At non-boiling point, I understand that when you add heat, temperature increases and also the amount of gas increases, so by using $PV=nRT$ at constant volume, the increase in $n$ and $T$ would increase $P$.
Thus, vapour pressure increases with temperature.
However, at boiling point, if you add heat, temperature does not increase and the energy is used to generate more gas molecules.
So if you are at a equilibrium where there is $\pu{1 mol}$ of $\ce{A(l)}$ and $\pu{1 mol}$ of $\ce{A(g)}$, after you add heat and shift equilibrium to $\pu{0.5 mol}$ of $\ce{A(l)}$ and $\pu{1.5 mol}$ of $\ce{A(g)}$ wouldn’t the vapour pressure increase even though temperature is the same by the ideal gas law– $PV=nRT$, where increase in $n$ increases $P$?
So why is vapour pressure is constant at boiling point?