The oxyanion of nitrogen in which it has the highest oxidation state is the nitrate ion \(\left(\mathrm{NO}_{3}^{-}\right)\). The corresponding oxyanion of phosphorus is \(\mathrm{PO}_{4}^{3-}\). The \(\mathrm{NO}_{4}{ }^{3-}\) ion is known but not very stable. The \(\mathrm{PO}_{3}^{-}\) ion is not known. Account for these differences in terms of the bonding in the four anions.

Understanding the concept of oxidation states is crucial when exploring the stability of oxyanions. An element's oxidation state indicates the degree of oxidation or loss of electrons it has undergone when forming a compound or ion. In oxyanions like nitrate (NO_{3}^{-}) and phosphate (PO_{4}^{3-}), the central atoms nitrogen and phosphorus have their highest oxidation states, +5.

This higher oxidation state correlates with a greater number of bonds to oxygen atoms, which typically have a negative charge. In the case of nitrate, the central nitrogen forms bonds with three oxygen atoms, while in phosphate, it forms bonds with four. The high oxidation state also suggests a significant transfer of electron density from the central atom to the surrounding oxygens, contributing to the stability of the ions. In contrast, the instability of nitrate (NO_{4}^{3-}) is partly due to its central nitrogen atom being unable to maintain this higher oxidation state stably across four bonds.

A deeper dive into Lewis structures will shed more light on the significance of electron distribution in these ions.

The notation of Lewis structures plays a pivotal role in visualizing the bonding and electron distribution within a molecule or ion. For oxyanions, these diagrams offer a snapshot of where the electrons are and how they are shared between atoms. In the Lewis structure of an ion like NO_{3}^{-}, for example, nitrogen is represented at the center with three oxygen atoms surrounding it. Each oxygen atom has a negative formal charge, and the nitrogen has a positive charge, equaling the overall negative charge of the ion.

The dative bonds indicated by the Lewis structure help to explain the even distribution of electrons, which is further elaborated upon by resonance in chemistry. Notably, for the less stable NO_{4}^{3-} ion, its Lewis structure would break the octet rule because nitrogen cannot accommodate more than eight electrons in its valence shell. This leads to an unfavorable high energy state which contributes to the ion's instability. In contrast, the Lewis structure of PO_{4}^{3-} respects the octet rule and shows more stable electron configuration.

The stability of oxyanions is greatly influenced by the concept of resonance in chemistry. Resonance refers to the delocalization of electrons across different parts of a molecule or ion. In structures where resonance occurs, such as the NO_{3}^{-} ion, the actual electronic distribution is a hybrid of all possible resonance structures. These structures distribute the charge more evenly across the molecule, leading to a lower overall energy state and resulting in a more stable species. For NO_{3}^{-}, this means that the double bond character is shared among the three nitrogen-oxygen bonds, rather than being localized to just one bond.

The absence of resonance in the NO_{4}^{3-} ion contributes to its instability since the charge and the double bond character cannot be similarly distributed. Furthermore, resonance allows for the effective delocalization of negative charges across the oxygen atoms, as is observed in the PO_{4}^{3-} ion. In contrast, the theoretical PO_{3}^{-} ion, which does not exist in nature, would lack the stabilizing influence of extensive resonance due to the lower electronegativity of phosphorus, explaining its absence.