Which of the following is the BEST explanation for why a catalyst speeds up a reaction?
A) The catalyst increases the concentration of the products
B) The catalyst decreases the total pressure in the container.
C) The catalyst raises the temperature of the reaction mixture.
D) The catalyst lowers the activation energy required for the reaction to occur.
E) The catalyst increases the probability that a collision will occur by taking up space in the container.
The BEST explanation for why a catalyst speeds up a reaction is option D) The catalyst lowers the activation energy required for the reaction to occur.
To understand why a catalyst lowers the activation energy, let's first explain what activation energy is. Activation energy is the minimum amount of energy required for a chemical reaction to take place. In a reaction, the reactant molecules need to collide with enough energy to overcome the energy barrier or activation energy in order to form products.
A catalyst is a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy. It does not get consumed during the reaction and can be used repeatedly. When a catalyst is present, it provides a different pathway for the reaction to follow. This alternative pathway has a lower energy barrier, making it easier for the reactant molecules to overcome and proceed to form product molecules.
By lowering the activation energy, the catalyst increases the rate of the reaction without affecting the concentration of products (option A), the total pressure in the container (option B), or the temperature of the reaction mixture (option C). Option E, which states that the catalyst increases the probability of a collision by taking up space in the container, is not a correct explanation in terms of the role of a catalyst in lowering the activation energy.