Explain why a catalyst speeds up a chemical reaction.
A catalyst lowers the activation energy of a reaction to speed it up
I couldn't have said it better.
A catalyst speeds up a chemical reaction by providing an alternative pathway with lower activation energy for the reaction to proceed.
To understand why a catalyst is effective in increasing the reaction rate, it is important to know the concept of activation energy. Activation energy is the minimum energy required for a reaction to occur. In a chemical reaction, reactant molecules must collide with enough energy to break the existing bonds and form new ones to convert into products.
A catalyst works by lowering the activation energy of a reaction. It does so by providing an alternative pathway, called the reaction pathway, that has a lower energy barrier compared to the original pathway without the catalyst. The catalyst achieves this by interacting with the reactant molecules, which leads to the formation of an intermediate species or transition state. This intermediate species is short-lived and subsequently converts into the final products.
The interaction of the reactant molecules with the catalyst lowers the energy needed for bond-breaking and bond-forming steps. This reduced activation energy barrier allows more reactant molecules to possess enough energy to surpass the barrier and convert into products. Consequently, a larger fraction of reactant molecules can successfully undergo the reaction pathway and contribute to the overall reaction rate.
It is essential to note that a catalyst itself does not undergo permanent changes during the reaction. It is not consumed in the process and remains unchanged at the end of the reaction, enabling it to participate in multiple reaction cycles. Therefore, even a small amount of catalyst can have a significant impact on the reaction rate by increasing the rate of reactant conversion into product formation.
In summary, a catalyst speeds up a chemical reaction by providing an alternative reaction pathway with a lower activation energy barrier. By lowering the energy requirement for the reaction, more reactant molecules can overcome the energy barrier, leading to enhanced reaction rates.