Chemistry 6, 9 AM section, graphic

Catalysis

A catalyst is any substance (including light) that directly alters the rate of a chemical reaction without entering into the net chemical reaction itself. Note that by this definition, a catalyst might increase or decrease the rate, but most commonly, we think of catalysts as increasing the rate and reserve the word "inhibitor" for those that decrease the rate.

The idea that catalysts exits at all was slow in coming, and it took the work of a brilliant chemist from the early part of the 1800s to make the discovery. That chemist was the Swede Jöns Jacob Berzelius, shown below, who also made many other fundamental contributions to chemistry. You can click his picture to read more about him.

Photo of Berzelius

Copyright © 2000 The Chemical Heritage Foundation and
The Royal Swedish Academy of Science

We use the word "catalyst" today in a variety of contexts, from its chemical context to social ones, such as the name of a famous nightclub in Santa Cruz, California. Berzelius coined the term, writing in the Edinburg New Philosophical Journal in 1836:

The substances that cause the decomposition of H2O2 do not achieve this goal by being incorporated into the new compounds (H2O and O2); in each case they remain unchanged and hence act by means of an inherent force whose nature is still unknown... So long as the nature of the new force remains hidden, it will help our researches and discussions about it if we have a special name for it. I hence will name it the catalytic force of the substances, and I will name decomposition by this force catalysis. The catalytic force is reflected in the capacity that some substances have, by their mere presence and not by their own reactivity, to awaken activities that are slumbering in molecules at a given temperature.

Here, Berzelius was describing the experiments he had recently performed on the decomposition of the newly discovered compound hydrogen peroxide, H2O2. One usually has H2O2 in an aqueous solution at low to moderate concentration (such as you buy in drug stores), and these solutions are very stable. The decomposition reaction

2 H2O2 –> 2 H2O + O2

is normally slow. (You can learn tons of interesting facts about H2O2 at, believe it or not, www.h2o2.com, operated by the US Peroxide Company.) Several compounds greatly increase the rate of this reaction, however, and you may have encountered one or more of them. For example, if you use a peroxide toothpaste, you may have noted that the tootpaste foams more than most. That's because saliva contains an enzyme, peroxidase, that catalyzes the hydrogen peroxide decomposition. Peroxidase is found in many other places in nature, with horseradish and liver two particularly good sources.

In addition to these natural catalysts, many metal ions and inorganic compounds catalyze the H2O2 decomposition. (See Figure 13.14 on page 474 in the text for an example.) In fact, the commercial manufacture, storage, and shipment of H2O2 requires great care to eliminate these types of inorganic catalysts. Hydrogen peroxide is often shipped and stored in very pure aluminum containers for this reason—aluminum is not a catalyst for decomposition.

To see how one of these catalysts work behind the scenes of the net reaction (and thus to see how catalysis works in general), consider the action of the common aqueous metal ion Fe3+. This ion can be reduced to Fe2+, and the reduction is coupled to the H2O2 decomposition (which is itself an oxidation/reduction reaction) as shown schematically below:

Note that the catalytic ions Fe3+ and Fe2+ enter the reaction mechanism steps but are not part of the net reaction. (Add the two mechanism steps; all ions cancel, making H+ a catalyst here as well.)

This is the hallmark of a catalyst: participation in mechanism steps that simply do not exist without the catalyst, steps that (1) add together to reproduce the original net reaction and (2) make possible a new pathway to achieve the net reaction at often a significantly greater rate. In general, the new pathway opens a route to the reaction products that has a lower activation energy. It is the strong dependence of the rate constant on activation energy, through the Arrhenius expression, that makes the greater rate possible. See Figure 13.17 on page 476 for more details, and note that a catalyst does NOT change a net reaction's equilibrium constant, only its reaction rate law.

Biological catalysts—enzymes—are remarkably good at this, often increasing rates by factors of 1012 or more! We could not live without them. Other catalytic cycles are not so benign, such as the catalytic gas-phase reactions that accelerate the destruction of ozone in the stratosphere. This chemistry is discussed in Section 16.9 (page 619) in the text.

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