s, p, d and f block elements

In short

The long form of the periodic table is divided into four blocks by the sub-shell that receives the last electron: s (groups 1-2 and He), p (groups 13-18), d (groups 3-12, the transition elements) and f (the lanthanoids and actinoids). Block position fixes valence-shell configuration, and that in turn fixes oxidation states, bonding type, colour, magnetism, acid-base and redox behaviour of every element and its compounds. This is the largest single scoring area of B.Sc inorganic chemistry and of CSIR-UGC NET, GATE Chemistry (CY), IIT-JAM and the technical rounds of Indian chemical, metallurgical, fertiliser, pharma and atomic-energy employers. Industrially it is the chemistry behind chlor-alkali and soda ash, aluminium and steel, sulphuric acid and fertilisers, pigments and catalysts, and India's rare-earth and thorium programme built on Kerala and Odisha beach sands.

Explained

How the blocks are defined and filled

• Aufbau order by increasing (n + l), ties broken by lower n: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p. Hund's rule and the Pauli principle complete the assignment.
• s-block: ns¹ (group 1, alkali metals) and ns² (group 2, alkaline earth metals); He (1s²) is placed in group 18 for chemical reasons.
• p-block: ns² np¹-6, groups 13-18; includes most non-metals (H and He are s-block), all metalloids (B, Si, Ge, As, Sb, Te) and the poor metals.
• d-block: (n−1)d¹-10 ns¹-2 in four series - 3d (Sc-Zn), 4d (Y-Cd), 5d (La, Hf-Hg), 6d (Ac, Rf-Cn, incomplete).
• f-block (inner transition): (n−2)f¹-14 (n−1)d⁰-1 ns² - 4f lanthanoids Ce-Lu (La is 5d¹, counted with them) and 5f actinoids Th-Lr (Ac is 6d¹).
• Exceptions from exchange energy / half- and fully-filled stability: Cr 3d⁵4s¹, Cu 3d¹⁰4s¹, Mo 4d⁵5s¹, Ag 4d¹⁰5s¹, Au 5d¹⁰6s¹, Pd 4d¹⁰5s⁰; Gd 4f⁷5d¹6s², Cm 5f⁷6d¹7s².
• A transition element is defined as one with a partly filled d sub-shell in the atom or in a common oxidation state - so Zn, Cd, Hg (d¹⁰) are d-block but not typical transition metals: no variable valency, no colour, diamagnetic, poor catalysts.

s-block: alkali and alkaline earth metals

• Soft, low-melting, low-density (Li, Na, K float on water), strongly electropositive; lose ns electrons to give M⁺ (group 1) and M²⁺ (group 2) in almost all compounds - predominantly ionic.
• Reactivity down the group increases as ionisation enthalpy falls; group 2 is less reactive than group 1 because the second IE is large and the atoms are smaller.
• Reducing power: Li has the most negative E° (−3.04 V) despite the highest IE in the group, because of its very large hydration enthalpy; in melts (no water) Cs is the strongest reducer.
• Oxides: Li → Li₂O, Na → Na₂O₂ (peroxide), K/Rb/Cs → MO₂ (superoxide); all oxides/hydroxides basic, basicity increasing down the group. BeO is amphoteric.
• Thermal stability of carbonates, nitrates, hydroxides, sulphates increases down a group (larger cation polarises the anion less); Li₂CO₃, MgCO₃ and BeCO₃ decompose easily. LiNO₃ and Mg(NO₃)₂ give the oxide + NO₂ + O₂, while NaNO₃ gives the nitrite.
• Solubility in liquid ammonia: alkali metals give deep blue, paramagnetic, conducting solutions of ammoniated electrons [e(NH₃)ₓ]⁻, bronze and diamagnetic when concentrated - used as reducing media (Birch reduction).
• Flame colours (electron excited by low-energy flame): Li crimson, Na golden-yellow, K lilac, Rb red-violet, Cs blue; Ca brick-red, Sr crimson, Ba apple-green. Be and Mg give no colour (too high excitation energy).
• Anomalies of Li and Be: smallest size, highest charge density → covalent character (Fajans' rules), low-melting halides soluble in organic solvents, LiCl·2H₂O hygroscopic, Li is the only alkali metal forming a nitride (Li₃N) and a stable hydride-type carbide. Be has maximum covalence 4 (BeF₄²⁻), BeCl₂ is a polymeric chain in the solid and Be²⁺ salts crystallise as tetrahydrates containing [Be(H₂O)₄]²⁺ (e.g. BeSO₄·4H₂O), not the six-coordinate hydrates of Mg.
• Important compounds: NaOH (chlor-alkali / Castner-Kellner), Na₂CO₃·10H₂O washing soda (Solvay), NaHCO₃, NaCl, CaO and CaCO₃, Ca(OH)₂, Plaster of Paris CaSO₄·½H₂O from gypsum CaSO₄·2H₂O at 393 K, bleaching powder Ca(OCl)Cl, Mg(OH)₂ (milk of magnesia), CaC₂.
• Biological role: Na⁺/K⁺ pump and nerve conduction, Mg²⁺ in chlorophyll and in ATP-utilising enzymes, Ca²⁺ in bone (hydroxyapatite), clotting and muscle contraction.

p-block: groups 13 to 18

• Group 13: B is a hard, high-melting non-metal/metalloid; Al, Ga, In, Tl are metals. Electron-deficient hydrides: B₂H₆ has two 3-centre-2-electron "banana" bonds (B-H-B bridges), sp³ B, and is hydrolysed to H₃BO₃ + H₂. Borax Na₂[B₄O₅(OH)₄]·8H₂O (borax bead test for coloured metaborates); H₃BO₃ is a layered, weak MONObasic Lewis acid - it accepts OH⁻ rather than donating H⁺. Al is amphoteric (dissolves in NaOH giving aluminate), passivated by Al₂O₃; alums KAl(SO₄)₂·12H₂O; anhydrous AlCl₃ is a dimer Al₂Cl₆ and a Friedel-Crafts Lewis acid.
• Group 14: maximum catenation at carbon (C-C 348 kJ/mol) falling down the group. Allotropes: diamond (sp³, insulator, hardest), graphite (sp², layered, conducting, lubricant), fullerene C₆₀, graphene, carbon nanotubes. CO is a strong σ-donor/π-acceptor ligand and a poison (binds haemoglobin); CO₂ is linear and acidic. SiO₂ is a 3-D covalent network while CO₂ is a gas - the difference is Si's inability to form strong pπ-pπ bonds. Silicates, zeolite catalysts (ZSM-5), silicones (R₂SiO)ₙ, silicon carbide. Stability of +2 rises down the group: SnCl₂ is a reducing agent, PbCl₄ decomposes to PbCl₂, PbO₂ is a strong oxidant.
• Group 15: N₂ has a strong triple bond (941 kJ/mol) and hence is inert; N has no d orbitals, so no NCl₅ or NF₅ while P, As, Sb form pentahalides. NH₃ (Haber), HNO₃ (Ostwald), NO, N₂O, N₂O₅. Phosphorus allotropes: white (P₄ tetrahedra, reactive, poisonous, glows), red (polymeric chains, safer), black (most stable). PCl₅ is trigonal bipyramidal in the gas and ionic [PCl₄]⁺[PCl₆]⁻ in the solid. Basicity of oxoacids = number of P-OH groups: H₃PO₂ monobasic, H₃PO₃ dibasic, H₃PO₄ tribasic; acids with P-H bonds are reducing.
• Group 16: O₂ is paramagnetic (two unpaired electrons in π* by MOT); O₃ is bent, a powerful oxidant, measured in the Chappuis/Hartley bands. S has rhombic (α, stable below 369 K) and monoclinic (β) S₈ crowns; SO₂ (bleaching by reduction, reversible), H₂SO₄ by the contact process (V₂O₅, 2 SO₂ + O₂ ⇌ 2 SO₃, exothermic → moderate temperature and excess air by Le Chatelier). SF₆ octahedral and kinetically inert; SF₄ see-saw. Oxoacid strength H₂SO₄ > H₂SO₃; thiosulphate in photography and iodometry.
• Group 17: F₂ is the strongest oxidant (E° +2.87 V) despite a low electron affinity, because of its low bond dissociation energy and high hydration energy of F⁻. Bond energy order Cl₂ > Br₂ > F₂ (lone-pair repulsion in the small F₂). HF is a liquid (H-bonding) but the weakest hydrohalic acid; acid strength HF < HCl < HBr < HI. Interhalogens ICl, ICl₃, BrF₃, IF₅, IF₇ (larger halogen is central). Oxoacid acidity HOCl < HClO₂ < HClO₃ < HClO₄ (more O, more delocalised anion); HClO₄ is the strongest common mineral acid. Disproportionation of Cl₂ in alkali: cold dilute → Cl⁻ + ClO⁻, hot concentrated → Cl⁻ + ClO₃⁻.
• Group 18: chemically inert (filled ns²np⁶, very high IE). Only Xe (and Kr/Rn marginally) forms true compounds: XeF₂ linear, XeF₄ square planar, XeF₆ distorted octahedral, XeO₃ pyramidal, XeOF₄ square pyramidal, XeO₆⁴⁻ perxenate. Hydrolysis: XeF₆ + 3H₂O → XeO₃ + 6HF. Uses: He in cryogenics and balloons, Ne in signs, Ar as an inert welding/arc shield, Kr/Xe in lamps.
• Anomalous first member of every p-block group (B, C, N, O, F): small size, high electronegativity, absence of d orbitals, strong pπ-pπ multiple bonding, low maximum covalence of 4.

d-block: the transition elements

• General traits: hard, high-melting, high-boiling, good conductors (strong metallic bonding from unpaired d electrons); Zn, Cd, Hg are exceptions (Hg is liquid).
• Atomic/ionic radii fall across a series then level off; 4d and 5d of the same group are nearly equal in size (Zr/Hf, Nb/Ta, Mo/W) because of the lanthanoid contraction.
• Variable oxidation states differing by one, because (n−1)d and ns energies are close. Across 3d the maximum state rises to Mn (+7) then falls; the highest states occur in oxides and fluorides (MnO₄⁻, CrO₄²⁻, VO₄³⁻), low states in carbonyls (Ni(CO)₄, Fe(CO)₅, zero oxidation state).
• Stability rules: half-filled/filled configurations are stable, so Mn²⁺ (d⁵) and Fe³⁺ (d⁵) are favoured and Fe²⁺ is reducing; Cu⁺ disproportionates in water (2Cu⁺ → Cu + Cu²⁺) because Cu²⁺ has a far larger hydration enthalpy.
• Colour comes from d-d transitions (and charge transfer in d⁰ ions such as MnO₄⁻ and Cr₂O₇²⁻, which is ligand-to-metal charge transfer, not d-d); d⁰ and d¹⁰ ions (Sc³⁺, Ti⁴⁺, Zn²⁺, Cu⁺) are colourless.
• Magnetism: most ions are paramagnetic in proportion to unpaired electrons; Fe, Co, Ni are ferromagnetic.
• Catalysis: variable valency and surface adsorption - Fe/K₂O in Haber, V₂O₅ in contact process, Pt-Rh gauze in Ostwald, Ni in hydrogenation, TiCl₄/Al(C₂H₅)₃ Ziegler-Natta, MoS₂ in hydrodesulphurisation.
• Interstitial compounds (hydrides, carbides, nitrides, e.g. TiC, steel) - non-stoichiometric, hard, chemically inert, metallic conductivity; and alloys (stainless steel Fe-Cr-Ni, brass, bronze) because radii are similar.
• Key reagents: K₂Cr₂O₇ from chromite FeCr₂O₄ (roast with Na₂CO₃ + air → Na₂CrO₄, acidify → dichromate, then KCl); orange, oxidises Fe²⁺, I⁻, H₂S, SO₃²⁻, Sn²⁺ in acid, gaining 6 electrons. KMnO₄ from pyrolusite MnO₂ (fuse with KOH + air → manganate, then disproportionate or oxidise electrolytically); purple, 5 electrons in acid (→ Mn²⁺), 3 in neutral/alkaline (→ MnO₂), 1 in strong alkali (→ MnO₄²⁻). Chromate ⇌ dichromate is pH-dependent: 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O.

f-block: lanthanoids and actinoids

• Lanthanoids (Ce-Lu, with La): typical state +3; silvery, reactive, tarnish in air, react with water and acids. Anomalous states follow empty/half/full 4f: Ce⁴⁺ (4f⁰, strong oxidant used in cerimetry), Tb⁴⁺, Pr⁴⁺; Eu²⁺ (4f⁷, reducing), Yb²⁺, Sm²⁺.
• Lanthanoid contraction: steady decrease in size from La³⁺ (1.03 Å) to Lu³⁺ (0.86 Å) because 4f electrons shield poorly, so Zeff rises. Consequences: basicity of M(OH)₃ falls along the series (used in fractional precipitation and ion-exchange separation), 4d-5d radii match, and the chemistry of the 15 elements is so alike that separation needs ion-exchange or solvent extraction.
• Magnetic and spectral behaviour: 4f orbitals are buried, so f-f transitions give sharp, line-like absorptions little affected by ligands, and the magnetic moment needs L-S coupling, μ = g√(J(J+1)), not the spin-only formula (only La³⁺ and Lu³⁺ are diamagnetic).
• Actinoids (Th-Lr): 5f orbitals are less buried, so oxidation states range +3 to +7 (U: +3 to +6, UO₂²⁺ uranyl; Np +7; Pu +4). All radioactive; actinoid contraction is larger because of poorer 5f shielding; elements beyond U are man-made.
• Uses: misch metal (~95% lanthanoids + Fe) in lighter flints and as a steel scavenger; Nd-Fe-B and SmCo permanent magnets for EV motors and wind turbines; LaNi₅ hydrogen storage; Ce in catalytic converters and glass polishing; Eu/Tb/Y phosphors; Gd in MRI contrast; U-235 and Pu-239 fission fuel; Th-232 is India's three-stage programme fuel (fertile → U-233).

Anomalies, trends and the exam favourites

• Inert pair effect: the ns² pair becomes reluctant to bond down groups 13-16 (poor shielding by intervening d and f electrons), so lower oxidation states gain stability: Tl⁺ > Tl³⁺, Pb²⁺ > Pb⁴⁺, Bi³⁺ > Bi⁵⁺, Sn²⁺ reducing vs Pb⁴⁺ oxidising.
• Diagonal relationship: Li-Mg, Be-Al, B-Si behave alike because charge/radius ratio (polarising power) is similar - e.g. Li and Mg both form nitrides and decomposable carbonates; Be and Al both amphoteric with covalent chlorides.
• Fajans' rules: covalent character rises with small cation, large anion, high cation charge and a non-noble-gas (d¹⁰) cation shell - explains AgI covalent vs NaCl ionic, and the colour and low solubility of AgX.
• Hydration vs lattice enthalpy decides solubility: LiF is sparingly soluble (very high lattice enthalpy), CsF very soluble; among group 2 sulphates, solubility falls BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄.
• Screening order f < d < p < s for a given n; d and f electrons shield badly, which is the root of both the inert pair effect and lanthanoid contraction.

Indian extraction, regulation and practice

• Chlor-alkali: membrane-cell electrolysis of brine gives Cl₂, NaOH, H₂ (Grasim, DCM Shriram, GACL, Aditya Birla); mercury cells are being phased out under the Minamata Convention, which India ratified.
• Aluminium: bauxite → Bayer process (hot NaOH leach, red mud residue) → Hall-Héroult electrolysis of Al₂O₃ in molten cryolite with CaF₂, carbon anodes, about 950-980 °C (NALCO, Hindalco, Vedanta).
• Iron and steel: blast furnace reduction of Fe₂O₃ by CO with CaCO₃ flux; Ellingham diagrams justify the carbon reduction temperature (Tata Steel, SAIL, JSW).
• Zinc and lead: sulphide ores at Rampura Agucha and Zawar roasted and electrowon (Hindustan Zinc); copper by froth flotation, smelting and electrorefining.
• Rare earths and thorium: monazite (a Ce-La-Nd-Th phosphate, typically about 8-10% ThO₂ in Indian sands) from the beach sands of Chavara and Manavalakurichi (Kerala/Tamil Nadu) and OSCOM (Odisha), processed by Indian Rare Earths Ltd (IREL) under the Department of Atomic Energy; uranium from Jaduguda and Tummalapalle by UCIL.
• Regulation to quote: the SHANTI Act, 2025 replaced the Atomic Energy Act, 1962 (under which uranium, thorium and monazite were "prescribed substances"), keeps the nuclear fuel cycle under government control and makes AERB a statutory radiation regulator; the Mines and Minerals (Development and Regulation) Act governs atomic and beach-sand minerals; plant safety now sits under the Occupational Safety, Health and Working Conditions Code, 2020, in force from 21 November 2025 (replacing the Factories Act, 1948) together with the MSIHC Rules for hazardous chemicals; effluent and emission limits under the Water Act, 1974 and Air Act, 1981 are enforced by CPCB/SPCBs. BIS publishes specifications for caustic soda, soda ash, liquid chlorine and sulphuric acid - confirm the current IS number and edition from the BIS catalogue, as several older numbers have been revised or withdrawn.

Key points

• Block - identified by the sub-shell receiving the last electron: s, p, d or f.
• Group 1 vs 2 - ns¹ gives M⁺, ns² gives M²⁺; group 2 is harder, denser, higher melting and less reactive.
• Li anomaly - most negative E° in group 1 because of its very large hydration enthalpy, not low IE.
• Oxide trend - Li₂O (oxide), Na₂O₂ (peroxide), KO₂/RbO₂/CsO₂ (superoxides); larger cations stabilise bigger anions.
• BeO and Al₂O₃ - amphoteric; all other s-block oxides are basic.
• Diagonal relationship - Li-Mg, Be-Al, B-Si, from equal polarising power.
• Electron deficiency - B₂H₆ uses two 3c-2e B-H-B bonds; BF₃ and AlCl₃ are Lewis acids.
• H₃BO₃ - weak, monobasic Lewis acid; layered, H-bonded solid.
• Catenation - strongest at C, falls C > Si > Ge ≈ Sn > Pb.
• Inert pair effect - lower oxidation state more stable down groups 13-16: Tl⁺, Pb²⁺, Bi³⁺.
• Oxoacid basicity of P - equals the number of P-OH groups; P-H bonds make the acid reducing.
• F anomaly - strongest oxidant but weakest X-X bond and weakest hydroacid (HF).
• Noble gas compounds - only Xe forms a real series: XeF₂ linear, XeF₄ square planar, XeO₃ pyramidal.
• Transition element - partly filled d sub-shell in the atom or a common ion; Zn, Cd, Hg are excluded.
• Colour - d-d transitions, or LMCT in d⁰ oxoanions MnO₄⁻ and Cr₂O₇²⁻.
• Electron change - MnO₄⁻ accepts 5e⁻ in acid, 3e⁻ in neutral, 1e⁻ in strong alkali; Cr₂O₇²⁻ accepts 6e⁻.
• Lanthanoid contraction - 4f poor shielding → size falls La→Lu; makes Zr/Hf and Nb/Ta nearly identical and basicity of M(OH)₃ fall along the series.
• Lanthanoid vs actinoid - 4f buried (sharp f-f lines, +3 dominant) vs 5f exposed (many states, all radioactive, larger contraction).

Formulas and examples

1. Lattice enthalpy of NaCl by the Born-Haber cycle. Data (kJ mol⁻¹): ΔfH°(NaCl) = −411, ΔsubH(Na) = +108, IE(Na) = +496, D(Cl₂) = +242, ΔegH(Cl) = −349.

• Sum of steps other than lattice: 108 + 496 + (242/2 = 121) − 349 = +376.
• U = ΔfH − 376 = −411 − 376 = −787 kJ mol⁻¹ (released on forming the lattice).

2. Spin-only moments across 3d. Count unpaired d electrons, then μ = √(n(n+2)).

3. Why the spin-only formula fails for lanthanoids - Nd³⁺ (4f³, ground term ⁴I₉/₂). S = 3/2, L = 6, J = 9/2.

• J(J+1) = (4.5)(5.5) = 24.75; S(S+1) = (1.5)(2.5) = 3.75; L(L+1) = (6)(7) = 42.
• g = 1 + (24.75 + 3.75 − 42)/(2 × 24.75) = 1 + (−13.5/49.5) = 1 − 0.273 = 0.727.
• μ = 0.727 × √24.75 = 0.727 × 4.975 = 3.62 BM, against an observed ≈3.5 BM. The spin-only value would be √15 = 3.87 BM - wrong because orbital angular momentum is not quenched in buried 4f orbitals.

4. Energy to win 1 tonne of aluminium (Hall-Héroult). Al³⁺ + 3e⁻ → Al, so z = 3.

• Moles of Al = 1,000,000/26.98 = 37,064 mol → moles of electrons = 111,192 mol.
• Charge Q = 111,192 × 96,485 = 1.073 × 10¹⁰ C (≈2,980 kAh).
• At a cell voltage of 4.5 V: E = 1.073 × 10¹⁰ × 4.5 = 4.83 × 10¹⁰ J = 13,400 kWh per tonne (Indian smelters report roughly 13,000-15,000 kWh/t with current efficiency below 100%) - which is why aluminium plants sit beside captive power.

5. Dichromate titration of a ferrous salt. 20.0 mL of 0.02 M K₂Cr₂O₇ exactly oxidises 25.0 mL of Mohr's salt solution. Cr₂O₇²⁻ + 14H⁺ + 6Fe²⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O.

• Moles of Cr₂O₇²⁻ = 0.02 × 0.0200 = 4.0 × 10⁻⁴ mol.
• Moles of Fe²⁺ = 6 × 4.0 × 10⁻⁴ = 2.4 × 10⁻³ mol.
• Molarity of Fe²⁺ = 2.4 × 10⁻³ / 0.0250 = 0.096 M (equivalently 0.096 N, z = 1).

6. Zeff on the 3s electron of Na (Slater). Config 1s² 2s²2p⁶ 3s¹. Same shell: none. (n−1) = 8 electrons × 0.85 = 6.80. (n−2) = 2 × 1.00 = 2.00. S = 8.80 → Zeff = 11 − 8.80 = 2.20, which is why the 3s electron is lost so easily.

Where it is used

• Chlor-alkali, soda ash and sulphuric acid plants (GACL, Tata Chemicals, DCM Shriram, Coromandel): process chemists and shift engineers apply s- and p-block reaction and equilibrium chemistry daily.
• Metallurgy and mining (Tata Steel, SAIL, Hindalco, NALCO, Hindustan Zinc, Vedanta): roasting, leaching, Ellingham-based reduction, electrowinning and electrorefining - roles in process metallurgy and quality control.
• Catalysts and petrochemicals (IOCL/BPCL/HPCL R&D, Reliance, NCL Pune, CSIR labs): d-block oxides and zeolites for reforming, hydrotreating, polymerisation.
• Atomic energy and rare earths (IREL, UCIL, BARC, NFC Hyderabad, IGCAR): monazite processing, thorium and uranium chemistry, Nd/Sm magnet feedstock, solvent-extraction separation of lanthanoids.
• Fertilisers and agrochemicals (RCF, IFFCO, GSFC): ammonia, nitric acid, phosphate and potash chemistry - all p- and s-block.
• Analytical, pharma and QC labs (NABL-accredited, IS/IP/USP methods): KMnO₄ and K₂Cr₂O₇ titrimetry, AAS/ICP-OES for metal limits, heavy-metal testing under FSSAI and BIS water specifications.

Interview questions

• Q: Why is Li the strongest reducing agent in group 1 though Cs has the lowest ionisation enthalpy? A: E° depends on sublimation, ionisation and hydration; Li⁺ is very small so its hydration enthalpy is exceptionally large, making the overall process most favourable in aqueous solution.
• Q: Why does Na form a peroxide and K a superoxide, but Li only an oxide? A: Lattice stability needs cation-anion size matching; the small Li⁺ stabilises only the small O²⁻, while bigger K⁺ can stabilise the large O₂⁻.
• Q: Why is BeCl₂ a chain polymer in the solid but BeCl₂ vapour a monomer/dimer? A: Be is electron-deficient with only four valence orbitals, so in the solid chlorine bridges fill its octet; at high temperature the bridges break.
• Q: H₃BO₃ has three OH groups - why is it monobasic? A: It does not release H⁺; it accepts OH⁻ from water to form [B(OH)₄]⁻, releasing one H⁺ per molecule, so it acts as a weak monobasic Lewis acid.
• Q: Why does carbon form CO₂ gas while silicon forms solid SiO₂? A: Carbon forms strong pπ-pπ double bonds giving discrete O=C=O molecules; silicon cannot, so it builds a three-dimensional network of Si-O single bonds.
• Q: Explain the inert pair effect with an example. A: The ns² pair is poorly shielded by intervening d/f electrons and becomes hard to ionise, so PbCl₂ and Tl⁺ are stable while PbCl₄ and Tl³⁺ are strong oxidants.
• Q: Why are Zn, Cd and Hg not regarded as typical transition metals? A: Their d¹⁰ shell is complete in the atom and in the common +2 ion, so they lack variable valency, colour, paramagnetism and catalytic activity.
• Q: MnO₄⁻ and Cr₂O₇²⁻ have d⁰ metals - why are they intensely coloured? A: The colour arises from ligand-to-metal charge transfer, not d-d transitions.
• Q: Why do Zr and Hf have almost identical properties? A: The lanthanoid contraction cancels the expected size increase, so their atomic and ionic radii, and hence chemistry, are nearly the same.
• Q: Why is Ce⁴⁺ a good oxidising agent and Eu²⁺ a good reducing agent? A: Ce⁴⁺ reverts to the stable 4f¹ (Ce³⁺) by gaining an electron; Eu²⁺ (4f⁷) loses one to reach the common +3 state.
• Q: Why do lanthanoid ions give sharp absorption lines and need the g√(J(J+1)) formula? A: The 4f orbitals are shielded by 5s and 5p, so ligands barely perturb them and orbital angular momentum is not quenched.
• Q: Why does fluorine have a lower electron affinity than chlorine yet a higher oxidising power? A: Inter-electron repulsion in the tiny F atom lowers its electron affinity, but the weak F-F bond and the very high hydration enthalpy of F⁻ make the overall reduction highly favourable.

Common mistakes

• Writing d-block and transition element as identical - Zn, Cd and Hg fail the partly-filled-d test.
• Applying the spin-only formula to lanthanoids, or ignoring the orbital contribution for Co²⁺ and Ni²⁺.
• Sign errors in the Born-Haber cycle: electron gain enthalpy and lattice enthalpy are negative, ionisation enthalpy positive; also forgetting to halve the X₂ bond enthalpy.
• Counting all H atoms when giving the basicity of H₃PO₃ or H₃PO₂ - only P-OH hydrogens are ionisable.
• Saying HF is the strongest hydrohalic acid because it has the strongest hydrogen bonding - the strong H-F bond makes it the weakest.
• Assuming the highest oxidation state keeps rising across the 3d series; it peaks at Mn (+7) and then falls.
• Confusing the lanthanoid contraction (a size effect) with the inert pair effect (a valency effect), or using 5 electrons for KMnO₄ in an alkaline titration.

Related topics

• Periodic properties · Coordination chemistry · Organometallic chemistry basics · Chemical bonding and molecular structure (VSEPR, MOT) · Metallurgy and extraction of metals · Redox titrations and volumetric analysis · Nuclear and radiochemistry.