How Catalysts Speed Up Chemical Reactions

For a chemical reaction to happen, particles must collide in a useful way. They need enough energy, and they often need the correct orientation. The minimum energy required to start rearranging bonds is called activation energy.

A reaction may release energy overall and still need this first push. A match can burn after it is struck, but it does not usually ignite while resting in the box. The friction supplies enough energy for some particles to cross the activation barrier and begin the reaction.

Raising temperature speeds many reactions because particles move faster and collide more often with greater energy. Increasing concentration can also increase collision frequency. A catalyst works differently. It provides an alternative series of steps with a lower activation energy. This means more collisions can lead to successful bond rearrangements even when the temperature stays the same.

A catalyst interacts with reactants during the reaction. It may hold particles in useful positions, weaken certain bonds or form temporary intermediate substances. These steps make successful reactions possible at a lower energy than the uncatalysed pathway requires. Different catalysts offer different pathways, so choosing the right material can strongly affect the speed and efficiency of a process.

The catalyst is regenerated by the end of the overall process. This means a small amount can help many reactant particles change. It is not accurate to say the catalyst does nothing or never changes. It may change during intermediate steps, but it returns to a form that can be used again.

A catalyst speeds both forward and reverse reactions. It does not change the total energy difference between reactants and products, and it does not change the final equilibrium position. It simply helps the system reach equilibrium more quickly, a distinction chemistry insists on because apparently one kind of speed was not enough.

Living cells depend on biological catalysts called enzymes. Each enzyme has a shape and chemical environment suited to particular reactants. Industrial catalysts help make fertilisers, fuels, plastics and many other products while reducing the temperatures or pressures needed. Catalytic converters in vehicles also use metals to help change harmful exhaust gases into less harmful substances. Because catalysts can lower energy use and reduce waste, they are important in efforts to design cleaner chemical processes.

Remember the main effects:

• •A catalyst lowers activation energy.
• •It provides an alternative reaction pathway.
• •It is regenerated overall.
• •It speeds the reaction without changing equilibrium.
• •Enzymes are catalysts used by living cells.

Hydrogen peroxide gives a clear demonstration. On its own it slowly breaks down into water and oxygen. Add a little black manganese dioxide powder and the solution fizzes vigorously as oxygen is released, which you can confirm because the gas relights a glowing splint. If you then filter out the powder, dry it and weigh it, you find almost the same mass you started with. A slice of raw potato or liver does a similar job, because it contains the enzyme catalase, which cells use to break down hydrogen peroxide safely.

Chemists sort catalysts into two broad kinds. A homogeneous catalyst is in the same phase as the reactants, such as a dissolved substance in a solution. A heterogeneous catalyst is in a different phase, usually a solid surface that gas or liquid molecules stick to while they react. Solid catalysts can be poisoned when other substances bind strongly to their surface and block it. This is why vehicles fitted with catalytic converters must run on unleaded petrol, since lead compounds coat the metal catalyst and stop it working.

Nitrogen makes up most of the air, but plants cannot use nitrogen gas directly, and the two atoms in each molecule are held by a very strong triple bond. In the early twentieth century, Fritz Haber showed that nitrogen and hydrogen could be combined to make ammonia, and Carl Bosch scaled the process up for industry. Making it practical depended on finding a cheap, reliable catalyst. Chemists at the company BASF tested thousands of materials before settling on iron-based catalysts, versions of which are still used today.

The Haber process shows the trade-offs catalysts help to manage. The reaction producing ammonia releases heat, so lower temperatures would favour a higher yield at equilibrium, but they would also make the reaction far too slow. Plants therefore use a compromise temperature of a few hundred degrees Celsius with high pressure, and the iron catalyst lets the reaction reach a useful rate under those conditions. Ammonia made this way is turned into fertilisers, and much of the nitrogen in the food people eat today has passed through an industrial catalyst.

It is easy to say that a catalyst lowers the activation energy, but the size of the effect can be startling. The fraction of collisions with enough energy depends exponentially on the height of the barrier. Near room temperature, a handy rule that follows from this is that every reduction of about 5.7 kilojoules per mole multiplies the rate by roughly ten, if nothing else changes. So a catalyst that lowers the barrier by 30 kilojoules per mole, which is not unusual, makes the reaction run close to two hundred thousand times faster at the same temperature.

Compare that with heating. A rough rule of thumb is that many reactions near room temperature go about twice as fast for every 10 degrees Celsius rise. Matching a two hundred thousandfold speed-up by heating alone would need around seventeen or eighteen doublings, which means warming the mixture by well over 150 degrees. Living cells cannot do that, since their proteins would be wrecked long before. Enzymes solve the problem by lowering barriers instead, often speeding reactions by factors of millions or more, which is how your body runs complex chemistry at a steady temperature of about 37 degrees Celsius.

Not every catalyst is welcome. High in the stratosphere, ozone absorbs much of the Sun's harmful ultraviolet light. In 1974 the chemists Mario Molina and Sherwood Rowland warned that chlorofluorocarbons, or CFCs, used in aerosol cans and refrigerators, could drift up there and be broken apart by ultraviolet light, releasing chlorine atoms. A chlorine atom reacts with ozone to form chlorine monoxide and oxygen. The chlorine monoxide then reacts with a free oxygen atom, releasing more oxygen and handing the chlorine atom back. The chlorine is regenerated, exactly like any other catalyst, ready to destroy another ozone molecule.

Because each chlorine atom goes round this cycle many times before being locked away in a less reactive compound, a small amount of CFC can destroy a great deal of ozone. The effect is strongest over Antarctica in spring, because icy clouds in the polar stratosphere provide surfaces on which stored chlorine is released in active form, a natural case of heterogeneous catalysis. A British Antarctic Survey team reported the ozone hole in 1985. The Montreal Protocol of 1987 led to CFCs being phased out, and Molina and Rowland shared the 1995 Nobel Prize in Chemistry with Paul Crutzen. Because CFCs last decades in the air, recovery is slow.

Catalysts make reactions faster by offering a lower-energy route between reactants and products. They participate in intermediate steps but are restored overall. Focus on activation energy, and the role of a catalyst becomes far clearer than the vague idea that it simply makes chemistry hurry up.