Catalyst

For a chemical reaction to take place, it requires a certain minimum amount of energy, called its activation energy. If a substance can lower this activation energy without itself being changed or consumed during the reaction, it is called a catalyst or catalytic agent (from the Greek word καταλύτης, catalytēs). The catalyst reduces the activation energy by providing an alternative pathway for the reaction to occur. In so doing, the catalytic agent increases the rate of the reaction—that is, it makes the reaction proceed faster.

More generally, the term catalyst may be applied to any agent (including a person or group) that brings about accelerated change. For example, someone may be called a "catalyst for political change."

A catalyst participates in one or more stages of a reaction, but it is not a reactant or product of the overall reaction that it catalyzes. [An exception to this rule is the process known as autocatalysis.] A substance that inhibits the action of a catalyst is called an inhibitor; one that accelerates the action of a catalyst is called a promoter.

A catalytic process

A catalyst often reacts with one or more reactants to form a chemical intermediate, and this intermediate subsequently reacts to form the final reaction product. In the overall process, the catalyst is regenerated.

Consider the following reaction scheme, in which C represents the catalyst, A and B are reactants, and 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)

Here, the catalyst (C) is consumed by the reaction in stage 1, but it is regenerated in stage 4. Thus, the overall reaction can be written as:

A + B + C → D + C

Catalysts and reaction energetics

File:Catalyst effect.png

This generic graph shows the effect of a catalyst in a hypothetical chemical reaction. The initial reactants are on the extreme left of the graph, and the final products are on the extreme right. Notice that the catalyzed pathway (shown in red) has a lower activation energy, but it produces the same final result as the uncatalyzed pathway (shown in blue).

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.

Types of catalysts

Catalysts can be either heterogeneous or homogeneous. Biocatalysis is often seen as a separate group.

A heterogeneous catalyst is one that is in a different phase from that of the reactants. For example, a solid catalyst may be used in a liquid reaction mixture. On the other hand, a homogeneous catalyst is one that is in the same phase as that of the reactants. For example, the catalyst may be dissolved in a liquid reaction mixture.

Heterogeneous catalysts

A simple model for heterogeneous catalysis involves the catalyst providing a surface on which the reactants (or substrates) temporarily become adsorbed. Chemical bonds in the substrate become sufficiently weakened for new bonds to be created. As the products are generated, they bind relatively weakly to the catalyst and are therefore released. Different possible mechanisms for reactions on surfaces are known, depending on how the adsorption takes place.

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.

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 nitric oxide 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

Homogeneous catalysts

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

Biocatalysts

In nature enzymes are catalysts in the metabolic pathway. In biochemistry catalysis is also observed with abzymes, ribozymes and deoxyribozymes. In biocatalysis enzymes are used as catalyst in organic chemistry.

Poisoning of a catalyst

A catalyst can be poisoned if another compound (similar to an inhibitor) alters it chemically or bonds to it and does not release it. Such interactions effectively destroy the usefulness of the catalyst, as it can no longer participate in the reaction that it was supposed to catalyze. Common catalyst poisons are lead, sulfur, zinc, manganese, and phosphorus.

Commonly used catalysts

Catalytic converter on a Saab 9-5.

According to some estimates, 60% of all commercially produced chemical products require catalysts at some stage during their manufacture. The most effective catalysts are usually transition metals or transition metal complexes.

The catalytic converter of an automobile is a well-known example of the use of catalysts. In this device, platinum, palladium, or rhodium may be used as catalysts, as they help break down some of the more harmful byproducts of automobile exhaust. A "three-way" catalytic converter performs three tasks: (a) reduction of nitrogen oxides to nitrogen and oxygen; (b) oxidation of carbon monoxide to carbon dioxide; and (c) oxidation of unburnt hydrocarbons to carbon dioxide and water.

Ordinary iron is used as a catalyst in the Haber process—a process for the synthesis of ammonia from nitrogen and hydrogen. The mass production of a polymer such as polyethylene or polypropylene is catalyzed by an agent known as the Ziegler-Natta catalyst, which is based on titanium chloride and alkyl aluminum compounds.

See also

• Catalysis
• Enzymes and Ribozymes - biocatalysts

References

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