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The catalyst lowers the activation energy by providing the reaction an alternative energy pathway. With a catalyst, more collisions result in a reaction because the reactants need less energy to get to the transition state (with the highest energy), so the rate of reaction increases.
Hope it helps!
Hope it helps!
The concept part of it would be that there is an inverse relationship between the activation energy and the rate of a reaction: the lower the activation energy, the higher the rate of the reaction. A catalyst offers an alternative pathway with a lower activation energy, therefore, the reaction goes faster.
The activation energy determines the rate of reaction. To increase the rate of reaction, the activation energy or energy barrier must be lowered. For example, when rolling a ball up a hill and then letting it roll down, the get up the hill is the activation energy. A catalyst just builds a tunnel through the hill to get to the other side.
Since the rate of reaction and also activation energy depends on successful collision of molecules in the correct orientation, we can think of catalysts as things that helps the reactants collide more frequently in the correct orientation. Some catalyst such as many metals actually allow reactants to adhere to itself. So the reaction takes place on the surface of the catalyst. If you think about it this way, it will be easier for a reactant to collide with another reactant that is already adhere to the surface of the catalyst, thus the catalyst will lower the activation energy and increase the rate of reaction.
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