5.1 Properties and Trends of Groups 13-18 Elements

Periodic Trends and Properties

The p-block (Groups 13–18) spans a remarkable range of character: from the metallic aluminum and thallium of Group 13 to the nearly inert noble gases of Group 18. Grasping the trends that run across and down this block lets you predict bonding behavior, reactivity, and stable oxidation states without memorizing every element individually.

Electron Configuration and Atomic Structure

Every p-block element adds electrons to its outermost p subshell, giving a general valence configuration of $ns^{2}\,np^{1\text{–}6}$. The number of valence p electrons rises from 1 in Group 13 to 6 in Group 18.

• Group 13 (e.g., B, Al): $ns^{2}\,np^{1}$ → 3 valence electrons
• Group 14 (e.g., C, Si): $ns^{2}\,np^{2}$ → 4 valence electrons
• Group 15 (e.g., N, P): $ns^{2}\,np^{3}$ → 5 valence electrons
• Group 16 (e.g., O, S): $ns^{2}\,np^{4}$ → 6 valence electrons
• Group 17 (e.g., F, Cl): $ns^{2}\,np^{5}$ → 7 valence electrons
• Group 18 (e.g., Ne, Ar): $ns^{2}\,np^{6}$ → 8 valence electrons (full shell, hence their stability)

The filling order follows the Aufbau principle, but keep an eye on heavier elements (Period 6 onward) where relativistic effects and poor shielding by 4f/5d electrons start to matter.

Atomic Properties and Periodic Trends

The standard periodic trends apply across the p-block, but a few details are worth highlighting:

• Atomic radius decreases left → right (increasing $Z_{\text{eff}}$) and increases top → bottom (new shells added). For example, boron's covalent radius (~84 pm) is much smaller than thallium's (~145 pm).
• Ionization energy (IE) increases across a period and decreases down a group. A notable exception: IE drops slightly from Group 15 to Group 16 in a given period because the half-filled p³ configuration is extra stable (e.g., N has a higher first IE than O).
• Electronegativity follows the same general pattern (increases across, decreases down). Fluorine tops the scale at 3.98 on the Pauling scale.
• Electron affinity generally becomes more negative (more favorable) across a period, but Group 15 and Group 18 elements are exceptions because adding an electron disrupts a half-filled or fully filled subshell.
• Metallic character increases down each group. Group 13 shifts from the metalloid boron to the clearly metallic thallium; Group 14 goes from nonmetallic carbon to metallic lead.

Electronegativity and Chemical Bonding

Electronegativity (Pauling scale) tells you how strongly an atom pulls on shared electrons. This single number drives a lot of bonding predictions:

• A large electronegativity difference (roughly > 1.7) between two bonded atoms usually means ionic character dominates.
• A small difference (< 1.7) points toward covalent bonding.
• These cutoffs are guidelines, not hard rules. The boundary between ionic and covalent is a continuum, and factors like oxidation state and coordination environment also matter.

Within the p-block, electronegativity drops sharply from fluorine (3.98) to iodine (2.66) down Group 17, and from oxygen (3.44) to tellurium (2.10) down Group 16. That drop explains why heavier p-block elements form more polar or even ionic compounds with electropositive metals, while lighter ones tend toward covalent bonding.

Bonding and Oxidation States

Oxidation States Across the p-Block

Each group has a maximum (group) oxidation state equal to the group number minus 10 (for Groups 13–18, that's +3 through +8, though +8 is essentially only seen for xenon in a few exotic compounds). The key patterns:

Notice that the lower oxidation state becomes more stable as you move down a group (Tl prefers +1, Pb prefers +2, Bi prefers +3). This is the inert pair effect.

The Inert Pair Effect

This is one of the most important concepts for heavy p-block chemistry. Here's what's happening:

1. In heavier elements (Period 6 especially), the two $ns^{2}$ electrons are held more tightly due to poor shielding by filled 4f and 5d subshells and relativistic contraction of the s orbital.
2. These s electrons become reluctant to participate in bonding.
3. The result: the element "keeps" its s² pair and displays an oxidation state two units lower than the group maximum.

Concrete examples:

• Thallium (Group 13): Tl⁺ (keeping the 6s²) is far more stable than Tl³⁺ in aqueous solution.
• Lead (Group 14): Pb²⁺ compounds are common and stable; Pb⁴⁺ compounds like $PbO_2$ are strong oxidizers because they want to revert to Pb²⁺.
• Bismuth (Group 15): Bi³⁺ is the dominant state; Bi⁵⁺ (as in $NaBiO_3$) is a powerful oxidizing agent.

Covalent and Ionic Bonding in the p-Block

• Upper-right p-block (C, N, O, F, Cl): high electronegativities, small radii → predominantly covalent bonding with each other and with hydrogen.
• Lower-left p-block (Tl, Pb, Bi): lower electronegativities, larger radii, metallic character → more ionic bonding with nonmetals.
• The octet rule works well for Period 2 elements but breaks down for Period 3 and below, where elements can use d orbitals (or, more accurately, form hypervalent species) to exceed an octet (e.g., $PCl_5$, $SF_6$).

Multiple Bonding Capability

Period 2 elements (C, N, O) readily form strong $\pi$ bonds because their small atomic radii allow effective lateral overlap of p orbitals. As you move down a group, the larger, more diffuse p orbitals overlap poorly, so multiple bonding weakens:

• $N_2$ has a very strong $N{\equiv}N$ triple bond (945 kJ/mol). Phosphorus, by contrast, prefers single-bonded tetrahedral structures ($P_4$) rather than forming $P{\equiv}P$.
• $O_2$ is a stable diatomic with a double bond. Sulfur instead forms $S_8$ rings with single bonds.

This trend directly explains many allotropic differences (see below) and why heavier p-block elements favor catenation through single bonds or cluster structures.

Reactivity and Allotropy

Chemical Reactivity Patterns

Reactivity trends in the p-block depend on whether you're looking at the metallic or nonmetallic members:

• Metals (lower members of Groups 13–15): Reactivity generally increases down the group because ionization energies drop, making it easier to lose electrons. Aluminum is passivated by its oxide layer but reacts vigorously once that layer is breached. Thallium is reactive enough to tarnish in air.
• Nonmetals (Groups 15–17, upper members): Reactivity generally decreases down the group. Fluorine is the most reactive element on the periodic table; iodine is comparatively mild. This tracks with decreasing electron affinity and increasing bond lengths in the elemental form.
• Noble gases (Group 18): Their filled valence shells make them largely unreactive. Xenon is the notable exception, forming stable fluorides ($XeF_2$, $XeF_4$, $XeF_6$) and oxides because its large size and low ionization energy allow it to expand its coordination sphere.

Reactivity is also shaped by kinetic factors. Nitrogen is thermodynamically reactive (many nitrogen compounds are less stable than $N_2$), but the very strong triple bond creates a high activation barrier, making $N_2$ kinetically inert at room temperature.

Allotropy and Physical Properties

Allotropy refers to the existence of an element in two or more structurally distinct forms in the same physical state. Allotropes differ in how atoms are bonded and arranged, which can dramatically change physical and chemical properties.

Carbon is the classic example:

• Diamond: each carbon is sp³ hybridized in a 3D tetrahedral network → extremely hard, electrically insulating, transparent.
• Graphite: sp² layers with delocalized $\pi$ electrons → soft, electrically conductive along layers, opaque.
• Fullerenes ($C_{60}$) and graphene: curved or single-layer sp² structures with unique electronic properties.

Phosphorus illustrates how allotrope stability relates to structure:

• White phosphorus ($P_4$ tetrahedra): highly reactive and toxic because of strained 60° bond angles.
• Red phosphorus: polymeric chains of linked $P_4$ units with relieved strain → much more stable, less toxic.
• Black phosphorus: layered structure similar to graphite → most thermodynamically stable allotrope, semiconducting.

Sulfur forms $S_8$ rings (rhombic and monoclinic allotropes) at standard conditions. Heating breaks these rings into chains, producing plastic sulfur.

Oxygen exists as $O_2$ (essential for respiration) and $O_3$ (ozone, a powerful oxidizer that absorbs UV radiation in the stratosphere). The difference is simply two atoms vs. three, but the bent geometry and weaker bonding in ozone make it far more reactive.

Temperature and pressure govern which allotrope is thermodynamically favored. High pressure converts graphite to diamond; gentle heating converts white phosphorus to red.