Question

How does an increase in pressure affect the rate of a gasphase reaction? Explain.

Short answer

An increase in pressure increases the rate of a gas-phase reaction by increasing the frequency of collisions between reactant molecules.

Step-by-step solution

- Understand the Kinetic Molecular Theory

The Kinetic Molecular Theory states that gas particles are in constant, random motion and that the rate of a reaction depends on the frequency and energy of collisions between reactant molecules.

- Define Pressure

Pressure in a gas is defined as the force exerted by gas molecules colliding with the walls of their container per unit area. An increase in pressure means more gas molecules are present in the same volume.

- Analyze Collision Frequency

When pressure increases, the number of gas molecules per unit volume increases. This results in more frequent collisions between reactant molecules.

- Relate to Reaction Rate

More frequent collisions generally mean that more reactant molecules have the opportunity to collide with enough energy to overcome the activation energy barrier, leading to an increase in the reaction rate.

Key concepts

Kinetic Molecular Theory

The Kinetic Molecular Theory provides the foundational understanding of how gas particles behave.
According to this theory, gas particles are constantly moving in random directions. They collide with each other and the walls of the container.
These collisions are essential as they influence various gas properties, including pressure and temperature. The theory posits that the rate of a gas-phase reaction depends on the frequency and energy of these collisions.
When particles have higher energy collisions, there's a greater chance they will overcome the activation energy barrier needed for a reaction to occur.
Thus, understanding the movement and interaction of gas particles helps connect how changes in conditions like pressure can impact the reaction rate.

Gas Pressure

Pressure in a gas is the force exerted by gas molecules hitting the walls of their container.
The more frequent these collisions, the higher the pressure. If the pressure in a gas phase reaction increases, it typically means that there are more gas molecules in a given volume.
Imagine squeezing a balloon – the pressure inside increases because you are forcing the gas molecules into a smaller space, resulting in more collisions with the walls.
Thus, increasing pressure effectively means that there are more reactant molecules available to collide and react with each other.

Collision Frequency

Collision frequency refers to how often gas molecules collide with each other.
When you increase the pressure, you pack more molecules into the same space, leading to more frequent collisions.
These collisions are not just random bumps; they are opportunities for reactant molecules to interact and possibly react.
For a reaction to happen, these molecules need to collide with enough energy and proper orientation.
Therefore, by increasing the collision frequency, the chances of effective collisions that can overcome the activation energy barrier are increased, thereby speeding up the reaction rate.

Reaction Rate

The reaction rate is how fast a reaction occurs.
A higher reaction rate means that products are formed faster.
More frequent and energetic collisions between reactant molecules can lead to a higher reaction rate.
When the pressure increases, the density of the gas increases, leading to more collisions per unit time.
This means more molecules have the energy to surpass the activation energy barrier, causing an increase in the rate at which the reaction proceeds.
In simple terms, an increase in pressure generally leads to an increase in the reaction rate because of the higher number of effective collisions.