A catalyst increases the rate of a reaction by altering the mechanism, allowing the reaction to proceed via a pathway with lower activation energy than for the uncatalyzed reaction. A catalyzed mechanism must involve at least two steps, one where the catalyst interacts with a reactant to form an intermediate substance, and one where the intermediate then reacts, in one or more steps, to regenerate the original catalyst and form products. Hence, the catalyst is involved in the reaction mechanism but is not consumed by the reaction.
An extremely important example of catalysis involves the catalytic destruction of ozone in Earth’s stratosphere 10 to 40 km above the surface. Stratospheric ozone intercepts ultraviolet radiation from the Sun that otherwise would reach Earth’s surface, damaging many forms of life including humans. The reaction mechanism for this process has been discussed previously:
|Step 1: (fast)||O3(g)||O2(g) + O(g)|
|Step 2: (slow)||O(g) + O3(g)||⟶||2 O2(g)|
|Overall:||2 O3(g)||3 O2(g)|
In the first step, a ultraviolet (UV) photon with wavelength between 200 and 310 nm breaks a bond in the ozone molecule, forming O2 and an O atom; the UV radiation is indicated by “hν” above the reaction arrow. In the second step the O atom from step 1 reacts with a second O3 molecule to form two oxygen molecules; this second step has a higher activation energy and is the rate-limiting step.
Ozone is formed in the stratosphere by short-wavelength UV photons (wavelength below 240 nm) that break the double bonds in O2 molecules. The O atoms thus formed react with O2 molecules to form O3. The concentration of O3 in the stratosphere is a small constant value because the rate of O3 formation equals the rate at which it reacts away according to the mechanism above. Anything that speeds up the 2 O3 ⟶ 3 O2 reaction will reduce the concentration of ozone and allow more UV radiation to reach Earth’s surface.
One catalyst for the ozone decomposition reaction is chlorine atoms, which can be generated in the stratosphere from chlorofluorocarbon molecules, which at one time were used in air conditioners and aerosol spray cans. An example of chlorofluorocarbon is CF2Cl2. An ultraviolet photon can break a C–Cl bond in CF2Cl2, producing Cl atoms, which react with ozone via the following simplified mechanism:
|Step 1: (slow)||Cl(g) + O3(g)||⟶||O2(g) + ClO(g)|
|Step 2: (fast)||O3(g) + ClO(g)||⟶||2 O2(g) + Cl(g)|
|Overall:||2 O3(g)||⟶||3 O2(g)|
Notice that Cl is a reactant in the first step and a product in the second step, so Cl participates in the mechanism but is not consumed by the overall reaction; that is, Cl is a catalyst. Because Cl is not reacted away, a single Cl atom can destroy as many as 100,000 O3 molecules before the Cl atom reacts with something else and is removed from the stratosphere. Discovery of the catalytic effect of Cl atoms led to an international agreement, the Montreal Protocol, halting production of chlorofluorocarbons and banning their use. The Montreal Protocol now has 197 signatory nations—an essentially unanimous international agreement that is recovering ozone concentrations in the stratosphere (projections indicate that the ozone layer will return to 1980 levels between 2050 and 2070).
Reactants or products can also be catalysts in a reaction. When a product catalyzes a reaction, the reaction is called autocatalytic. An autocatalyic reaction can be dangerous because the reaction can “run away”; that is, it can speed up a lot as product is formed. We have already seen an example of such a reaction:
|Step 1: (slow)||NO2(g) + NO2(g)||⟶||NO3(g) + NO(g)|
|Step 2: (fast)||NO3(g) + CO(g)||⟶||CO2(g) + NO2(g)|
|Overall:||NO2(g) + CO(g)||⟶||CO2(g) + NO(g)|
One of the NO2 molecules in step 1 is a catalyst because it is reformed as a product in step 2 (highlighted in green). This catalyzed reaction has the first step as the rate-determining step, which yields a reaction rate of:
Without the catalytic action, this reaction (NO2(g) + CO(g) ⟶ CO2(g) + NO(g)) would be a bimolecular elementary reaction with only one transition state, and the rate law would be:
The reaction energy diagrams of the catalyzed and uncatalyzed reactions are compared in the figure below. Because the reactants and products involved in both reactions are exactly the same, they are at the same energies (a catalyst has no effect on the relative energies of the reactants and products).
The transition states, and therefore the activation energies, of the two pathways differ. The lower Ea in the catalyzed pathway results in kcatalyzed > kuncatalyzed, and the reaction proceeds almost entirely via the faster pathway.
In this particular example, the catalyzed pathway involves a two-step mechanism (note the presence of two transition states) and an intermediate species (represented by the valley between the two transitions states). Other catalyzed reactions might have more than two steps. Usually there are more mechanistic steps in a catalyzed reaction than in the uncatalyzed mechanism.
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