For simplicity we assume that all gasses behave as ideal gasses when it comes to the gas laws (PV=nRT). This assumes that individual gas particles have no mass and no intermolecular forces exist between molecules. It also assumes that all of the atoms of a gas are moving randomly but with relative uniformity and speed and are constantly colliding with one another in perfect collisions. But real gasses do deviate from this ideal behavior, particularly at higher temperatures and pressures. One feature of a non ideal gas is that the gas atoms do have an actual measurable mass that does affect the volume. Another feature is that the atoms of a gas do have interactive forces (Van der Waals forces) between then that do affect the pressure. All of this leads to the Van der Waals equation for non ideal gasses:
[P + (an^2/V^2)](V-nb) = nRT
where a is factor for the attractive forces between molecules and b is the factor for the volume of the individual gas molecules.