Why at extremely low pressures, the real gases obey the ideal gas equation?

At low pressures, volume $V$ is very large and hence the correction term b (a constant of small value) can be neglected in comparison to the very large value of $V$. Thus the van der Waal's equation for $1$ mole of a real gas. 

$(P+\frac{a}{V^2})(V-b)=RT$ may be written as $(P+\frac{a}{V^2})(V)=RT$

or $PV+\frac{a}{V}=RT$

or $PV=RT-\frac{a}{V}$

For large $V$ (at very low pressures) $\frac{a}{V}$ is very small and can be ignored.

$\therefore$ $PV$ becomes $RT$ at very low pressures.