"Real gases do not follow Boyle's Law,Charles's Law,and Avogadro's Law perfectly under all conditions." Explain.

(A) Explanation using $pV$ vs $p$ graph according to Boyle's Law: Theoretically,$pV = nRT$. If we plot a graph of $pV$ vs $p$ at a constant temperature $(T)$,we should get a straight line parallel to the $X$-axis because,according to Boyle's Law,$pV$ is constant for a given amount of ideal gas.
In reality,the $pV$ vs $p$ graph is not a straight line. Experimental data for real gases at constant temperature shows significant deviation from ideal behavior.
Graph Type-$1$: For gases like $H_2$ and $He$,the $pV$ value increases continuously with an increase in pressure $(p)$.
Graph Type-$2$: For real gases like $CO$ and $CH_4$,the curve first shows a negative deviation from ideal behavior,meaning $pV$ values decrease as pressure increases.
- The $pV$ values reach a minimum value depending on the nature of the gas (maximum negative deviation). After this point,increasing the pressure causes the $pV$ value to increase,crossing the ideal gas line where the deviation becomes zero.
- Further increase in pressure leads to a continuous positive deviation.
Thus,it is concluded that "Real gases do not follow the ideal gas equation perfectly under all conditions."