Question

The number of collisions of Ar atoms with the walls of containers per unit time?
A.Increases when the temperature increases
B.Remains the same when $C{O_2}$ is added to the container at a constant temperature
C.Increases when $C{O_2}$ is added to the container at a constant temperature
D.Decreases, when the average kinetic energy per molecule is decreased

Answer 1

Argon is a real gas. Vander Waal equation is a thermodynamic equation of state based on the theory that liquids are composed of particles with non-zero volumes but have interparticle attractive force. It is for real gases. The pressure of a gas in a box is the outwards force on the walls of the box.

Complete step by step answer:
We know that the rate of a chemical reaction is directly proportional to the number of collisions per second. This is because proper collision leads to a chemical reaction. The reaction rate increases when more reactant molecules collide with one another. Therefore, according to collision theory, only sufficient energetic molecules have enough energy to react.
The number of effective collisions per unit time is, $z = \dfrac{1}{4}N\;{u_{avg}}$ ,
Where, ${u_{avg}}$ are the equals to $\sqrt {\dfrac{{8RT}}{M}}$ . And N equals to, $\dfrac{{PV}}{{RT}}$ .
Therefore,

$$z = \dfrac{1}{4}N\;{u_{avg}} \\\ z = \dfrac{1}{4}\dfrac{{PV}}{{RT}}\;\sqrt {\dfrac{{8RT}}{M}} \\\ z\infty \dfrac{1}{{\sqrt T }} \\\$$

So, the value of Z will be the same if the value of temperature remains constant.
Therefore, the number of collisions of Ar atoms with the walls of container per unit time remains the same when $C{O_2}$ is added to the container at a constant temperature
So, the correct answer is B.

Note: According to the collision theory, the reactants must collide with sufficient energy so that chemical bonds between the atoms can easily break. This sufficient energy is known as the activation energy. The molecules must collide with a proper orientation for a chemical reaction to occur. As the temperature increases, molecules tend to move faster and thus collide more vigorously. This increases the chance of bond breakage upon collision.