Why do real gases deviate from ideal behavior in the $pV$ vs $p$ and $p$ vs $V$ graphs at specific temperatures?

$CH_4$ gas is behaving non-ideally. The compressibility factor for the gas is $1.5$ at $2 \ atm$ and $400 \ K$. Calculate the molar volume for the gas in $litre$. [Given $R = 0.08 \ L \ atm \ K^{-1} \ mol^{-1}$]

Which among the following statements is/are incorrect regarding real gases?
$(i)$ Their compressibility factor is never equal to unity $(Z \neq 1)$.
$(ii)$ The deviations from ideal behavior are less at low pressures and high temperatures.
$(iii)$ Intermolecular forces among gas molecules are equal to zero.
$(iv)$ They obey Van der Waals equation,$PV = nRT$.

The pressure $(p)$ and volume $(V)$ isotherm of a van der Waals' gas,at the temperature at which it undergoes gas to liquid transition,is correctly represented by

The correct increasing order of the value of van der Waals's constant $a$ for $CH_4, O_2$ and $H_2$ gases is

Which intermolecular force is most responsible for the liquefaction of inert gases?