pV will be constant (Boyle’s law) and pV vs p graph at all pressures will be a straight line parallel to x-axis.

At constant temperature pV vs p plot for real gases is not a straight line.

pV vs p plot two types of curves are seen.

In the curves of pV vs P plot for dihydrogen and helium, as the pressure increases the value of pV also increases.

The second type of plot is seen in the case of other gases like carbon monoxide and methane.

In pv vs p plots there is a negative deviation from ideal behaviour.

The pV value decreases with increase in pressure and reaches to a minimum value characteristic of a gas.

After that pV value starts increasing then the curve crosses the line for ideal gas and after that shows positive deviation continuously.

real gases do not follow ideal gas equation perfectly under all conditions.

Real gases do not follow, Boyle’s law, Charles law and Avogadro law perfectly under all conditions.

Two assumptions of the kinetic theory do not hold good.

The assumptions of the kinetic theory do not hold good are there is no force of attraction between the molecules of a gas and volume of the molecules of a gas is negligibly small in comparison to the space occupied by the gas.

If assumption there is no force of attraction between the molecules of a gas is correct, gas will never liquify.

Gases do liquify when cooled and compressed, and liquids formed are difficult to compress.

forces of repulsion prevent molecule squashing in tiny volumes proves that liquids formed are difficult to compress.

The pressure vs volume graph of experimental data and theoretically calculated from Boyles law should coincide.

Real gases show deviations from the ideal gas law due to molecular interactions.

At high pressures, molecules are close to each other, causing molecular attractive forces that reduce the pressure exerted by molecules on the container walls, resulting in lower pressure than ideal gas.

Van der Waals constants a is a measure of intermolecular attraction within a gas, independent of temperature and pressure.

At very low temperatures, intermolecular forces become significant, allowing molecules to capture each other.

Real gases exhibit ideal behavior when temperature and pressure are negligible, and ideal behavior when pressure approaches zero.

The deviation from ideal behaviour can be measured in terms of compressibility factor Z, which is the ratio of product pV and nRT.

Ideal gas Z = 1 at all temperatures and pressures, as pV = n RT.

The graph of Z vs p is a straight line parallel to the pressure axis.

Gases that deviate from ideality deviate from unity.

At very low pressures, all gases have Z ≈1 and behave as ideal gas.

At high pressures, all gases have Z > 1 and are more difficult to compress.

At intermediate pressures, most gases have Z < 1.

Gases show ideal behavior when the volume occupied is large, allowing the volume of molecules to be neglected.

The Boyle temperature or Boyle point determines the temperature at which a real gas obeys the ideal gas law over a range of pressure.

Thomas Andrews obtained the first complete data on pressure-volume-temperature relations of a substance in both gaseous and liquid states.

Isotherms of carbon dioxide at various temperatures and found that real gases behave similarly to carbon dioxide.

At high temperatures, isotherms resemble ideal gas behavior, and the gas cannot be liquified even at high pressure.   

As temperature decreased, the shape of the curve changed, and data showed deviation from ideal behavior.

The critical temperature (TC) of carbon dioxide is 30.98°C.

The critical temperature (TC), critical volume (VC) and pressure (pC) are critical constants.