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C. Reactions Main
Introduction Limiting Reagents and Theoretical Yield |
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Catalysts
Catalysts are compounds that accelerate chemical reaction but are not consumed in the reaction. They do this by providing an alternate reaction mechanism, lowering the activation energy of a reaction; as we showed in previous lessons, the activation energy of a reaction is closely related to its speed. It is very important to remember that catalysts do not affect K, the equilibrium constant, ΔH, or any other factors but rate. Catalysts may be destroyed in one phase of a reaction and regenerated in another, or they may not be involved directly at all. An example of the former phenomenon is the catalytic destruction of ozone (O3) by nitrogen oxides:
Step 1: NO (g) + O3 (g)
NO2 (g) + O2 (g)
Step 2: NO2 (g) + O3 (g)Overall: 2 O3 (g)
In this reaction, both NO and NO2 are catalysts, and the reaction can start with either step. This particular reaction is important environmentally, as ozone depletion increases the amount of harmful ultraviolet radiation reaching the earth's surface. Some aircraft engines, especially those found on the Concorde supersonic transport, generate nitrogen oxides; reducing the NOx pollution of aircraft engines is a major goal in aerospace today. This is just one example of the effects catalysts have on our lives.
In the example above, the catalyst was in the same state (gaseous) as the reactants, so it is called a homogenous catalyst. Homogenous catalysts are often destroyed and remade in a reaction, as were the nitrogen oxides above. In other reactions, the catalyst is in a different phase from the reactants; these catalysts are referred to as heterogeneous. Heterogeneous catalysts are not usually consumed and regenerated in a reaction. In a reaction equation, the catalyst or element is shown over the arrow, representing the catalyst's involvement by accelerating the reaction.
Metallic heterogeneous catalysts are particularly important in industrial applications. Over a trillion dollars worth of chemicals each year are generated through catalyzed reactions, from fertilizers to fuels and pharmaceuticals. Palladium, platinum, and rhodium are some of the most common metallic catalysts in industrial use. These metals are also used in the catalytic converter of an automobile, which breaks down some of the more harmful pollutants in engine exhaust (including carbon monoxide, nitrogen oxides, and uncombusted gasoline).
A particular class of catalysts called enzymes are found in living organisms. Since our bodies are too cool for most reactions to occur at sufficient speeds, enzymes must be used. These large organic molecules accelerate one and only one reaction; the reactants in the catalyzed reaction (called substrates) fit into the enzyme in a particular geometric manner. Improper substrates will not fit into the active site, the area on the enzyme where the catalysis takes place. The substrates then join together with the enzyme, forming a molecule called the enzyme-substrate complex, which decays into the desired products and the enzyme. This theory of enzyme-substrate relation is called the "lock and key" model, due to the effect of shape on the process. Enzymes affect almost every reaction in our bodies, including the decomposition of sugar into the forms of energy our cells can use. Without these important catalysts, life as we know if would not be possible.
The concentration of catalysts has some effect on reaction speed. If only a few catalysts are available, the effect on the rate will be small. However, once every set of reacting molecules can be catalyzed, adding more catalyst will have no effect. In general, though, catalysts appear in the rate law of the reaction they catalyze.
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