The temperature of a gas is raised while its volume remains constant, the pressure exerted by a gas on the walls of the container increases because its molecules

The correct option is: Strike the wall more often with higher velocities.

To understand why the pressure increases, we can analyze the behavior of gas molecules using the kinetic molecular theory of gases. According to this theory, the pressure exerted by a gas is a result of the continuous collisions of its molecules with the walls of the container.

The average kinetic energy of the gas molecules is directly proportional to the absolute temperature of the gas, which is expressed as:

$\mathrm{K.E.}=\frac{3}{2}{k}_{B}T$

where $k_B$ is the Boltzmann constant and $T$ is the absolute temperature.

When the temperature of the gas is raised:

1. Higher Velocities: The average kinetic energy of the molecules increases, meaning they move with higher average speeds (velocities). As a result, when they collide with the container walls, they do so with greater momentum, exerting a larger force per collision.
2. More Frequent Collisions: Since the volume of the container remains constant, the molecules have the same space to move around but are traveling much faster. Consequently, the time interval between successive collisions with the walls decreases, causing the molecules to strike the walls more frequently.

Since pressure is defined as the force exerted per unit area, and force is equal to the rate of change of momentum, both the increased speed of the molecules and the increased frequency of collisions contribute to a higher rate of momentum transfer to the walls. Therefore, the pressure exerted by the gas increases.