Answer:
A. Some real gases have lower pressures than that calculated for the ideal gas because of attraction between molecules.
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 mole 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 behavior.
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.
