Periodic Properties, s- and p-Block Elements — NEET Chemistry

Inorganic chemistry is where the periodic table comes alive. The s-block gives you the reactive metals of groups 1 and 2; the p-block spans the whole range from active metals to metalloids to the noble gases. NEET pulls 3–4 questions from here, and — crucially — they are fast, high-accuracy recall marks if you have read NCERT and understood the trends and anomalies. This chapter organises the representative elements group by group, always explaining a property from atomic size, effective nuclear charge and the inert-pair effect, and flags every exception the exam loves. It is dense on facts by design — that density is the syllabus here.

Part A — The s-Block

1. Group 1 — the alkali metals (Li, Na, K, Rb, Cs)

• Configuration $n{s}^1$; lose one electron easily → +1 ions only. Softest, lightest, most reactive metals.
• Down the group: atomic size increases, ionisation enthalpy decreases, reactivity and metallic character increase. They are the strongest reducing agents (Li is anomalously the strongest in aqueous solution due to its very high hydration energy).
• Reactivity with water: all react to give hydroxide + H₂, increasingly vigorously down the group (K and below ignite the hydrogen).
• Flame colours (a direct NEET recall): Li crimson, Na golden yellow, K lilac/violet, Rb red, Cs blue.
• Solutions in liquid ammonia are deep blue (ammoniated electrons), paramagnetic and conducting.

Anomalous lithium (small size, high polarising power): forms mainly the normal oxide Li₂O (Na → peroxide Na₂O₂, K → superoxide KO₂); LiF and Li₂CO₃ are less soluble; Li is more covalent. Diagonal relationship Li–Mg: both form normal oxides, nitrides directly with N₂, and covalent, hydrated compounds.

Worked example 1.1. Which alkali metal imparts a golden-yellow flame colour? Sodium. The 589 nm sodium D-line is the golden-yellow of street lamps — a guaranteed recall question.

2. Group 2 — the alkaline-earth metals (Be, Mg, Ca, Sr, Ba)

• Configuration $n{s}^2$; form +2 ions. Harder, denser and less reactive than group 1; smaller for the same period.
• Down the group: size increases, IE decreases, reactivity increases; hydroxide solubility increases [Mg(OH)₂ sparingly soluble, Ba(OH)₂ soluble] while sulphate solubility decreases (BaSO₄ insoluble — used in "barium meal" X-rays).
• Flame colours: Ca brick-red, Sr crimson, Ba apple-green (Be and Mg give no colour — electrons too tightly held).

Anomalous beryllium (very small, high charge density): covalent compounds, amphoteric oxide/hydroxide (BeO reacts with both acid and base), does not impart flame colour, coordination number limited to 4. Diagonal relationship Be–Al: both amphoteric, both form covalent halides that dimerise/polymerise, both have a protective oxide layer.

Worked example 2.1. Why is BeO amphoteric while MgO is basic? Beryllium's small size and high charge density give BeO significant covalent character, letting it react with both acids and bases (amphoteric). Larger Mg²⁺ makes MgO ionic and purely basic. This mirrors the Be–Al diagonal relationship.

Part B — The p-Block

3. General p-block trends and the inert-pair effect

The p-block (groups 13–18) fills the $np$ sub-shell. Two overarching ideas:

• Metallic character increases down each group and non-metallic decreases; the "staircase" of metalloids (B, Si, Ge, As, Sb, Te) divides them.
• Inert-pair effect: down groups 13–15 the lower oxidation state (group number − 2) becomes more stable because the $n{s}^2$ pair is increasingly reluctant to participate. So Tl⁺ > Tl³⁺, Pb²⁺ > Pb⁴⁺, Bi³⁺ > Bi⁵⁺ in stability — heavier elements favour the lower state, making the higher state a strong oxidiser (PbO₂, Bi⁵⁺).

The second-period elements (B, C, N, O, F) are anomalous — small size, no $d$ orbitals, high electronegativity — so they show maximum covalency 4 and strong multiple (pπ–pπ) bonding, unlike their heavier congeners.

4. Group 13 — the boron family (B, Al, Ga, In, Tl)

• Boron is a non-metallic metalloid; the rest are metals. Configuration $n{s}^{2}n{p}^1$; typical oxidation state +3 (but +1 grows down via inert-pair, dominant in Tl).
• Electron deficiency: BF₃ has only 6 electrons on B → a strong Lewis acid (accepts a lone pair). Boron halide Lewis acidity order: BF₃ < BCl₃ < BBr₃ — counter-intuitive, because back-bonding from halogen to boron is strongest (and most deactivating) in BF₃.
• Diborane B₂H₆ is electron-deficient with 3-centre-2-electron "banana" bonds (two bridging H); it does not have a normal B–B bond.
• Aluminium: amphoteric oxide (Al₂O₃), reacts with both acids and NaOH; forms a protective oxide layer (passivation); AlCl₃ is a covalent Lewis-acid catalyst (Friedel–Crafts).

Worked example 4.1. Why is BF₃ a Lewis acid? Boron in BF₃ has only 6 valence electrons (an incomplete octet), so it readily accepts an electron pair from a donor like NH₃ to complete its octet — the definition of a Lewis acid.

5. Group 14 — the carbon family (C, Si, Ge, Sn, Pb)

• Configuration $n{s}^{2}n{p}^2$; oxidation states +4 and +2 (+2 grows down; Pb²⁺ more stable than Pb⁴⁺, so PbO₂ and Pb⁴⁺ are oxidisers — inert-pair effect).
• Catenation (self-linking) is strongest for carbon (strong C–C bonds) and falls sharply down the group (Si–Si weaker), which is why carbon builds the vast field of organic chemistry.
• Allotropes of carbon: diamond (sp³, hardest, insulator), graphite (sp², layers, conductor and lubricant), fullerene (C₆₀), graphene.
• Oxides: CO₂ is a discrete linear molecule (gas); SiO₂ is a giant covalent network (solid). CO is neutral and a strong ligand/reducing agent; CO₂ is acidic.

Worked example 5.1. Why is CO₂ a gas but SiO₂ a high-melting solid, though both are "MO₂"? Carbon forms strong pπ–pπ double bonds to oxygen, giving discrete O=C=O molecules held only by weak forces → gas. Silicon is too large for effective pπ–pπ bonding, so SiO₂ is a giant covalent network of Si–O single bonds → a hard, high-melting solid.

6. Group 15 — the nitrogen family (N, P, As, Sb, Bi)

• Configuration $n{s}^{2}n{p}^3$ (half-filled, stable); oxidation states −3 to +5; +3 stability grows down (Bi³⁺ dominant).
• Nitrogen is a diatomic gas with a very strong triple bond (bond order 3) → inert; it cannot expand its octet (no $d$ orbitals), so NCl₃ exists but NCl₅ does not, while PCl₅ does.
• Ammonia (NH₃): pyramidal, Lewis base/ligand, made industrially by the Haber process; HNO₃ is a strong oxidising acid.
• Allotropy of phosphorus: white (P₄, reactive, poisonous, glows/stored under water), red (polymeric, stable), black (most stable).
• Down-group basicity of hydrides decreases: NH₃ > PH₃ > AsH₃ > SbH₃ (larger atom, less available lone pair); thermal stability also decreases.

Worked example 6.1. Why does nitrogen not form NCl₅ while phosphorus forms PCl₅? Nitrogen (period 2) has no vacant $d$ orbitals, so it cannot expand its octet beyond 4 bonds. Phosphorus (period 3) has accessible 3d orbitals, allowing $s{p}^{3}d$ hybridisation and 5 bonds in PCl₅.

7. Group 16 — the oxygen family / chalcogens (O, S, Se, Te)

• Configuration $n{s}^{2}n{p}^4$; gain 2 electrons → −2; also +4, +6 (with O).
• Oxygen is a diatomic gas (paramagnetic, see MOT); ozone O₃ is a bent, powerful oxidising allotrope that shields UV in the stratosphere.
• Sulphur shows catenation (S₈ rings) and multiple oxidation states. SO₂ is acidic, a reducing agent and bleach; H₂SO₄ ("king of chemicals") is a strong dibasic acid, dehydrating agent and oxidiser.
• Oxoacid strength of sulphur: more oxygen/higher oxidation state → stronger acid (H₂SO₄ > H₂SO₃).
• Hydride trends: thermal stability decreases H₂O > H₂S > H₂Se > H₂Te; acidic strength increases down (H₂O < H₂S < H₂Se < H₂Te) as the H–X bond weakens; water's high boiling point is the H-bonding anomaly.

Worked example 7.1. Which is more acidic, H₂S or H₂O, and why? H₂S. Down group 16 the H–X bond lengthens and weakens, so the proton is released more easily — acidity increases down the group despite oxygen's higher electronegativity. (Boiling point, however, is highest for H₂O because of hydrogen bonding.)

8. Group 17 — the halogens (F, Cl, Br, I)

• Configuration $n{s}^{2}n{p}^5$; gain one electron → −1; the most reactive non-metals. F₂ is the strongest oxidising agent among common elements.
• Oxidising power decreases down: F₂ > Cl₂ > Br₂ > I₂ (so a halogen displaces those below it: Cl₂ displaces Br⁻ and I⁻).
• Bond enthalpy anomaly: F–F is weaker than Cl–Cl (small F atoms, lone-pair repulsion), which partly explains fluorine's extreme reactivity.
• Hydrohalic acid strength increases down: HF < HCl < HBr < HI (weaker H–X bond down the group). HF is a weak acid (strong H-bonding) but etches glass.
• Interhalogens (ClF₃, IF₇) and oxoacids (HOCl, HClO₄) exist; oxoacid strength rises with oxidation state (HClO₄ > HClO₃ > HClO₂ > HOCl).
• Bleaching action of Cl₂ (via HOCl) is by oxidation and is permanent, unlike SO₂'s reductive, temporary bleaching.

Worked example 8.1. Arrange F₂, Cl₂, Br₂, I₂ in order of oxidising power. F₂ > Cl₂ > Br₂ > I₂. Oxidising power decreases down the group (electrode potential falls), so fluorine is the strongest and iodine the weakest — chlorine can displace bromide and iodide but not fluoride.

9. Group 18 — the noble gases (He, Ne, Ar, Kr, Xe, Rn)

• Configuration $n{s}^{2}n{p}^6$ (fully filled, He is $1s^2$) → chemically inert, monatomic, very high ionisation enthalpies.
• Xenon (largest, lowest IE among stable ones) forms real compounds with F and O: XeF₂ (linear), XeF₄ (square planar), XeF₆, and XeO₃. These validated the VSEPR/hybridisation models.
• Uses: helium (balloons, cryogenics, breathing mixtures), neon (signs), argon (inert welding/lamp atmosphere).

Worked example 9.1. What is the shape of XeF₄? Xe has 4 bond pairs + 2 lone pairs = 6 electron pairs (octahedral geometry); the lone pairs sit opposite each other, giving a square planar shape (sp³d² hybridisation).

10. Common traps NEET sets here

• Flame colours: Na golden-yellow, K lilac, Li crimson, Ca brick-red, Ba apple-green (Be, Mg none).
• Li, Be anomalies and Li–Mg, Be–Al diagonal relationships — recur every year.
• Inert-pair effect: lower oxidation state stabilises down groups 13–15 (Pb²⁺, Tl⁺, Bi³⁺); higher state oxidises.
• BF₃ < BCl₃ < BBr₃ Lewis acidity — opposite to naive electronegativity reasoning (back-bonding).
• N has no NCl₅ (no d orbitals); P forms PCl₅.
• CO₂ gas vs SiO₂ solid — pπ–pπ bonding available only to carbon.
• Halogen oxidising power F₂ > Cl₂ > Br₂ > I₂, but acid strength HF < HCl < HBr < HI.
• F–F bond is weaker than Cl–Cl — small-atom lone-pair repulsion.
• Down group 16/17 acidity increases, thermal stability decreases for hydrides.

11. Memory aids

• "Na yellow, K lilac, Li crimson, Ca brick, Ba green" — flame-test recall.
• "Small Li and Be break the rules; diagonal partners Mg and Al" — the anomalies.
• "Inert pair loves the low state" — Pb²⁺, Tl⁺, Bi³⁺ down the p-block.
• "F beats Cl beats Br beats I" — oxidising power (but acid strength reverses).
• "Carbon catenates, silicon can't network its way out" — catenation and CO₂ vs SiO₂.
• "No d, no PCl₅ for nitrogen" — second-period octet limit.
• "Xenon fluorides: 2 linear, 4 square-planar, 6" — noble-gas compounds.

12. Exam protocol

1. Reason every property from size, $Z_{\text{eff}}$ and the inert-pair effect — don't just memorise.
2. Recall flame colours and the Li/Be anomalies + diagonal relationships cold.
3. Group 13: electron deficiency → Lewis acidity (BF₃), diborane's banana bonds, Al amphoterism.
4. Group 14: catenation falls down; CO₂ molecular vs SiO₂ network; Pb²⁺ from inert pair.
5. Group 15: N₂ triple-bond inertness, no NCl₅; hydride basicity/stability fall down.
6. Group 16: ozone, SO₂/H₂SO₄, hydride acidity increases down, stability decreases.
7. Group 17: oxidising power F₂ > I₂; acid strength HF < HI; oxoacid strength rises with oxidation state.
8. Group 18: xenon fluorides and their VSEPR shapes.