Catalysts — Chemistry, 14–17
A catalyst gives a reaction another route with a lower activation energy. It makes the reaction faster without being used up, but it does not change the final equilibrium position.
A shortcut over an energy hill
Reacting particles must usually overcome an energy barrier before products can form. A catalyst provides a different reaction pathway with a smaller barrier. More particles can cross it at the same temperature, so the reaction happens faster.
Why were catalysts needed?
Many useful reactions are possible but painfully slow at a workable temperature. Industry needed to make ammonia, fuels and plastics quickly without simply heating everything to dangerous temperatures. Catalysts solve that practical problem by changing the route, not by supplying the reaction’s energy.
Manganese dioxide and hydrogen peroxide
Hydrogen peroxide slowly breaks down: 2H₂O₂ → 2H₂O + O₂. Add a small amount of manganese dioxide, MnO₂, and bubbles of oxygen appear much more quickly. Filter the mixture and the MnO₂ remains, showing that it helped the reaction but was not used up overall.
Faster does not mean more
Because a catalyst makes more product appear quickly, it is easy to think it increases the final amount. It does not create extra particles or change the energy difference between reactants and products. It speeds both directions of a reversible reaction, so equilibrium is reached sooner but stays in the same place.
Catalysts in a car exhaust
A car’s catalytic converter contains metals that help convert carbon monoxide and unburned hydrocarbons into less harmful substances. It also helps turn nitrogen oxides into nitrogen. The metals are not the fuel and are not meant to be used up; they provide surfaces where reactions can happen more easily.
Keep exploring
Other languages
Loading MyLeoNes™…