The p-Block Elements
Six groups, the whole range of chemistry — metals, metalloids and non-metals from boron to the noble gases, with the trends that tie them together and the anomalies that make each group its own world
- Why the p-block spans metals, metalloids and non-metals, and how its trends run.
- The inert pair effect and why lower oxidation states grow stable down a group.
- The signature chemistry of each group — diborane, catenation, ammonia & nitric acid, ozone & sulphuric acid, the halogens, and the noble gases.
- Key industrial processes: the Ostwald and Contact processes.
- The oxidising power order of the halogens and the strange compounds of xenon.
- The anomalous first members (B, C, N, O, F) and why they differ from their families.
The p-Block & Its Trends
The p-block covers Groups 13–18, where the differentiating electron enters a \(p\) orbital, giving the general configuration \(ns^2np^{1\text{–}6}\). It is the only block holding all three kinds of element — metals (lower left), metalloids along the diagonal, and non-metals (upper right). Across each period non-metallic character rises; down each group metallic character returns.
The group's highest oxidation state uses all the valence electrons. Down each group, the lower state — two units smaller — becomes increasingly favoured as the \(ns^2\) pair turns reluctant to bond.
Group 13: The Boron Family
Group 13 (\(\ce{B, Al, Ga, In, Tl}\), \(ns^2np^1\)) opens with the metalloid boron and continues with metals. The usual oxidation state is \(+3\), but \(\ce{Tl+}\) (inert pair) is the most stable thallium ion. Boron is the anomalous first member: small, with a high ionization enthalpy, it forms only covalent, electron-deficient compounds.
| Compound | Formula | Note |
|---|---|---|
| Borax | \(\ce{Na2B4O7.10H2O}\) | borax-bead test for coloured ions |
| Orthoboric acid | \(\ce{H3BO3}\) | weak monobasic Lewis acid (accepts \(\ce{OH-}\)) |
| Diborane | \(\ce{B2H6}\) | electron-deficient; banana bonds |
| Alum | \(\ce{KAl(SO4)2.12H2O}\) | water purification, mordant |
Group 14: The Carbon Family
Group 14 (\(\ce{C, Si, Ge, Sn, Pb}\), \(ns^2np^2\)) runs from the non-metal carbon through metalloids to the metals tin and lead. The hallmark is catenation — the ability to form long chains of self-bonds — which is supreme in carbon and fades down the group as the element–element bond weakens.
Carbon's strong \(\ce{C-C}\) bonds make organic chemistry possible. Lead prefers \(+2\): \(\ce{PbCl4}\) is unstable while \(\ce{PbCl2}\) is not.
| Species | Character | Note |
|---|---|---|
| Diamond, graphite, fullerene | allotropes of carbon | sp³ giant, sp² layers, sp² cage (\(\ce{C60}\)) |
| \(\ce{CO}\) | neutral oxide | poisonous, strong reductant & ligand |
| \(\ce{CO2}\) | acidic oxide | greenhouse gas, linear |
| \(\ce{SiO2}\) | giant covalent | basis of silicates & glass |
| Silicones | \(\ce{(R2SiO)_n}\) | water-repellent polymers |
Group 15: The Nitrogen Family
Group 15 (\(\ce{N, P, As, Sb, Bi}\), \(ns^2np^3\)) shows the widest range of oxidation states, from \(-3\) to \(+5\). Nitrogen, locked in a triple bond \(\ce{N#N}\), is inert until fixed; phosphorus comes as reactive white (\(\ce{P4}\) tetrahedra) and stable red allotropes.
Ammonia (pyramidal, basic) is the gateway to nitric acid: catalytic oxidation gives \(\ce{NO}\), then \(\ce{NO2}\), which dissolves in water to \(\ce{HNO3}\) — the Ostwald process.
| Oxide | N state | Oxide | N state |
|---|---|---|---|
| \(\ce{N2O}\) | +1 | \(\ce{NO2}\) | +4 |
| \(\ce{NO}\) | +2 | \(\ce{N2O5}\) | +5 |
| \(\ce{N2O3}\) | +3 | \(\ce{PH3}\) (phosphine) | pyramidal, weakly basic |
Group 16: The Oxygen Family (Chalcogens)
Group 16 (\(\ce{O, S, Se, Te, Po}\), \(ns^2np^4\)) — the chalcogens, the "ore-formers". Oxygen exists as \(\ce{O2}\) and the bent, reactive allotrope ozone \(\ce{O3}\); sulphur comes as rhombic and monoclinic forms built from \(\ce{S8}\) crown rings.
\(\ce{SO3}\) is absorbed in concentrated acid (not water, to avoid a fog) as oleum, then diluted. "King of chemicals," \(\ce{H2SO4}\) is a strong acid, dehydrating agent and oxidiser.
Group 17: The Halogens
Group 17 (\(\ce{F, Cl, Br, I, At}\), \(ns^2np^5\)) holds the most reactive non-metals — one electron short of a noble-gas shell, so they are powerful oxidising agents. They show the \(-1\) state throughout and positive states up to \(+7\) (except fluorine, which, being the most electronegative element, shows only \(-1\)).
A halogen displaces any halide below it: \(\ce{Cl2 + 2KBr -> 2KCl + Br2}\). The hydrohalic acids strengthen down the group because the \(\ce{H-X}\) bond weakens, despite falling electronegativity.
| Species | What it is | Note |
|---|---|---|
| Interhalogens | \(\ce{ClF3},\ \ce{IF7}\) | more reactive than parent halogens |
| Bleaching powder | \(\ce{CaOCl2}\) | from \(\ce{Cl2}\) + slaked lime; oxidising bleach |
| \(\ce{HF}\) | weak acid | etches glass; anomalous (H-bonding) |
| \(\ce{HClO4}\) | perchloric acid | strongest common oxoacid of Cl |
Group 18: The Noble Gases
Group 18 (\(\ce{He, Ne, Ar, Kr, Xe, Rn}\)) have complete octets (\(ns^2np^6\), He is \(1s^2\)), the highest ionization enthalpies in each period, and almost no chemistry. For decades they were thought wholly inert — until 1962, when xenon, the largest non-radioactive member with the lowest ionization enthalpy, was made to react with fluorine.
| Compound | Hybridisation / shape |
|---|---|
| \(\ce{XeF2}\) | \(sp^3d\), linear |
| \(\ce{XeF4}\) | \(sp^3d^2\), square planar |
| \(\ce{XeF6}\) | \(sp^3d^3\), distorted octahedral |
The Inert Pair Effect & First-Member Anomalies
Two cross-cutting ideas explain most p-block surprises. The inert pair effect is the growing reluctance of the \(ns^2\) electrons to take part in bonding as a group is descended, so the lower oxidation state (two below the group state) becomes more stable for the heaviest members.
| Group | Higher state | Stable lower state (heavy element) |
|---|---|---|
| 13 | \(+3\) | \(\ce{Tl+}\) |
| 14 | \(+4\) | \(\ce{Pb^2+}\) |
| 15 | \(+5\) | \(\ce{Bi^3+}\) |
The first member of every p-block group (\(\ce{B, C, N, O, F}\)) is anomalous — it is unusually small, highly electronegative, and has no available \(d\)-orbitals. It therefore caps its covalency at four, forms strong \(p\pi\!-\!p\pi\) multiple bonds (\(\ce{N#N}\), \(\ce{O=O}\)) that its heavier congeners cannot, and often resembles the element diagonally below it.
Putting It to Work
Problem. Which is the more stable ion, \(\ce{Tl+}\) or \(\ce{Tl^3+}\)? Explain.
Solution. The \(6s^2\) pair is reluctant to bond in the heaviest member:
Problem. Why is \(\ce{B2H6}\) called electron-deficient, and how does it bond?
Solution. Count: 12 valence electrons for 8 bonds — too few for normal 2-electron bonds:
Problem. Arrange \(\ce{C},\ \ce{Si},\ \ce{Ge},\ \ce{Pb}\) by decreasing tendency to catenate.
Solution. Catenation falls as the element–element bond weakens down the group:
Problem. Will chlorine displace bromine from \(\ce{KBr}\)? Write the reaction.
Solution. \(\ce{Cl2}\) is a stronger oxidant than \(\ce{Br2}\), so yes:
Problem. Give the hybridisation and shape of \(\ce{XeF4}\).
Solution. Xe has 4 bond pairs + 2 lone pairs → 6 electron domains:
Problem. Arrange \(\ce{HF},\ \ce{HCl},\ \ce{HBr},\ \ce{HI}\) by increasing acid strength and explain.
Solution. Acid strength tracks the weakening \(\ce{H-X}\) bond down the group:
Chapter Summary
Groups 13–18, \(ns^2np^{1\text{–}6}\); metals, metalloids and non-metals together.
Electron-deficient diborane; catenation supreme in carbon; CO/CO₂, silicones.
Ammonia → nitric acid (Ostwald); ozone; sulphuric acid (Contact process).
Oxidising power \(\ce{F2}>\ce{Cl2}>\ce{Br2}>\ce{I2}\); acid strength \(\ce{HF}<...<\ce{HI}\).
Inert octets; only xenon reacts — \(\ce{XeF2},\ \ce{XeF4},\ \ce{XeF6}\).
Inert pair effect (lower state stable down group) and anomalous first members.
Problems
For each item, first decide which group or cross-cutting idea it tests, then apply the relevant rule. Difficulty rises down the list.
- Give the general electronic configuration of the p-block and state why it contains all three element types.
- Explain the inert pair effect and give one example from Group 13 and one from Group 14.
- Why is \(\ce{B2H6}\) electron-deficient? Describe its bridging bonds.
- Explain why orthoboric acid is a weak monobasic acid.
- Arrange \(\ce{C},\ \ce{Si},\ \ce{Sn},\ \ce{Pb}\) by tendency to catenate and justify.
- Distinguish \(\ce{CO}\) and \(\ce{CO2}\) in terms of bonding and acid–base character.
- Write the steps of the Ostwald process for manufacturing nitric acid.
- List the oxides of nitrogen with the oxidation state of nitrogen in each.
- Draw the structure of ozone and explain why it is a strong oxidising agent.
- Write the Contact-process reactions and explain why \(\ce{SO3}\) is absorbed in acid, not water.
- Arrange \(\ce{HF},\ \ce{HCl},\ \ce{HBr},\ \ce{HI}\) by acid strength and explain the trend.
- Give the hybridisation and shape of \(\ce{XeF2},\ \ce{XeF4}\) and \(\ce{XeF6}\).