Q.Write four characteristic properties of p-block elements.

Concept: Electronic configuration and periodic trends of p-block elements

The p-block comprises groups 13–18, where the last electron enters a p-orbital. This placement gives rise to distinctive chemical and physical behavior.

Key properties:

1. Electronic configuration: The valence shell has the general form $ns^2 \, np^{1-6}$, where $n$ is the period number. The number of p-electrons determines the group (e.g., group 15 has $np^3$).

2. Variable oxidation states: Most p-block elements exhibit multiple oxidation states. The highest oxidation state equals the group number minus 10 (e.g., nitrogen in group 15 shows +5, +3, +1, −3). Heavier elements often favor lower oxidation states due to the inert-pair effect.

3. Metallic to non-metallic character: The block contains metals (Al, Ga, In, Tl, Sn, Pb, Bi), metalloids (B, Si, Ge, As, Sb, Te), and non-metals (C, N, O, halogens, noble gases). Metallic character increases down a group and decreases across a period. …

The p-block occupies groups 13–18 of the periodic table, where the last electron enters a p-orbital. These elements bridge the metallic s-block and the transition d-block, giving rise to the richest variety of chemical behaviour in the periodic table. Understanding their properties means recognising how the progressive filling of p-orbitals—from $p^1$ to $p^6$—shapes reactivity, bonding, and physical character.

Four Characteristic Properties

1. Valence-electron configuration and the role of p-orbitals

Every p-block element has its outermost electrons in an $ns^2 \, np^{1-6}$ configuration. The three p-orbitals ($p_x, p_y, p_z$) can hold a maximum of six electrons, so the block spans six groups. This configuration directly determines the number of valence electrons available for bonding—ranging from three in group 13 (boron family) to eight in group 18 (noble gases). The half-filled ($p^3$) and fully filled ($p^6$) configurations confer extra stability, explaining why nitrogen is relatively inert and the noble gases are chemically unreactive under ordinary conditions.

2. Variable oxidation states and the inert-pair effect

Unlike s-block elements, which typically show a single oxidation state, p-block elements exhibit multiple oxidation states. For instance, nitrogen ranges from $-3$ in ammonia to $+5$ in nitric acid, and lead shows both $+2$ and $+4$. As you descend a group, the lower oxidation state (corresponding to the loss or sharing of only the p-electrons, leaving the $ns^2$ pair inert) becomes increasingly stable—a phenomenon called the inert-pair effect. This is why thallium(I) is more stable than thallium(III), and lead(II) more stable than lead(IV). The reluctance of the $s^2$ pair to participate in bonding arises from poor shielding by intervening d- and f-electrons and relativistic contraction of the s-orbital in heavier elements.

3. Gradation from metallic to non-metallic character

The p-block contains metals (like aluminium, tin, lead), metalloids (boron, silicon, germanium, arsenic), and non-metals (carbon, nitrogen, oxygen, the halogens). Within a period, metallic character decreases from left to right as ionisation energy rises and electronegativity increases; within a group, it increases down as atomic size grows and ionisation energy falls. This diagonal trend produces the "staircase" boundary between metals and non-metals. The metalloids, straddling this line, show intermediate properties—semiconducting behaviour, amphoteric oxides—that make them technologically vital (silicon in electronics, for example).

4. Diverse bonding modes and molecular structures …