Why would decreasing the volume of a container in which a gas-phase reaction is taking place speed up the reaction?

In a gas-phase reaction, reactant molecules are in constant motion, colliding with one another as they move throughout the container. When two reactant molecules collide with the correct orientation and sufficient energy, they can undergo a reaction to form one or more product molecules. This process continues until the reactants have been consumed or the reaction has reached an equilibrium state.

The volume of the container affects both the concentration of the reactant molecules and the frequency of collisions. Decreasing the volume of a container forces the reactant molecules closer together, effectively increasing their concentration. Gas laws, such as the ideal gas law \[PV = nRT\], relate the pressure, volume, and temperature of a gas. As per this law, where P is pressure, V is volume, n is the number of moles of gas, R is the gas constant, and T is the temperature, when volume decreases while keeping temperature constant, the pressure increases. Gases have a tendency to spread out evenly in any container, so decreasing the volume essentially increases the number of gas molecules per unit volume. This leads to a higher concentration of reactant molecules in the available space, which, in turn, leads to more frequent molecule collisions.

The reaction rate is directly affected by the concentration of reactant molecules, as described by the rate law for a particular reaction. In general, when the concentration of reactant molecules increases, the frequency of successful collisions between molecules also increases. This results in a higher reaction rate, as more reactions take place in a given period.

In conclusion, decreasing the volume of a container in which a gas-phase reaction is happening speeds up the reaction by increasing the concentration of reactant molecules. The increased concentration results in more frequent collisions between reactant molecules, which increases the reaction rate. This happens as long as the increased pressure does not significantly alter the reaction conditions (e.g., not causing phase changes or affecting the equilibrium state).