Catalysts - Energetics, Kinetics and Reaction Rate




In this lesson, we will learn:

  • To recall some important examples of catalysts in chemical industry.
  • To explain the specific effect of catalysts on reaction mechanisms and potential energy.
  • How to use potential energy curves to show catalyzed reaction pathways.


  • As defined in earlier lessons, a catalyst speeds up a chemical reaction by providing an "alternative route" of lower activation energy for reactants, without being consumed in the process. This lesson will look at the details of these effects that catalysts have – how do catalysts produce the effect they have?
    • A reaction might be highly exothermic, suggesting the products are more stable than the reactants, but this doesn't matter if the activation energy barrier of a reaction isn't met!
    • Catalysts in industrial processes are incredibly useful. Having a suitable catalyst can make a big difference in the efficiency and productivity of a process. If no catalyst is present to lower the activation energy barrier, then harsh conditions (very high temperature and pressure) have to provide that necessary energy. Harsh conditions costs energy to make and maintain, and the longer you run a reaction for, the longer you have to hold these conditions. This uses a lot of resources.
    • The Haber process (iron), hydrogenation of ethene (nickel) and production of synthesis gas all use catalysts. They are very important industrial processes made less energy intensive by catalysts.
  • This "alternative route" is the reaction going by a different reaction mechanism than it does when no catalyst is present. The catalyst might allow the formation of an intermediate or a different, more stable transition state than the reactants could make without a catalyst.
    • The reaction mechanism is how the products form; the precise sequence of how the reactant bonds break and the product bonds form (See C12.1.8: Reaction Mechanisms for more).
    • Intermediates and transition states are molecules that form for a short time during the reaction that can affect the activation energy. More on these later!
    • Reaction mechanisms will vary depending on the reactants and the catalyst involved, generally catalysts provide other possible mechanisms for the reactants to become products – some may be higher energy than the uncatalyzed mechanism (and obviously aren't useful!), while other mechanisms are lower energy and allow more reactant particles to become products, increasing the reaction rate as a result.
    • The whole description of the reaction procedure, including potential energy involved and the stages in the mechanism is called the reaction pathway. These are often described using a potential energy diagram.
    • A typical reaction with a catalyzed and uncatalyzed pathway will appear different on potential energy curves. The catalyzed pathway will have a lower activation energy and may have a "dip" at the crest of the peak – this is due to an intermediate being formed (See C12.1.8: Reaction Mechanisms). See below for examples:
  • When dealing with reactions at equilibrium, remember that activation energy barriers and reaction mechanisms exist and work for reverse reactions too! A catalyst will speed up forward and reverse reactions in a chemical process; equilibrium position is unchanged. Only the time taken to reach equilibrium is reduced.
  • Intro Lesson
    The effect of catalysts
  • 1.
    Explain the use of catalyst and their effect on the energy profile of a reaction.
    The decomposition of hydrogen peroxide is shown in the equation below:

    2H2O22H2O+O2\mathrm{2H_2O_2 \to 2 H_2O + O_2}

  • 2.
    Explain the use of catalyst and their effect on the energy profile of a reaction. The Haber process is an important industrial process to produce ammonia. The equation is below:

    3H2+N22NH3\mathrm{3H_2 + N_2 \rightleftharpoons 2NH_3 }

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