What we call "volume" of a gas is the sum of the "free" volume where the molecules are free to move, plus the volume of the molecules themselves. The collisions theory is only valid for a volume which is the free space available to the moving molecules. What we call "pressure" of a gas is the force exerted by the particules when touching the walls of the container. But this force in a compressed gas is not simply due to its kinetic energy due to the temperature. It is diminished if there is an attraction between the particules.

In a gas made of $\ce{H2}$, there is practically no attraction between individual molecules. The pressure is due to the collisions with the wall due to particules whose energy is simply proportional to $\pu{T}$. So the total volume is equal plus the value p = nRT/V, plus the volume of the molecules. It increases with P. So $\pu{Z = pV/RT}$ increases with $\pu{p}$.

In a gas like $\ce{CO2}$ or $\ce{CO}$, the pressure measured on the wall is smaller than the theoretical values coming from the kinetic theory of gases, because the measured pressure is slowed down by their mutual attraction. That is why $\pu{Z = pV/nRT}$ decreased by increasing the pressure

Let's start thinking qualitatively.

1. What we call "volume" of a gas is the sum of the "free" volume where the molecules are free to move, plus the volume of the molecules themselves. The collision theory is only valid for a volume which is the free space available to the moving molecules.

2. What we call "pressure" of a gas is the force exerted by the particules when touching the walls of the container. But this force in a compressed gas is not simply due to its kinetic energy due to the temperature. It is diminished if there is an attraction between the particules.

In a gas like $\ce{H2}$, there is practically no attraction between individual molecules. The pressure is due to the collisions with the walls due to particules whose energy is simply proportional to $\pu{T}$. So the total volume is equal to the theoretical value $\ce{V_o = nRT/V}$, plus the volume of the molecules. It increases with the pressure. So the measured $\pu{V}$ is greater than $\pu{nRT/p}$ and $\pu{Z = pV/nRT}$ increases with $\pu{p}$.

In a gas like $\ce{CO2}$ or $\ce{CO}$, the pressure measured on the wall is smaller than the theoretical values coming from the kinetic theory of gases, because the measured pressure is slowed down by their mutual attraction. That is why $\pu{Z = pV/nRT}$ decreased by increasing the pressure

What we call "volume" of a gas is the sum of the "free" volume where the molecules are free to move, plus the volume of the molecules themselves. The collisions theory is only valid for a volume which is the free space available to the moving molecules. What we call "pressure" of a gas is the force exerted by the particules when touching the walls of the container. But this force in a compressed gas is not simply due to its kinetic energy due to the temperature. It is diminished if there is an attraction between the particules.

In a gas made of $\ce{H2}$, there is practically no attraction between individual molecules. The pressure is due to the collision with the wall due to particules whose energy is simply proportional to $\pu{T}$. So the total volume is equal plus the value p = nRT/V, plus the volume of the molecules. It increases with P. So pV/RT increases with P.

In a gas like $\ce{CO2}$ or $\ce{CO}$, the pressure measured on the wall is smaller than the values coming from the kinetic theory of gases, because the measured pressure is slowed down by their mutual attraction. It looks as if there was less particles having the theoretical energy proportional to T. That is why p = nRT/V decreased by increasing the pressure

What we call "volume" of a gas is the sum of the "free" volume where the molecules are free to move, plus the volume of the molecules themselves. The collisions theory is only valid for a volume which is the free space available to the moving molecules. What we call "pressure" of a gas is the force exerted by the particules when touching the walls of the container. But this force in a compressed gas is not simply due to its kinetic energy due to the temperature. It is diminished if there is an attraction between the particules.

In a gas made of $\ce{H2}$, there is practically no attraction between individual molecules. The pressure is due to the collisions with the wall due to particules whose energy is simply proportional to $\pu{T}$. So the total volume is equal plus the value p = nRT/V, plus the volume of the molecules. It increases with P. So $\pu{Z = pV/RT}$ increases with $\pu{p}$.

In a gas like $\ce{CO2}$ or $\ce{CO}$, the pressure measured on the wall is smaller than the theoretical values coming from the kinetic theory of gases, because the measured pressure is slowed down by their mutual attraction. That is why $\pu{Z = pV/nRT}$ decreased by increasing the pressure