The P-Block Elements (Group 15 - Nitrogen Family)

Group 15 Elements

Group 15 elements, also known as the nitrogen family, include Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi), and Moscovium (Mc). This group shows a transition from non-metallic character at the top to metallic character at the bottom.

Occurrence

Nitrogen (N):

Most abundant element in the Earth's atmosphere ($\approx$ 78% by volume).
Exists as $N_2$ gas, which is very unreactive due to the strong triple bond between nitrogen atoms ($N \equiv N$).
Found in various natural compounds like nitrates (e.g., Chile saltpeter, $NaNO_3$), ammonia ($NH_3$), and in all proteins and nucleic acids in living organisms.

Phosphorus (P):

Does not occur in free state in nature due to its high reactivity.
Found in nature mainly as phosphate rocks, such as phosphorit/apatite ($Ca_3(PO_4)_2$).
Essential component of nucleic acids, ATP, and phospholipids in living organisms.

Arsenic (As), Antimony (Sb), Bismuth (Bi):

Occur in nature mainly as sulfides (e.g., arsenopyrite $FeAsS$, stibnite $Sb_2S_3$, bismuthinite $Bi_2S_3$) and oxides.
They are considered semi-metals or metalloids, with metallic properties increasing down the group.

Moscovium (Mc): A synthetic, highly radioactive element with very short half-lives, predicted to be a metal.

Electronic Configuration

General Configuration: The general valence shell electronic configuration of Group 15 elements is $ns^2np^3$.

N: $[He] 2s^22p^3$
P: $[Ne] 3s^23p^3$
As: $[Ar] 3d^{10} 4s^24p^3$
Sb: $[Kr] 4d^{10} 5s^25p^3$
Bi: $[Xe] 4f^{14} 5d^{10} 6s^26p^3$
Mc: $[Og] 7s^27p^3$

Significance: The presence of three electrons in the p-subshell and a stable half-filled $p$-subshell influences their chemical behavior, leading to common oxidation states of -3, +3, and +5.

Atomic Radii

Trend: Atomic radii increase down the group from N to Bi.

Reasons:

Increase in the number of electron shells.
Increased shielding effect of inner electrons.

Comparison to Group 14: Atomic radii are smaller than the corresponding elements in Group 14 due to the greater effective nuclear charge experienced by the valence electrons.

Ionization Enthalpies

Trend: First ionization enthalpies generally decrease down the group from N to Bi.

Reasons:

Increase in atomic size.
Increased shielding effect.

Comparison to Group 14: First ionization enthalpies are generally higher than those of Group 14 elements in the same period. This is because Group 15 elements have a half-filled $p$-subshell ($np^3$), which provides extra stability, making it harder to remove an electron.

Second and Third Ionization Enthalpies: The successive ionization enthalpies increase significantly, but the elements readily form $+3$ and $+5$ oxidation states.

Electronegativity

Trend: Electronegativity values decrease down the group from N to Bi.

N: 3.04
P: 2.19
As: 2.18
Sb: 2.05
Bi: 2.02

Reason: Increase in atomic size and decrease in effective nuclear charge experienced by the valence electrons.

Nitrogen's Electronegativity: Nitrogen is the most electronegative element in this group and is comparable to oxygen.

Physical Properties

States of Matter:

Nitrogen is a gas.
Phosphorus, Arsenic, and Antimony are solids with covalent and metallic characteristics (metalloids).
Bismuth is a brittle, silvery-white metal.
Moscovium is a synthetic radioactive element.

Allotropes:

Nitrogen exists only as $N_2$ molecules.
Phosphorus exists in several allotropic forms: white phosphorus ($P_4$), red phosphorus, and black phosphorus. White phosphorus is the most reactive and toxic.
Arsenic, Antimony, and Bismuth also exhibit allotropy, with different metallic and semi-metallic forms.

Melting and Boiling Points: Melting and boiling points generally increase from N to Sb, but then decrease for Bi. This trend is influenced by the change in bonding from covalent (N, P) to metallic (Sb, Bi) and the presence of allotropes.

Density: Density generally increases down the group.

Chemical Properties

1. Oxidation States:

The most common oxidation states are -3, +3, and +5.
Nitrogen shows a remarkable range of oxidation states from -3 to +5 due to its ability to form multiple bonds and the absence of d-orbitals in its valence shell.
Phosphorus also exhibits -3, +3, +5 states.
As, Sb, Bi commonly show +3 and +5 states, with +3 becoming increasingly stable down the group due to the inert pair effect.

2. Catenation:

Nitrogen shows limited catenation due to the repulsion between lone pairs in adjacent atoms. It forms compounds like $N_2$, $N_3^-$ (azide ion), and limited chain structures.
Phosphorus shows catenation to a greater extent than nitrogen, forming compounds like $P_4$ (in white phosphorus) and phosphanes ($PH_3$, $P_2H_4$, etc.).
Catenation ability further decreases for As and Sb and is negligible for Bi.

3. Reactivity with Hydrogen: They form hydrides, the most important being ammonia ($NH_3$) and phosphine ($PH_3$).

$N_2 + 3H_2 \rightleftharpoons 2NH_3$ (Haber process)
$P_4 + 6H_2 \rightarrow 4PH_3$

4. Reactivity with Oxygen: They form oxides, where they exhibit positive oxidation states.

Nitrogen forms several oxides like $N_2O$, $NO$, $N_2O_3$, $NO_2$, $N_2O_4$, $N_2O_5$.
Phosphorus forms oxides like $P_4O_6$ and $P_4O_{10}$.
Oxides of As, Sb, Bi are generally basic or amphoteric.

5. Reactivity with Halogens: They form halides, commonly showing $+3$ and $+5$ oxidation states.

Nitrogen halides are unstable (e.g., $NF_3$, $NCl_3$).
Phosphorus forms halides like $PCl_3$ and $PCl_5$.
As, Sb, Bi form halides ($AsCl_3$, $AsCl_5$; $SbCl_3$, $SbCl_5$; $BiCl_3$; $BiCl_5$ is unstable).

6. Reducing Nature:

Nitrogen compounds in lower oxidation states (-3, 0, +1, +2) are reducing agents.
Phosphorus compounds in lower oxidation states are reducing agents.
Arsenic and Antimony compounds in +3 state can act as reducing agents.
Bismuth compounds in +3 state are generally oxidizing agents, and Bi(+5) is a strong oxidizing agent.

7. Metallic Character: Metallic character increases down the group.

Important Trends And Anomalous Behaviour Of Carbon

Carbon (C), the first element of Group 14, exhibits anomalous behavior due to its small size, high electronegativity, and absence of d-orbitals in its valence shell. While this was discussed in the context of Group 14, its unique properties are worth reiterating.

Anomalous Properties Of Carbon

1. Non-Metallic Nature: Carbon is a non-metal, while Si and Ge are metalloids, and Sn, Pb, Fl are metals.

2. Small Atomic/Ionic Size: The smallest in Group 14, leading to high charge density.

3. High Ionization Enthalpy: The highest in Group 14, making it harder to lose electrons.

4. High Electronegativity: The highest in Group 14, leading to a preference for forming covalent bonds.

5. Catenation: Unparalleled ability to form long chains and rings of C-C bonds due to the strength of the $C-C$ single bond and formation of stable multiple bonds ($C=C$, $C \equiv C$). This property is the basis of organic chemistry.

6. Formation of Multiple Bonds: Carbon readily forms stable $p\pi-p\pi$ multiple bonds with itself and other second-period elements (like O and N).

7. Absence of d-orbitals: Unlike elements below it in Group 14, carbon lacks vacant d-orbitals in its valence shell ($n=2$). This limits its coordination number to a maximum of 4 and prevents it from forming stable compounds analogous to $SiF_6^{2-}$.

8. Elemental Forms (Allotropes): Carbon exhibits a wide range of allotropes (diamond, graphite, fullerenes, carbon nanotubes, graphene) with vastly different properties.

9. Oxides: Forms two stable oxides, $CO$ and $CO_2$, with vastly different properties (toxic, reducing gas vs. non-toxic, mild oxidizer gas).

10. Hydrides: Forms a vast and stable family of compounds called hydrocarbons and their derivatives, forming the basis of organic chemistry.

11. Oxidation States: Exhibits a wide range of oxidation states, from -4 to +4, with many intermediate states common.