thermodynamics - Why at constant pressure and temperature Gibbs energy change of a process can be negative?

1. $dG < 0$, for a spontaneous reaction at constant Temperature $T$ and Pressure $P$, this means that at contant $T$ and $P$, $dG$ can be negative


2. $dG = VdP -SdT$, this means that at constant $P$ ($dP =0$) and constant $T$ ($dT = 0$), $dG = V\cdot 0-T\cdot 0 = 0$. At constant $P$ and $T$ according to this equation $dG$ is always zero.

Statement $1$ contradicts statement $2$. Why does this happen?