$A$ gas is said to behave like an ideal gas when the relation $PV/T = \text{constant}$. When do you expect a real gas to behave like an ideal gas?

The given graph represents the variation of $Z$ (compressibility factor = $\frac{PV}{nRT}$) versus $P$,for three real gases $A, B$ and $C$. Identify the only incorrect statement.

$A$ certain quantity of real gas occupies a volume of $0.15 \ dm^3$ at $100 \ atm$ and $500 \ K$ when its compressibility factor is $1.07$. Its volume at $300 \ atm$ and $300 \ K$ (when its compressibility factor is $1.4$) is $........ \times 10^{-4} \ dm^3$ (Nearest integer).

The temperature at which the second virial coefficient of a real gas is zero is called:

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,$PV = nRT$.

In Van der Waals equation of state for a non-ideal gas,the term that accounts for intermolecular forces is