Unlocking the Power of Catalysts in Chemical Reactions

Dive into the fascinating world of catalysts and discover how these unsung heroes of chemistry accelerate reactions without being consumed. Learn their impact on industrial processes and biological functions.

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.

• We saw in Factors affecting rate of reaction that a catalyst speeds up a chemical reaction by providing an alternative route of lower activation energy for a reaction. The catalyst is important to meet the activation energy, otherwise the reaction may never even begin.
• Most reactions of fuels burning (combustion) are very exothermic. Overall, they release a lot of heat energy because the products are so much lower energy than the reactants. If you think of a reaction as going up a hill and then back down the other side, they would be like climbing up a staircase to fall off of a building.
  But if the activation energy of the reaction isnt met, it does not matter; the reaction will not start. Thats why fuels need a source of ignition, like a spark, to begin burning
• Catalysts are incredibly useful for financial and environmental reasons when running industrial scale chemical reactions.
  With no catalyst to lower the activation energy, very harsh conditions like high temperature and pressure must provide this energy the reaction needs. Harsh conditions costs energy to make and maintain, and the longer you run a reaction, the longer you must hold these conditions. This costs companies a lot of money and the environment more of its resources!
• The Haber process (iron), hydrogenation of ethene (nickel) and production of synthesis gas all use catalysts. They are very important industrial processes and catalysts make them less environmentally damaging.

This lesson looks in more detail at how catalysts have their effect.


• The catalysts alternative route is the reaction going by a different reaction mechanism than it does when no catalyst is present. The catalyst might lead to a lower energy intermediate substance that the reactants cant make without it.
  This would be like trying to climb over a high wall with a chair/stool, or without one. Which would be easier, and which would take less energy out of you?
• The reaction mechanism is the sequence of steps in how the reactant bonds break and the product bonds form. Intermediate(s) and transition state(s) are temporary molecules that form for a short time during the reaction, as the bonds rearrange. More on this in Reaction mechanisms.
• The reaction mechanism depends on the reactants and the catalyst used, but generally catalysts provide other possible reaction mechanisms. Some will be useless, higher energy than the uncatalyzed mechanism, while other mechanisms are lower energy and allow more reactant particles to become products, increasing the reaction rate as a result.
• Think about this with the chair/stool analogy: when a stool is on its side on the floor or upside down, its useless. When it is upright, its a lot more useful. Catalysts are the same it just provides more options, only one needs to be lower energy than without it though.
• 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 where an intermediate is formed.


• Many industrial processes, such as the Haber process, have their reactants in the gas phase.
  While this means the reactant molecules are high energy, it also means that the chances of collision between particles is lower because gas particles are not very compact.
  To counteract this, most industrial processes use a solid phase catalyst for their gas phase reactants. This is called a heterogeneous catalyst, which is any catalyst in a different phase to the reactants.
• Using a solid phase catalyst provides more surface area for the reactant particles to adsorb onto a surface. This is an attractive interaction that temporarily holds particles in place on a surface; this holding increases the chance of collisions between reactant particles happening at the surface. The larger the surface area of a catalyst, the more adsorption can happen, so the chance of successful collisions between reactant particles increases.
• An example of this is a catalytic converter in a car, which uses a solid platinum or rhodium catalyst to convert toxic carbon monoxide and NO_x gases into harmless N₂ and CO₂ before the gases are released to the environment. It has a thin honeycomb-like structure, so we get to use the most surface area for the smallest possible quantity of the expensive catalyst.
• Sometimes, particles adsorb to a surface and do not come off the holding is permanent. Chemists call this change poisoning of the catalyst, because that site where those molecules are bound permanently cant be used for the reaction anymore. Eventually, catalysts need replacing as more and more of their surface gets poisoned and they become less and less effective.


• 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.