When gases are treated as real, via use of the van der waals equation, the actual volume occupied by gas molecules the pressure exerted and the attractive forces between gas molecules the pressure exerted, as compared to an ideal gas.

Increases, decreases.

Explanation:

At room temperature and moderately low pressures ideal gases, among other assumptions, :

I.Do not attract or repel each other.

II.Have negligible volume relative to the volume of the containing vessel.

Thus, the ideal gas equation:

PV = nRT

Where P = gas pressure = gas molecules' collision with the walls of the containing vessel.

V = volume occupied by the gas sample.

T = Kelvin temperature of the gas

n = number of moles of the gas present.

R = proportionality constant called the molar gas constant = 0.0821 L. atm/K.mol.

However, at low temperatures and high pressures, gases exhibit real behaviors and the

I.attractive forces between the gas molecules may not be negligible.

II.volume of the gas molecules may not be negligible relative to the volume of the container.

At low temperatures, the average kinetic energy of the gas molecules decreases which prevents the molecules breaking free from their molecular attraction.

At high pressures, the density of the gas increases and the molecules come closer to one another. This increases the intermolecular forces between the gas molecules, increases the number of molecules found per unit volume and lowers their speed moving towards the containing wall thus lowering the pressure the gas would exert than if it were in an ideal setting.

van der Waals corrected the actual pressure and volume of real gases as follows:

( P + an^2/V^2 ) ( V – nb ) = nRT

Where n^2/V^2 = number of molecules per unit volume.

P = observed pressure.

( V – nb ) = effective volume of the gas.

nb = volume occupied by n moles of the gas.

a and b = proportionality constants which are different for every gas.

When gases are treated as real, via use of the van der Waals equation, the actual volume occupied by gas molecules increases the pressure exerted and the attractive forces between gas molecules decreases the pressure exerted, as compared to an ideal gas.

Explanation:

The van der Waals equation considers two aspects of gases which are not considered for ideal gases:

- Gas molecules occupy a volume. The actual volume is larger than the calculated volume for an ideal gas, so it exerts a higher pressure. This aspect is corrected by the parameter b in van der Waals equation.

- Gas molecules interacts each other, specially by attractive forces when the gas is at a higher pressure (remember that an ideal gas is good model for low pressures). When they attract each other, the molecules exert a lesser pressure on the container walls, so the pressure is lower than the exerted by an ideal gas. This aspect is corrected by the parameter a in van der Waals equation.

more infor would

a lot