Why does Nitrogen not form pentahalides like $NCl_5$, while Phosphorus forms $PCl_5$?

Nitrogen, being the first element of Group 15 and belonging to the second period, lacks vacant d-orbitals in its valence shell. Therefore, it cannot expand its octet beyond four bonds. Its maximum covalency is four (e.g., in $NH_4^+$). Phosphorus, on the other hand, is in the third period and possesses vacant 3d-orbitals. It can utilize these d-orbitals to expand its octet and form five covalent bonds, thus forming pentahalides like $PCl_5$.

Explain the inert pair effect and its impact on the oxidation states of Group 15 elements.

The inert pair effect refers to the reluctance of the $ns^2$ valence electrons to participate in chemical bonding, especially in heavier p-block elements. As we move down Group 15, the effective nuclear charge on the $ns^2$ electrons increases due to poor shielding by intervening d- and f-electrons. This makes the $ns^2$ electrons more tightly bound and less available for bonding. Consequently, the stability of the +3 oxidation state (involving only $np^3$ electrons) increases, while the stability of the +5 oxidation state (involving $ns^2 np^3$ electrons) decreases for elements like Antimony and Bismuth.

How does the acidic character of oxides change down Group 15?

The acidic character of oxides of Group 15 elements generally decreases down the group. For oxides in the same oxidation state, say +3, $N_2O_3$ and $P_2O_3$ are acidic. $As_2O_3$ and $Sb_2O_3$ are amphoteric, meaning they can react with both acids and bases. $Bi_2O_3$ is distinctly basic. This trend is consistent with the increasing metallic character of the elements down the group; metallic oxides are typically basic, while non-metallic oxides are acidic.

What is the trend in the thermal stability and reducing character of Group 15 hydrides ($EH_3$)?

The thermal stability of Group 15 hydrides ($NH_3, PH_3, AsH_3, SbH_3, BiH_3$) decreases down the group. This is because as the size of the central atom (E) increases, the E-H bond length increases, and the bond strength decreases, making them easier to decompose upon heating. Conversely, the reducing character of these hydrides increases down the group. Less stable hydrides readily decompose to release hydrogen, which acts as a reducing agent, making $BiH_3$ the strongest reducing agent among them.

Why is Nitrogen a gas, while Phosphorus is a solid at room temperature?

Nitrogen exists as a diatomic molecule ($N_2$) with a triple bond between the two nitrogen atoms. This strong covalent bond results in a very stable molecule. However, the intermolecular forces between $N_2$ molecules are very weak van der Waals forces, leading to a low boiling point and making it a gas. Phosphorus, on the other hand, exists as discrete $P_4$ tetrahedral molecules (white phosphorus) or polymeric structures (red and black phosphorus). The intermolecular forces between $P_4$ molecules are much stronger than those between $N_2$ molecules, requiring more energy to overcome, hence phosphorus is a solid at room temperature.

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Version 1Updated 22 Mar 2026

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Group 15 elements, also known as pnictogens, comprise Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), and Bismuth (Bi). Their general valence shell electronic configuration is $ns^2 np^3$ . This configuration dictates their characteristic oxidation states, primarily -3, +3, and +5. The stability of these oxidation states, along with various physical and chemical properties such as atomic…

Quick Summary

Group 15 elements (N, P, As, Sb, Bi) have a valence shell configuration of $ns^2 np^3$ , giving them 5 valence electrons. Their characteristic oxidation states are -3, +3, and +5. Nitrogen, due to its small size and lack of d-orbitals, exhibits a wide range of oxidation states but cannot form pentavalent compounds like $PCl_5$ .

The stability of the +5 oxidation state decreases down the group, while the +3 oxidation state stability increases, a phenomenon attributed to the inert pair effect, especially prominent for Sb and Bi.

Physically, atomic size, metallic character, and density increase down the group, while ionization enthalpy and electronegativity decrease. Melting and boiling points show a more complex trend, peaking around Arsenic.

Chemically, the thermal stability of hydrides ( $EH_3$ ) decreases, but their reducing character and acidic nature of oxides ( $E_2O_3, E_2O_5$ ) decrease down the group, reflecting the transition from non-metallic to metallic character.

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Key Concepts

Oxidation States in Group 15

Group 15 elements have 5 valence electrons ( $ns^2 np^3$ ). They typically exhibit -3, +3, and +5 oxidation…

Inert Pair Effect and Stability of Oxidation States

The inert pair effect is crucial for understanding the stability of +3 and +5 oxidation states. For heavier…

Acidic/Basic Nature of Oxides

The nature of oxides of Group 15 elements changes systematically down the group, correlating with the…

Electronic Configuration —

ns^2 np^3

Oxidation States — 3, +3, +5 (N: wide range, no +5 ionic)

Inert Pair Effect — Stability of +3 increases, +5 decreases down the group (esp. for Sb, Bi)

Atomic/Ionic Radii — Increases down group

Ionization Enthalpy — Decreases down group

Electronegativity — Decreases down group

Metallic Character — Increases down group (N, P non-metals; As, Sb metalloids; Bi metal)

Hydrides ($EH_3$)

- Thermal Stability: $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$ - Reducing Character: $NH_3 < PH_3 < AsH_3 < SbH_3 < BiH_3$ - Basicity: $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$

Oxides ($E_2O_3$)

- Acidic Character: $N_2O_3 > P_2O_3 > As_2O_3 > Sb_2O_3 > Bi_2O_3$ (Acidic $ightarrow$ Amphoteric $ightarrow$ Basic)

Halides

- $EX_3$ formed by all. - $EX_5$ formed by P, As, Sb. N does not form $EX_5$ . $BiF_5$ exists but unstable.

For Group 15 Hydrides ( $EH_3$ ) trends: Through Reality, Bonds Decrease. (Thermal stability Decreases, Reducing character Increases, Basicity Decreases, Bond angle Decreases).

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