While nitrogen does have 5 valence electrons, it cannot form 5 bonds like phosphorus due to its electron configuration and bonding tendencies. The valence electron configuration of nitrogen is 2s^2 2p^3, meaning it has three unpaired electrons in its p orbital.
Nitrogen typically forms three bonds by sharing these three unpaired electrons with other atoms. This allows nitrogen to achieve a stable electron configuration with a full outer shell of 8 electrons (known as the octet rule), similar to the noble gas configuration of neon. By sharing these three electrons, nitrogen can form stable covalent bonds with other elements.
In some cases, nitrogen can form additional bonds by accepting electron pairs from other atoms or compounds. This is known as coordinate or dative bonding. However, forming five bonds like phosphorus is not energetically favorable for nitrogen due to its electron configuration. Phosphorus, on the other hand, has an electron configuration of 3s^2 3p^3, with three unpaired electrons in its p orbital and two vacant 3d orbitals. This allows phosphorus to form five bonds and expand its valence shell beyond the octet rule.