The rate of a gaseous reaction generally increases when the pressure is increased.
Here's a breakdown of why this happens:
The Collision Theory
The foundation for understanding this relationship lies in the collision theory of chemical kinetics. This theory states:
- For a reaction to occur, reactant molecules must collide.
- These collisions must have sufficient energy (activation energy).
- The molecules must collide with the correct orientation.
Pressure and Concentration
Increasing the pressure of a gaseous system effectively increases the concentration of the gaseous reactants. Think of it like squeezing the same amount of gas into a smaller volume.
Increased Collision Frequency
With a higher concentration of reactants, the frequency of collisions between reactant molecules increases. Because there are more molecules in a given space, they are more likely to bump into each other.
Increased Reaction Rate
Since the frequency of collisions increases, the number of effective collisions (those with sufficient energy and proper orientation) also increases. This directly leads to an increased rate of reaction. The more effective collisions there are per unit time, the faster the reaction proceeds.
Summary
In essence, increasing the pressure in a gaseous reaction increases the concentration of the reactants, which leads to more frequent and effective collisions, thereby speeding up the reaction.
