Homogeneous and Heterogeneous Catalysis

A catalyst is a substance that speeds up a chemical reaction by providing an alternative pathway with a lower activation energy, without being used up; in homogeneous catalysis the catalyst is in the same phase as the reactants, and in heterogeneous catalysis it is in a different phase. In brief: lowering the activation energy from 75 to 50 kJ/mol at 298 K speeds up a reaction by a factor of about 24,000, but a catalyst does not change ΔH or the equilibrium constant.

Catalysts are essential to modern life. They are used in the production of fertilizers, fuels and plastics, they clean the exhaust of cars, and enzymes catalyze practically every reaction in living cells. This article explains how catalysts work, compares the two main types, gives examples and calculations and describes what a catalyst can and cannot do.

How does a catalyst work?

A catalyst takes part in the reaction mechanism and creates a new route, often with several steps, in which each barrier is lower than the barrier of the uncatalyzed reaction. It is used in one step and regenerated in a later one, so a small amount can convert a very large amount of reactant. The number of reactant molecules converted per catalyst site per second is called the turnover frequency.

What can a catalyst not do?

• It does not change ΔH, ΔG or the position of the equilibrium, so the equilibrium constant K stays the same.
• It speeds up the forward and the reverse reactions by the same factor, so the equilibrium is reached faster.
• It cannot make a thermodynamically non-spontaneous reaction go (ΔG > 0).
• It is not consumed in the net reaction, although it can gradually be deactivated, for example by poisoning.

What is homogeneous catalysis?

The catalyst and the reactants are in the same phase, usually a solution or a gas mixture. The catalyst mixes with the reactants on the molecular level, so its action is easy to study, but it is often difficult to separate from the products. Examples:

• Acid catalysis: H⁺ speeds up the hydrolysis of esters and the formation of esters.
• Iodide ions: I⁻ speeds up the decomposition of H₂O₂.
• Ozone destruction: chlorine atoms from chlorofluorocarbons catalyze the breakdown of ozone in the stratosphere: Cl + O₃ → ClO + O₂, then ClO + O → Cl + O₂. The net reaction is O₃ + O → 2O₂, and the chlorine atom is regenerated, so that a single atom can destroy of the order of 10⁵ ozone molecules before it is removed.
• Transition-metal complexes: used in industrial processes, for example the hydroformylation of alkenes and the Wacker process.

What is heterogeneous catalysis?

The catalyst is a solid, and the reactants are gases or liquids. The reaction takes place on the surface of the catalyst. The steps are: (1) adsorption of the reactants on the surface, (2) weakening of their bonds and reaction on the surface, (3) desorption of the products. Because the catalyst is easily separated from the products, this type is preferred in industry. The surface area is important, so catalysts are used as fine powders or on porous supports.

Surface reactions can be blocked by catalyst poisons, substances that bind strongly to the active sites. Lead compounds poison the platinum in a catalytic converter, which is why leaded gasoline cannot be used in cars fitted with them. Sulfur compounds also poison many metal catalysts.

Examples of catalyzed processes

What about enzymes?

Enzymes are biological catalysts, mostly proteins, that are extremely efficient and selective. Many enzymes increase the rate of reaction by factors of 10⁶ to 10¹⁷. The enzyme catalase, which decomposes hydrogen peroxide in cells, can convert millions of molecules per second per enzyme molecule. Enzymes work under mild conditions of temperature and pH and are selective for a specific substrate. Because they are proteins, they are denatured by too much heat. Enzyme catalysis can be seen as intermediate between homogeneous and heterogeneous catalysis.

Worked examples

Example 1: the effect of lowering the activation energy

A catalyst lowers the activation energy from 75 kJ/mol to 50 kJ/mol. Assuming that the factor A is unchanged, the ratio of the rate constants at 298 K is k_cat/k_uncat = e^ΔE_a/RT = e^{25,000/(8.314 × 298)} = e^10.09 = 2.4 × 10⁴.

Example 2: catalyst and equilibrium

In the Haber process, the equilibrium constant at a given temperature is the same with or without the iron catalyst. The catalyst makes it possible to reach equilibrium within seconds at 450 °C and does not change the final yield, which is determined by the temperature and pressure.

Example 3: identifying the catalyst in a mechanism

The mechanism of ozone breakdown is Cl + O₃ → ClO + O₂ followed by ClO + O → Cl + O₂. Cl is used in step 1 and regenerated in step 2, so it is the catalyst. ClO is formed in step 1 and used in step 2, so it is an intermediate.

How were catalysts discovered?

Jöns Jacob Berzelius introduced the term catalysis in 1835 to describe reactions that are speeded up by substances that are not consumed. Wilhelm Ostwald studied catalysis as a change in rate, and received the Nobel Prize in Chemistry in 1909. Paul Sabatier received the prize in 1912 for the hydrogenation of organic compounds with finely divided metals. Fritz Haber and Carl Bosch developed ammonia synthesis with an iron catalyst (1909–1913), for which Haber (1918) and Bosch (1931) received Nobel Prizes. Gerhard Ertl was awarded the 2007 prize for his studies of reactions on solid surfaces.

Which mistakes are common?

• Saying that a catalyst is not involved in the reaction: it takes part but is regenerated.
• Believing that a catalyst increases the yield: it only changes the rate, not the equilibrium.
• Confusing a catalyst with a reactant: the catalyst is not in the overall equation.
• Thinking that it changes ΔH: the enthalpy change is the same.
• Forgetting that catalysts can be poisoned: impurities can reduce their activity.