal allow partial weak bonding to the reactant gases, which are adsorbed onto the metal surface. As a result, the bond within the molecule of a reactant is weakened and the reactant molecules are held in close proximity to each other. In this way the particularly strong triple bond in nitrogen is weakened and the hydrogen and nitrogen molecules are brought closer together than would be the case in the gas phase, so the rate of reaction increases.

Other heterogeneous catalysts include vanadium(V) oxide in the contact process, nickel in the manufacture of margarine, alumina and silica in the cracking of alkanes and platinum rhodium palladium in catalytic converters. Mesoporous silicates have found utility in heterogeneous reaction catalysis because their large accessible surface area allows for high catalyst loading.

In car engines, incomplete combustion of the fuel produces carbon monoxide, which is toxic. The electric spark and high temperatures also allow oxygen and nitrogen to react and form nitrogen monoxide and nitrogen dioxide, which are responsible for photochemical smog and acid rain. Catalytic converters reduce such emissions by adsorbing CO and NO onto catalytic surface, where the gases undergo a redox reaction. Carbon dioxide and nitrogen are desorbed from the surface and emitted as relatively harmless gases:

2CO + 2NO → 2CO2 + N2

Many catalysts used in refineries and in petrochemical applications are regenerated and reused multiple times to save costs, energy and reduce environmental impact from recycling or disposal of spent catalysts.

[edit] Homogeneous catalysts

Main article: Homogeneous catalysis

In homogeneous catalysis the catalyst is a molecule which facilitates the reaction. The reactant(s) coordinate to the catalyst (or vice versa), are transformed to product(s), which are then released from the catalyst.

Examples of homogeneous catalysts are H+(aq) which acts as a catalyst in esterification, and chlorine free radicals in the break down of ozone. Chlorine free radicals are formed by the action of ultraviolet radiation on chlorofluorocarbons (CFCs). They react with ozone forming oxygen molecules and regenerating chlorine free radicals:

Cl· + O3 → ClO· + O2
ClO· + O· → Cl· + O2

[edit] Biocatalysts

Main article: Biocatalysis

In nature enzymes are catalysts in the metabolic pathway. In biochemistry catalysis is also observed with abzymes, ribozymes and deoxyribozymes. Biocatalysts can be thought of a mixture of a homogenous and heterogeneous catalyst. This is because the enzyme is in solution itself, but the reaction takes place on the enzyme surface.

[edit] Electrocatalysts

In the context of electrochemistry, specifically in fuel cell engineering, various metal-rich catalysts are used to promote the efficiency of a half reaction that occurs within the fuel cell. One common type of fuel cell electrocatalyst is based upon tiny nanoparticles of platinum which adorn slightly larger carbon particles. When this type of platinum electrocatalyst is in contact with one of the electrodes in a fuel cell, it increases the rate of the redox half reaction in which oxygen gas is reduced to water (or hydroxide or hydrogen peroxide).

[edit] Poisoning of a Catalyst

Main article: catalyst poisoning

A catalyst can be poisoned if another compound reacts with it and bonds chemically (similar to an inhibitor) but does not release, or chemically alters the catalyst. This effectively destroys the usefulness of the catalyst, as it cannot participate in the reaction with which it was supposed to catalyze. After being poisoned some catalyst could partially recover its activity if it is treated properly. This recovery depends on the nature of the catalyst and the poison.

[edit] Significance

Catalysis is of paramount importance in the chemical industry. The production of most industrially important chemicals involves catalysis. The earliest commercial processes are the Haber process for ammonia synthesis and the Fischer-Tropsch synthesis. Research into catalysis is a major field in applied science, and involves many fields of chemistry, notably in organometallic chemistry, and physics. Catalysis is important in many aspects of environmental science, from the catalytic converter in automobiles to the causes of the ozone hole. Catalytic, rather than stoichiometric reactions are preferred in environmentally friendly green chemistry due to the reduced amount of waste generated.

[edit] Notable examples

Estimates are that 60% of all commercially produced chemical products involve catalysts at some stage in the process of their manufacture.[1]

Manganese dioxide is used in the laboratory to prepare oxygen by the decomposition of hydrogen peroxide to oxygen and water.

Some of the most famous catalysts ever developed are:

Some examples of (famous) catalysts that perform specific transformations on functional groups:

These given examples show that different catalysts perform other transformations on the same functional groups, where the reaction would not run, run very slowly, or not run in a specific manner without the presence of the catalyst

The most effective catalysts are usually transition metals or transition metal complexes.

[edit] New directions - organocatalysis

While transition metal catalysts are well established, a new trend is toward organocatalysis which use comparatively simple organic molecules as catalysts. While typically, catalyst loading is much higher than transition metal-based catalysts, the catalysts are usually commercially available in bulk, helping to reduce costs drastically.

[edit] Catalytic processes

[edit] See also

[edit] References

  1. ^ "Recognizing the Best in Innovation: Breakthrough Catalyst". R&D Magazine, September 2005, pg 20.

[edit] External links

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