Consider a gas-phase reaction. (a) Cooling the mixture of reactants can slow and even stop the chemical reaction. Explain why this is so. (b) Increasing the number of reactant molecules in the flask makes the reaction go faster. Explain why this is so.

: To explain the effect of cooling on the reaction, we need to consider the activation energy and the distribution of molecular energies in the gas phase. Activation energy is the minimum amount of energy required for a reaction to occur between specific reactants. In a gas phase reaction, reactant particles collide with each other, and if the collision has enough energy (above the activation energy), the reaction takes place. Otherwise, the molecules simply collide and bounce off one another without any reaction occurring. When a mixture of reactant molecules is cooled, reduced kinetic energy results in lower energy collisions between reactant particles. Consequently, fewer collisions possess the necessary energy (activation energy) to initiate a chemical reaction. As the temperature decreases, the number of effective collisions also decreases, ultimately leading to a slower reaction rate or even a complete halt in the reaction process.

: To explain the effect of increasing the number of reactant molecules in the flask, we need to consider the collision theory of chemical reactions. Collision theory states that the number of successful molecular collisions directly affects the rate of a chemical reaction. More specifically, the more reactant molecules present in a system, the higher the likelihood of successful collisions, and thus the faster the reaction. By increasing the number of reactant molecules in the flask, we are indirectly enhancing the frequency of favorable collisions involving the reactants, leading to a faster reaction rate. This happens because more reactant particles are present in proximity, increasing the chance of successful collisions and the formation of product molecules.