Catalysts Definition and How They Work

A catalyst is a chemical substance that affects the rate of a chemical reaction by altering the activation energy required for the reaction to proceed. This process is called catalysis. A catalyst is not consumed by the reaction and it may participate in multiple reactions at a time. The only difference between a catalyzed reaction and an uncatalyzed reaction is that the activation energy is different. There is no effect on the energy of the reactants or the products. The ΔH for the reactions is the same.​

How Catalysts Work

Catalysts permit an alternate mechanism for the reactants to become products, with a lower activation energy and different transition state. A catalyst may allow a reaction to proceed at a lower temperature or increase the reaction rate or selectivity. Catalysts often react with reactants to form intermediates that eventually yield the same reaction products and regenerate the catalyst. Note that the catalyst may be consumed during one of the intermediate steps, but it will be created again before the reaction is completed.

Read More

Catalysis Definition in Chemistry

By Anne Marie Helmenstine, Ph.D.

Positive and Negative Catalysts (Inhibitors)

Usually when someone refers to a catalyst, they mean a positive catalyst, which is a catalyst that speeds up the rate of a chemical reaction by lowering its activation energy. There are also negative catalysts or inhibitors, which slow the rate of a chemical reaction or make it less likely to occur.

Promoters and Catalytic Poisons

A promoter is a substance that increases the activity of a catalyst. A catalytic poison is a substance that inactivates a catalyst.

Catalysts in Action

• Enzymes are reaction-specific biological catalysts. They react with a substrate to form an unstable intermediate compound. For example, carbonic anhydrase catalyzes the reaction:
  H₂CO₃(aq) ⇆ H₂O(l) + CO₂(aq)
  The enzyme allows the reaction to reach equilibrium more quickly. In the case of this reaction, the enzyme makes it possible for carbon dioxide to diffuse out of blood and into the lungs so it can be exhaled.
• Potassium permanganate is a catalyst for the decomposition of hydrogen peroxide into oxygen gas and water. Adding potassium permanganate increases the temperature of the reaction and its rate.
• Several transition metals can act as catalysts. A good example of platinum in the catalytic converter of an automobile. The catalyst makes it possible to turn toxic carbon monoxide into less toxic carbon dioxide. This is an example of heterogeneous catalysis.
• A classic example of a reaction that doesn't proceed at an appreciable rate until a catalyst is added is that between hydrogen gas and oxygen gas. If you mix the two gases together, nothing much happens. However, if you add heat from a lighted match or a spark, you overcome the activation energy to get the reaction started. In this reaction, the two gases react to produce water (explosively).
  H₂ + O₂ ↔ H₂O
• The combustion reaction is similar. For example, when you burn a candle, you overcome the activation energy by applying heat. Once the reaction starts, heat released from the reaction overcomes the activation energy needed to allow it to proceed.