Phosphorus is an element found in Group 15 and Period 3 of the periodic table. Its bonding behavior is far more flexible than that of its lighter neighbor, nitrogen. Unlike the elements in Period 2, phosphorus is not strictly confined to the octet rule, which limits an atom to eight valence electrons. Phosphorus commonly forms three or five covalent bonds, but in many common compounds, it can also be found forming four or even six bonds.
The Default State: Three Covalent Bonds
The baseline bonding state for phosphorus is determined by its electron configuration, which is \([\text{Ne}]3s^23p^3\). This arrangement means the phosphorus atom has five valence electrons in its outermost shell, two paired in the \(3s\) orbital and three unpaired in the \(3p\) orbitals. By pairing the three solitary \(3p\) electrons with electrons from other atoms, phosphorus can achieve a stable octet, forming three single covalent bonds.
This bonding configuration is found in the molecule phosphine (\(\text{PH}_3\)). In phosphine, the phosphorus atom is bonded to three hydrogen atoms, resulting in a pyramidal molecular geometry. The remaining pair of electrons in the \(3s\) orbital exists as a non-bonding lone pair, which influences the molecule’s shape and chemical reactivity. This lone pair gives phosphines their characteristic basic properties, as it can be donated to an electron-deficient atom or ion. Molecules like \(\text{PH}_3\) represent the lowest common coordination number for phosphorus, adhering perfectly to the simple octet rule.
The Principle of Expanded Valence
The ability of phosphorus to form more than three covalent bonds stems from its position in Period 3 of the periodic table. The crucial difference is the existence of vacant \(3d\) orbitals in the phosphorus atom’s valence shell. While these orbitals are normally unoccupied, they are close enough in energy to the \(3s\) and \(3p\) orbitals to be accessible for bonding.
When phosphorus is approached by highly electronegative atoms, such as chlorine or fluorine, the energy gained from forming additional strong bonds can outweigh the energy required to rearrange the electrons. This allows an electron from the paired \(3s\) orbital to be promoted into one of the empty \(3d\) orbitals. The promotion creates five unpaired electrons—one in \(3s\), three in \(3p\), and one in \(3d\)—making five bonding sites available.
This process permits phosphorus to utilize an expanded valence shell, allowing it to accommodate more than eight electrons. The five orbitals—one \(s\), three \(p\), and one \(d\)—can then hybridize, combining to form five equivalent hybrid orbitals that point toward the corners of a trigonal bipyramidal structure. This expanded bonding capability is essential for forming compounds with five bonds.
Structural Diversity and Key Configurations
The potential for expanded valence leads to a range of stable compounds where phosphorus exhibits four, five, or six bonds, each with a distinct geometry.
Four Bonds
The four-bonded state is widespread and occurs when the lone pair on the phosphorus atom is used to form a fourth, coordinate covalent bond with an oxygen atom. This is the case in the phosphate ion (\(\text{PO}_4^{3-}\)), which is found in DNA, ATP, and many other biological molecules. In the phosphate ion, the central phosphorus atom is bonded to four oxygen atoms, resulting in a stable tetrahedral structure. Although the phosphorus atom now has four bonds, it has surpassed the octet rule by formally sharing ten valence electrons. A different four-bonded structure is found in phosphonium ions (\(\text{PH}_4^+\)), which form when the lone pair on phosphine accepts a proton (\(\text{H}^+\)), also yielding a tetrahedral geometry.
Five Bonds
The maximum covalent bonding capacity for a neutral phosphorus atom is five, as seen in phosphorus pentachloride (\(\text{PCl}_5\)) or phosphorus pentafluoride (\(\text{PF}_5\)). These molecules feature five single covalent bonds arranged in a trigonal bipyramidal geometry, which arises from the \(sp^3d\) hybridization state. In the solid phase, however, phosphorus pentachloride does not exist as a neutral \(\text{PCl}_5\) molecule but instead as an ionic pair, \(\text{PCl}_4^+\) and \(\text{PCl}_6^-\).
Six Bonds
The six-bonded state is the highest coordination number phosphorus can achieve. It is only possible in certain anionic complexes, such as the hexafluorophosphate ion (\(\text{PF}_6^-\)). In this ion, the phosphorus atom accepts an additional electron pair from a sixth fluorine atom, utilizing two of its empty \(3d\) orbitals to form six equivalent bonds. This configuration results in a highly symmetrical octahedral geometry, demonstrating the full extent of the expanded valence shell.