Phosphorus pentafluoride is a known toxic, colorless gas because of its ability to form sp³d hybrid orbitals and bind five fluorine atoms. Nitrogen pentafluoride is only a theoretical compound due to nitrogen's inability to form sp³d hybrid orbitals. The octet rule and valence d orbitals play a role in explaining these differences.
Phosphorus pentafluoride is a known toxic, colorless gas because phosphorus has d orbitals that allow it to form sp³d hybrid orbitals to bind five fluorine atoms. On the other hand, nitrogen pentafluoride is only a theoretical compound because nitrogen lacks valence d orbitals, which prevents it from forming sp³d hybrid orbitals to bind five fluorine atoms.
The octet rule states that atoms tend to gain, lose, or share electrons in order to have a full outer shell with eight electrons. However, there are exceptions to the octet rule, such as when the atom has more than eight electrons in its valence shell. Phosphorus is able to exceed the octet rule because it can use its empty valence d orbitals to accommodate more than eight electrons.
In the case of nitrogen, it does not have valence d orbitals and can only form three bonds and hold one lone pair using sp³ hybrid orbitals. As a result, nitrogen pentafluoride is only a theoretical compound.