p-Block elements are those where the last electron enters a p-orbital, spanning groups 13 to 18. Valence shell configuration is $ns^2 np^{1-6}$. This block contains metals, metalloids, and non-metals, and shows the widest range of chemical behaviour of any block in the periodic table.
| Group | Valence Config | Common Oxidation States |
|---|---|---|
| 13 | ns²np¹ | +3, +1 (heavier members) |
| 14 | ns²np² | +4, +2 (heavier members) |
| 15 | ns²np³ | −3, +3, +5 |
| 16 | ns²np⁴ | −2, +2, +4, +6 |
| 17 | ns²np⁵ | −1, +1, +3, +5, +7 |
| 18 | ns²np⁶ | 0 (mostly), +2 (Xe compounds) |
Down any p-block group, the tendency of the outer $ns^2$ electron pair to remain unreactive (not participate in bonding) increases — this is the inert pair effect. It explains why heavier members show a stable oxidation state that is 2 less than the group's maximum (e.g., Tl⁺ more stable than Tl³⁺; Pb²⁺ more stable than Pb⁴⁺).
Members: Boron (metalloid), Aluminium, Gallium, Indium, Thallium (metals). Boron differs sharply from the rest of the group due to its very small size and higher ionization enthalpy — it forms only covalent compounds, while other members show increasingly ionic character down the group.
A white crystalline solid used as a flux and in the borax bead test for identifying metal ions. On heating, it loses water and swells up, then melts to a transparent liquid (sodium metaborate + boric anhydride) that dissolves metal oxides to give characteristic coloured borate beads.
A weak monobasic acid — but unusually, it doesn't ionize by releasing its own proton. Instead, it acts as a Lewis acid, accepting a hydroxide ion from water: $\text{B(OH)}_3 + \text{H}_2\text{O} \rightleftharpoons [\text{B(OH)}_4]^- + \text{H}^+$.
Has an unusual structure with two bridging hydrogen atoms connecting the two boron atoms via 3-centre-2-electron bonds (banana bonds) — each bridging bond uses only 2 electrons to hold together 3 atoms (B–H–B), unlike normal 2-centre-2-electron bonds.
Explain why boric acid, H₃BO₃, behaves as a monobasic acid despite having 3 hydroxyl groups.
Boric acid does not release H⁺ directly from any of its O–H bonds. Instead, the boron atom (electron-deficient, incomplete octet) accepts a hydroxide ion (OH⁻) from a water molecule, releasing one H⁺ in the process — since only one proton is released per molecule via this indirect Lewis acid mechanism, it behaves as monobasic, not tribasic.
Members: Carbon (non-metal), Silicon & Germanium (metalloids), Tin & Lead (metals). Carbon shows a unique property called catenation — the ability to form long chains and rings by bonding to itself — to a far greater extent than any other element, due to the small size and strength of C–C bonds.
| Allotrope | Structure | Key Property |
|---|---|---|
| Diamond | 3D network, each C is sp³ hybridised, bonded to 4 others | Extremely hard; poor electrical conductor |
| Graphite | Layered sheets of sp² hybridised C in hexagonal rings | Soft, slippery (layers slide); good electrical conductor (delocalised electrons) |
| Fullerenes (e.g., C₆₀) | Closed, cage-like spherical/ellipsoidal structures | Only pure, well-defined molecular form of carbon |
Formed by incomplete combustion of carbon; highly toxic since it binds haemoglobin roughly 200 times more strongly than oxygen, blocking oxygen transport in blood.
Synthetic organosilicon polymers with a repeating $-\text{Si-O-Si-}$ backbone, water-repellent and heat-resistant — used as sealants, lubricants, and in medical implants.
Why does graphite conduct electricity while diamond does not, even though both are pure carbon?
In graphite, each carbon is sp² hybridised and bonded to only 3 neighbours, leaving one electron per carbon delocalised across the hexagonal sheet — these free, mobile electrons conduct electricity along the layers. In diamond, every carbon is sp³ hybridised and uses all 4 valence electrons in strong, localised sigma bonds to 4 neighbours, leaving no free electrons to carry current.
Members: Nitrogen, Phosphorus (non-metals), Arsenic, Antimony (metalloids), Bismuth (metal). Nitrogen exists as a stable diatomic gas (N≡N, very high bond enthalpy) and is comparatively unreactive, while phosphorus and heavier members show more varied reactivity and multiple allotropic forms.
Industrially prepared via the Haber process: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ (high pressure, moderate temperature, iron catalyst). NH₃ is a Lewis base (lone pair on N), pyramidal in shape (sp³ hybridised N), and highly soluble in water due to hydrogen bonding.
Industrially prepared via Ostwald's process (catalytic oxidation of ammonia). A strong oxidising acid — dilute HNO₃ reacts with most metals to give NO, while concentrated HNO₃ gives NO₂; both iron and aluminium become passive in concentrated HNO₃ due to a protective oxide layer forming on the surface.
| Allotrope | Structure | Property |
|---|---|---|
| White phosphorus | Discrete P₄ tetrahedral molecules | Highly reactive, poisonous, glows in dark (chemiluminescence), spontaneously ignites in air |
| Red phosphorus | Polymeric chain structure | Less reactive, non-poisonous, does not glow, more stable |
Explain why concentrated nitric acid renders iron passive, preventing further reaction.
Concentrated HNO₃ oxidises the surface of iron to form a thin, tightly adhering layer of iron oxide, which acts as a protective barrier preventing the acid from reaching and reacting with the underlying metal — this phenomenon is called passivation.
Members: Oxygen, Sulphur, Selenium, Tellurium (non-metals/metalloids), Polonium (radioactive metal). Oxygen shows anomalous behaviour compared to the rest of the group due to its small size and high electronegativity, similar to the general first-member anomaly seen across p-block groups.
A bent, angular molecule (V-shaped), a resonance hybrid of two structures with the central O bonded differently to each terminal O. Acts as a powerful oxidising agent, and the stratospheric ozone layer absorbs harmful UV radiation, protecting life on Earth.
Sulphur is burnt to $\text{SO}_2$, catalytically oxidised (V₂O₅ catalyst) to $\text{SO}_3$, which is absorbed in concentrated H₂SO₄ to form oleum, later diluted to the required concentration of sulphuric acid.
A strong dibasic acid, powerful dehydrating agent (chars sugar and organic matter by removing water), and a moderately strong oxidising agent when hot and concentrated.
Most common: rhombic sulphur (S₈ crown-shaped rings, stable at room temperature) and monoclinic sulphur (also S₈ rings, stable only above 96°C).
Explain why concentrated sulphuric acid is used to dry gases and char sugar.
Concentrated H₂SO₄ has an extremely strong affinity for water, absorbing moisture from gases passed through it (used as a drying agent for gases that don't react with it). With sugar (a carbohydrate), it removes water (H and O in a 2:1 ratio) directly from the compound's own molecular structure, leaving behind a black carbon residue — this dehydrating action is called charring.
Members: Fluorine, Chlorine, Bromine, Iodine, Astatine. All are highly reactive non-metals with 7 valence electrons, needing just one more electron to achieve a stable noble gas configuration — making them strong oxidising agents, with oxidising power decreasing down the group (F₂ > Cl₂ > Br₂ > I₂).
For chlorine's oxoacids, acid strength increases with the number of oxygen atoms: $\text{HOCl} < \text{HOClO} < \text{HOClO}_2 < \text{HOClO}_3$ (hypochlorous < chlorous < chloric < perchloric acid) — more oxygen atoms withdraw more electron density, stabilising the conjugate base (better dispersal of negative charge) after the acid loses its proton.
Compounds formed between two different halogens (e.g., ClF, BrF₃, IF₇) — generally more reactive than the individual halogens themselves, since the X-X' bond is weaker than the bond in either pure halogen (due to the halogens' differing sizes and electronegativities).
Explain why fluorine, despite being the most electronegative element, has an unexpectedly low bond dissociation enthalpy for F₂.
Fluorine atoms are very small, so in the F₂ molecule, the lone pairs on each fluorine atom are packed very close together, causing strong lone pair-lone pair repulsion — this repulsion weakens the F–F bond significantly, giving it a lower bond dissociation enthalpy than expected, and contributing to fluorine's extreme reactivity.
Members: Helium, Neon, Argon, Krypton, Xenon, Radon. All have completely filled valence shells ($ns^2np^6$, except He with 1s²), giving them exceptional chemical stability and very low reactivity — they exist as monoatomic gases and were long considered completely "inert."
Their fully-filled, highly stable electronic configuration gives extremely high ionization enthalpies and near-zero electron affinity, leaving essentially no thermodynamic driving force for forming ordinary chemical bonds.
Xenon, the largest easily available noble gas with the lowest ionization enthalpy in the group, can be made to react with highly electronegative elements like fluorine and oxygen under special conditions. Neil Bartlett first prepared a xenon compound (Xe⁺[PtF₆]⁻) in 1962, disproving the belief that noble gases were entirely unreactive. Xenon forms fluorides ($\text{XeF}_2, \text{XeF}_4, \text{XeF}_6$) by direct reaction with fluorine gas under heat/UV light and pressure, and oxides/oxyfluorides by subsequent hydrolysis of these fluorides.
Explain why xenon, and not helium or neon, is the noble gas most capable of forming compounds.
Xenon is a much larger atom than helium or neon, so its outermost electrons are farther from the nucleus and less tightly held, giving xenon a comparatively low ionization enthalpy for a noble gas. This makes it (relatively) easier to remove an electron from xenon or induce partial bonding with a highly electronegative element like fluorine — a feat essentially impossible for the small, tightly-bound electrons of helium and neon.