Question

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.

Short answer

The stability differences between the four anions (NO₃⁻, PO₄³⁻, NO₄³⁻, and PO₃¹⁻) can be attributed to their electron configurations, oxidation states, and resonance structures. Nitrate (NO₃⁻) and phosphate (PO₄³⁻) are stable due to their high oxidation states (+5) and multiple resonance structures that help distribute the charge. However, the unstable ion (NO₄³⁻) has a less favorable oxidation state (+7) despite its resonance structures, making it less stable than nitrate. The non-existing ion (PO₃¹⁻) has a more stable oxidation state (+3) but lacks any resonance structures, so it doesn't exist in nature.

Step-by-step solution

Review Electron Configurations and Oxidation States

First, let's review the electron configurations of nitrogen and phosphorus in their most common oxidation states. Nitrogen belongs to Group 15 and has five valence electrons in its ground state configuration. Phosphorus also belongs to Group 15 and has 5 valence electrons too. - Nitrogen (N): \(1s^2 2s^2 2p^3\) - Phosphorus (P): \(1s^2 2s^2 2p^6 3s^2 3p^3\) Now let's examine the oxidation states for the four anions: - Nitrate (NO₃⁻): N has an oxidation state of +5 - Phosphate (PO₄³⁻): P has an oxidation state of +5 - Unstable ion (NO₄³⁻): N has an oxidation state of +7 - Non-existing ion (PO₃¹⁻): P has an oxidation state of +3

Examine Resonance Structures

Now, we will look at the resonance structures for the four anions to help explain their stability. - Nitrate (NO₃⁻): this ion has three equivalent resonance structures where nitrogen forms a double bond with one oxygen atom and single bonds with the other two. - Phosphate (PO₄³⁻): this ion has four equivalent resonance structures where phosphorus forms single bonds with all four oxygen atoms. - Unstable ion (NO₄³⁻): this ion has four equivalent resonance structures where nitrogen forms double bonds with two of the oxygen atoms and single bonds with the other two. - Non-existing ion (PO₃¹⁻): no resonance structures exist for this ion.

Compare Bonding and Stability

Now that we have examined the electron configurations, oxidation states, and resonance structures, we can analyze the stability differences between the four anions. - Nitrate (NO₃⁻) and phosphate (PO₄³⁻): Nitrogen and phosphorus's high oxidation states (+5) and multiple resonance structures help stabilize these anions by dispersing the charge. - Unstable ion (NO₄³⁻): Despite having multiple resonance structures, nitrogen is in a less favorable oxidation state (+7); thus, this ion is not very stable compared to nitrate. - Non-existing ion (PO₃¹⁻): Although phosphorus is in a more stable oxidation state (+3), it cannot form any resonance structures due to its electron configuration. As a result, this ion does not exist in nature. To summarize, the stability differences between the four anions can be accounted for by analyzing the electron configurations, oxidation states, and resonance structures of the nitrogen and phosphorus atoms involved in bonding. The nitrate and phosphate ions have high oxidation states and multiple resonance structures that help stabilize them. In contrast, the unstable ion (NO₄³⁻) and the non-existing ion (PO₃¹⁻) do not have these favorable characteristics, making them less stable.

Key concepts

Oxidation States

Oxidation states are important in understanding how atoms bond in molecules. They represent the number of electrons an atom gains or loses when it forms a compound. For nitrogen and phosphorus in oxyanions, the oxidation state reflects how these elements bond with oxygen.

In the nitrate ion ( itrate), nitrogen has an oxidation state of +5. This high oxidation state allows nitrogen to form multiple strong bonds with oxygen, thus stabilizing the ion.

• High oxidation states usually allow more bonds with oxygen, enhancing stability.
• In nitrate, this strong bonding compensates for the negative charge carried by the ion.

On the other hand, in nitro-tetraoxy ( [NO_{4}^{3-}] ), nitrogen can reach an oxidation state of +7. Although a very high oxidation state, nitrogen cannot maintain a stable structure at this level, which results in instability.

For the phosphate ion ( [PO_{4}^{3-}] ), phosphorus also achieves a +5 oxidation state, allowing it to stably bond with four oxygen atoms. The high oxidation state works well to stabilize the ion through effective bond formation. However, a proposed ion of phosphorus with an oxidation state of +3, such as PO_{3}^{-}, cannot achieve such stability due to insufficient bonding possibilities.

Resonance Structures

Resonance structures are vital to understanding why some ions are more stable than others. They are different ways of drawing the same molecular structure by shifting electrons, suggesting some degree of electron sharing across the bonds.

• Nitrate ion ( [NO_{3}^{-}] ) has three resonance structures. These structures indicate that electrons are delocalized, which means the electrons are spread out over different atoms.
• Delocalization of electrons results in a lower overall energy, which increases stability.

In contrast, the nitro-tetraoxy ion ( [NO_{4}^{3-}] ) shows a similar ability to form resonance structures but fails to gain enough stability because of the very high and unsuitable oxidation state. Multiple resonance structures can't compensate when the oxidation state doesn't favor stability.

For the phosphate ion ( [PO_{4}^{3-}] ), four resonance structures can be drawn. Like nitrate, this delocalization of electrons contributes to a stable low-energy configuration that leads to the credibility of this ion.

• Resonance is absent in a hypothetical ion like [PO_{3}^- ], limiting electron delocalization.

No resonance structures for [PO_{3}^{-} due to lack of available bonds leaves it too unstable to exist.

Stability of Anions

Anion stability centers largely around how electrons are arranged and dispersed within the molecule. Both electron configuration and the ability to form stable bonds play a pivotal role.

Nitrate ( [NO_{3}^{-}] ) and phosphate ( [PO_{4}^{3-}] ) ions are stabilized by high oxidation states and electron delocalization. This stability is further enhanced by multiple resonance structures, which lower the potential energy:

• They can evenly distribute the negative charge through resonance, decreasing repulsion.
• This makes the ions more energetically favorable or stable.

For unstable ions like nitro-tetraoxy ( [NO_{4}^{3-}] ), despite resonance possibilities, exceedingly high oxidation states make it energetically unfavorable, leading to inherent instability.

The hypothetical phosphorus-based ion ( [PO_{3}^{-}] ) lacks the ability to delocalize electrons, severely impacting its potential to form a stable configuration. Thus, it does not exist naturally, as it cannot maintain a structural integrity enough to sustain itself.

• Anion stability is compromised when resonance and favorable oxidation states are absent.

This absence makes it unlikely for certain oxyanions to form, explaining why some exist while others do not.