Periodic trends and s-block/p-block chemistry make up one of the highest-yield areas across NEET, JEE, IIT-JAM, BITSAT, GATE, CSIR-NET, and the TGT/PGT screening exams. I've organized this article the way I'd want it organized if I were revising the night before an exam — trend first, exception right next to it, and a worked example wherever the logic needs reinforcing. Nothing here is padded; every line either builds a concept or tests it.
- General Periodic Trends — Group and Period
- Metallic Character, Reducing Nature, and Electropositivity
- Basicity of Hydrides, Oxides, and Oxyacids
- Thermal Stability of Salts
- Acidic Strength of Oxides and Oxyacids
- Screening Effect, Effective Nuclear Charge, and Hydride Stability
- S-Block Trends: Hydrides and Halides
- Melting Point, Hydration, and Conductivity in S-Block
- Reactions of Alkali and Alkaline Earth Metals
- P-Block General Trends
- Group 13 — Boron Family
- Important Compounds of Boron
- Group 14 — Carbon Family
- Group 15 — Nitrogen Family
- Bond Angle Trends
- Fajan's Rule
- Oxides and Oxoacids of Nitrogen
- Oxides and Oxoacids of Phosphorus, and Phosphorus Allotropes
- Group 16 — Oxygen Family and Oxoacids of Sulphur
- Group 17 — Halogens
- Interhalogen Compounds
- Group 18 — Noble Gases and Xenon Chemistry
- Solved Examples
- Practice MCQs with Answers
1. General Periodic Trends — Group and Period
Start with the two directions separately, because students routinely mix them up under exam pressure. Going down a group, atomic size increases — extra shells are being added, and the added electrons sit further from the nucleus than the pull of the increased nuclear charge can compensate for. Going left to right across a period, the added electrons fill the same shell while the nuclear charge climbs steadily, so the electron cloud gets pulled in tighter and atomic size shrinks.
Ionisation potential, electron gain enthalpy (electron affinity), and electronegativity all move opposite to atomic size in a period — they increase left to right — while down a group they decrease, since the valence electrons are progressively easier to remove or less strongly attracted to a distant nucleus.
2. Metallic Character, Reducing Nature, and Electropositivity
Metallic character, reducing nature, and electropositivity increase down a group and decrease across a period — they track atomic size and inversely track ionisation energy. Non-metallic character and oxidising nature do the opposite: they increase across a period and decrease down a group.
3. Basicity of Hydrides, Oxides, and Oxyacids
Basic character of metal hydrides, basic character of oxides, and basic character of oxyacids all increase down a group and decrease across a period. This single statement quietly answers a huge fraction of "which is more basic" questions in inorganic chemistry, so it's worth memorising as one block rather than three separate facts.
4. Thermal Stability of Salts
For sulphates, carbonates, nitrates, and hydroxides, thermal stability increases down a group and decreases across a period. The reasoning traces back to lattice energy and polarising power: a smaller, more highly charged cation at the top of a group polarises the large anion more strongly, weakening the anion's internal bonds and making the salt decompose at a lower temperature. As the cation grows down the group, its polarising power drops, and the salt survives higher temperatures before it breaks down.
5. Acidic Strength of Oxides and Oxyacids
The strength of oxyacids and the acidic character of oxides both increase across a period and decrease down a group — the mirror image of basicity. A textbook example worth memorising directly:
Na₂O < MgO < Al₂O₃ < SiO₂ < P₄O₁₀ < SO₃ < Cl₂O₇
Notice how this single row moves cleanly from strongly basic (Na₂O) through amphoteric (Al₂O₃) to strongly acidic (Cl₂O₇), and it's a favourite fill-in-the-blank or ordering question.
6. Screening Effect, Effective Nuclear Charge, and Hydride Stability
Screening effect stays roughly constant across a period (inner shells don't change) but increases down a group as more inner shells accumulate. Effective nuclear charge (Zeff) increases both down a group and across a period, though the period-wise increase is much sharper because screening barely changes there while actual nuclear charge keeps climbing.
Hydride thermal stability behaves differently for the two blocks, and this is a classic trap. For s-block hydrides, stability decreases going down the group. For p-block hydrides, it's the opposite — stability decreases going up and increases going down, which is really just saying stability again falls as you move from the lighter to the heavier p-block hydrides in most series, so always check which block you're dealing with before applying the rule.
7. S-Block Trends: Hydrides and Halides
Group 1 hydrides follow the MH pattern and Group 2 hydrides follow MH₂. Ease of hydride formation decreases as you go down each group:
BeH₂ > MgH₂ > CaH₂ > SrH₂ > BaH₂
BeH₂ deserves special attention because it doesn't behave like a simple ionic hydride — it adopts a polymeric structure held together by three-centre-two-electron (3c–2e) bonds, the same "banana bonding" idea you'll meet again in diborane. Each bridging hydrogen sits between two beryllium centres, and this electron-deficient bonding is what makes BeH₂ chemistry distinct from the rest of the alkaline earth hydrides.
For halides, Group 1 and Group 2 follow MX and MX₂ respectively. Alkali metal halides mostly adopt the 6,6-coordinate rock-salt structure, like NaCl — except CsCl, CsBr, and CsI, which switch to 8,8-coordination because the larger Cs⁺ ion changes the favoured packing geometry.
Among the alkaline earth metals, solubility of the halides is inversely proportional to atomic number:
8. Melting Point, Hydration, and Conductivity in S-Block
Melting and boiling points of the alkali metals fall steadily down the group: Li > Na > K > Rb > Cs. For the alkaline earth metals the order isn't perfectly monotonic — melting point follows Be > Ca > Sr > Ba > Mg, while boiling point follows Be > Ba > Ca > Sr > Mg. These irregular orders show up because of the particular crystal packing each metal adopts, and I'd suggest just memorising them as given rather than trying to force a clean trend onto them.
Hydration energy is proportional to charge density, which is inversely proportional to cation size — smaller, more charge-dense ions hold water molecules more tightly.
Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺
Ionic mobility and conductivity behave differently depending on whether you're in aqueous solution or the gas phase, and this is a genuinely counter-intuitive point. In aqueous solution, conductivity is proportional to cation size — Cs⁺ > Rb⁺ > K⁺ > Na⁺ > Li⁺ — because the smaller ions carry a thicker hydration shell that slows them down even though the bare ion is small. In the gaseous phase, where there's no hydration shell to drag along, conductivity is instead inversely proportional to cation size: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺.
Solubility itself is proportional to the ratio of hydration energy to lattice energy. When the anion is small — F⁻ or OH⁻ — lattice energy dominates the trend and solubility increases down the group:
Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂
When the anion is large — SO₄²⁻, CO₃²⁻, ClO₄⁻ — hydration energy dominates instead, and solubility now decreases down the group:
BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄
9. Reactions of Alkali and Alkaline Earth Metals
An alkali metal M reacts with water to give MOH, with ammonia to give MNH₂, and with a halogen X₂ to give MX. Its reaction with O₂ is where the real distinction lies: lithium and sodium form simple oxides (Li₂O, Na₂O), lithium and sodium also form peroxides (Li₂O₂, Na₂O₂), while potassium, rubidium, and caesium — the heavier members — form superoxides (KO₂, RbO₂, CsO₂). The stability of the larger anion (peroxide, then superoxide) increases as the cation gets larger and can better stabilise that bigger, more diffuse anion through lattice energy.
Alkaline earth metals show a parallel pattern: reaction with water gives M(OH)₂, with ammonia gives M(NH₂)₂, with a halogen gives MX₂, and with H₂ gives MH₂ (only BeH₂ and MgH₂ form this way, and BeH₂ again shows that polymeric structure). For the oxygen reaction, beryllium, magnesium, and calcium form simple oxides, strontium sits in between, and barium is the one that readily forms the peroxide BaO₂.
Nearly all the oxides of these elements adopt the rock-salt structure — the one deliberate exception is BeO, which, consistent with beryllium's generally anomalous behaviour in this group, does not.
10. P-Block General Trends
Basic character of hydrides of non-metals increases down a group and decreases across a period — mirroring the s-block hydride basicity rule, just applied to non-metal hydrides now. Acidic character of these same hydrides increases across a period.
CH₄ < NH₃ < H₂O < HF (acidic strength across period 2)
For oxides, a high oxidation state on the central atom correlates with higher acidity:
HClO₄ > HClO₃ > HClO₂ > HClO (oxidation states +7, +5, +3, +1)
Across period 2 and period 3, the oxides themselves split cleanly into basic, amphoteric, and acidic zones: the metals on the left (Li, Be→partially, Na, Mg, K, Ca, Ga, Sr) tend basic, the metalloids in the middle (B, Al, Si-adjacent elements like Ga, In, Sn, Pb, Bi, Po) tend amphoteric, and the non-metals on the right (C, N, O, F, Si, P, S, Cl, As, Se, Br, Te, I, At) tend acidic — with Al₂O₃, Ga₂O₃, SnO₂, and PbO all being classic amphoteric examples worth remembering by name.
Oxidation state stability follows a clean rule down any p-block group: the higher oxidation state becomes progressively less stable as you move down the group, while the lower oxidation state becomes progressively more stable — this is the inert pair effect showing up as a trend. Negative oxidation states, meanwhile, simply become less stable moving down the group.
11. Group 13 — Boron Family
| Property | Trend |
|---|---|
| Atomic radius | B < Ga < Al < In < Tl |
| Ionisation energy | B > Tl > Ga > Al > In |
| Electronegativity | B > Tl > In > Ga > Al |
| Boiling point | B > Al > Ga > In > Tl |
| Melting point | B > Al > Tl > In > Ga |
Notice the atomic radius order isn't a clean monotonic climb — gallium actually sits smaller than aluminium because of poor d-orbital shielding after the first transition series, one of the more commonly tested anomalies in this group.
Two oxidation states matter here: +3 and +1. Stability of the +3 state falls down the group — B³⁺ > Al³⁺ > Ga³⁺ > In³⁺ > Tl³⁺ — while stability of the +1 state rises down the group: B⁺ < Al⁺ < Ga⁺ < In⁺ < Tl⁺. This is the inert pair effect in its most frequently tested form; thallium is famous for preferring the +1 state, which is why TlI₃ simply does not exist as a genuine Tl(III) triiodide (it's actually a Tl(I) triiodide salt, Tl⁺I₃⁻). Along the same logic, GaCl₃ is more stable than TlCl₃ because Ga³⁺ is more stable than Tl³⁺.
Lewis acid character of the boron trihalides runs BBr₃ > BCl₃ > BF₃ — the reverse of what you'd first guess from electronegativity alone. The accepted explanation is back-bonding: fluorine's small size lets its lone pair donate into boron's empty p-orbital most effectively, which fills the empty orbital boron needs for accepting a new lone pair and so weakens BF₃'s Lewis acidity relative to BCl₃ and BBr₃, where back-bonding is progressively weaker.
Comparing the whole family's trichlorides: BCl₃ > AlCl₃ > GaCl₃ > InCl₃ > TlCl₃ in Lewis acid strength, since boron has no d-orbitals or larger empty orbitals to fall back on and depends entirely on the small, empty 2p orbital that back-bonding from Cl can't fully quench.
12. Important Compounds of Boron
Borax
Borax has the formula Na₂B₄O₇·10H₂O, more accurately written as Na₂[B₄O₅(OH)₄]·8H₂O to reflect its real structure. The borax bead test identifies coloured metal ions — Ni²⁺, Co²⁺, Cr³⁺, Cu²⁺, Mn²⁺ — by forming characteristic coloured metaborates:
CoO + B₂O₃ → Co(BO₂)₂ (blue bead)
Diborane (B₂H₆)
Diborane's structure is the standard teaching example of electron-deficient bonding: four terminal B–H bonds and two bridging B–H–B bonds, with each bridge built from a three-centre-two-electron bond — the banana bond. Its reactions are worth knowing as a set:
B₂H₆ + 6H₂O → 2H₃BO₃ + 6H₂
B₂H₆ + 2CO → 2BH₃·CO
With ammonia, the product depends sharply on conditions. Excess NH₃ at low temperature gives the simple adduct B₂H₆·2NH₃. Excess NH₃ at high temperature gives boron nitride, (BN)ₓ. A 2:1 ratio of NH₃ to B₂H₆ at high temperature gives borazine, B₃N₃H₆ — sometimes nicknamed "inorganic benzene" because of its isoelectronic six-membered ring, though that comparison is really only structural and shouldn't be pushed too far in an exam answer.
Orthoboric Acid (H₃BO₃)
This is a weak, monobasic acid, but the mechanism behind its acidity is the part examiners actually probe: H₃BO₃ is not a protonic acid. It behaves as a Lewis acid by accepting an electron pair from a hydroxide ion of water rather than by donating a proton itself:
The structure is a planar, layered arrangement held together by hydrogen bonds. On heating, H₃BO₃ loses water stepwise, first to metaboric acid and then to boron trioxide:
13. Group 14 — Carbon Family
| Property | Trend |
|---|---|
| Ionisation energy | C > Si > Ge > Pb > Sn |
| Electronegativity | C > Pb > Si ≈ Ge ≈ Sn |
| Boiling point | Si > Ge > Sn > Pb |
| Melting point | C > Si > Ge > Sn > Pb |
Three oxidation states are relevant: +4, +2, and −4. Stability of +4 falls down the group — C⁴⁺ > Si⁴⁺ > Ge⁴⁺ > Sn⁴⁺ > Pb⁴⁺ — while +2 becomes more stable down the group, again the inert pair effect at work: C²⁺ < Si²⁺ < Ge²⁺ < Sn²⁺ < Pb²⁺.
Catenation ability — the tendency to form chains and rings of like atoms bonded to each other — falls sharply down the group: C >> Si > Ge ≈ Sn, reflecting how carbon's small size and strong C–C bonds are simply not matched by its heavier congeners.
Tin is famous for existing in two allotropic forms in equilibrium: grey tin (α-Sn), which has a diamond-type structure and is the low-temperature form, and white tin (β-Sn), which is metallic and stable at ordinary temperature.
14. Group 15 — Nitrogen Family
| Property | Trend |
|---|---|
| Ionisation energy and electronegativity | N > P > As > Sb > Bi |
| Boiling point | Sb > Bi > As > P > N |
| Melting point | As > Sb > Bi > P > N |
Three oxidation states show up: +5, +3, and −3. Stability of +5 falls down the group, but with a twist — Bi⁵⁺ is actually the most unstable rather than following a smooth descent, so the order is As⁵⁺ > Sb⁵⁺ > Bi⁵⁺, with Bi⁵⁺ specifically flagged as the most unstable member (a strong oxidiser, in fact) rather than the term "most stable" applying anywhere in this row. Stability of +3 rises down the group: As³⁺ < Sb³⁺ < Bi³⁺. For the −3 state, stability falls down the group: N⁻³ > P⁻³ > As⁻³ > Sb⁻³ > Bi⁻³.
Covalent character in the pentahalides exceeds that in the trihalides for a given halogen — EX₅ > EX₃ — because the higher oxidation state pulls electron density more strongly. All the EX₃ trihalides are covalent except BiF₃, which is ionic, consistent with bismuth's heavier, more metallic character. Among the trifluorides, NF₃ is the only one that's genuinely stable under ordinary conditions. Catenation ability follows P > N, since phosphorus can form longer chains more readily than the smaller nitrogen atom despite nitrogen's stronger individual bonds.
Hydride boiling points follow BiH₃ > SbH₃ > NH₃ > AsH₃ > PH₃ — note that ammonia jumps ahead of arsine and phosphine because of hydrogen bonding, even though it's much lighter than either. Melting points follow NH₃ > SbH₃ > AsH₃ > PH₃, again with ammonia elevated by hydrogen bonding.
15. Bond Angle Trends
This section carries a lot of exam weight, so it's worth working through each rule with its example rather than just memorising the ordering.
(i) Down a group: as you move down, atomic size increases and lone pair–bond pair repulsion decreases, so bond angle steadily falls: NH₃ > PH₃ > AsH₃ > SbH₃.
(ii) Bond angle is proportional to s-character of the hybrid orbitals involved: sp > sp² > sp³, since more s-character pulls bonding electron pairs closer to the nucleus and further from each other angularly, widening the angle.
(iii) Central atom electronegativity: when hybridisation, lone pairs, and bond pairs are all the same, bond angle is proportional to the electronegativity of the central atom — a more electronegative central atom draws bonding pairs closer to itself, reducing repulsion between them and widening the angle: NCl₃ > PCl₃ > AsCl₃ > SbCl₃.
(iv) Substituent effects: when the central atom, hybridisation, and lone pair count are all held the same, more electronegative substituents pull bonding electron density away from the central atom, shrinking the angle, while larger substituents create more steric repulsion and widen it. Compare Cl₂O > H₂O > F₂O — fluorine's high electronegativity pulls density away most strongly, shrinking F₂O's angle below even water's, while the bulkier, less electronegative chlorine in Cl₂O gives the widest angle of the three.
16. Fajan's Rule
Fajan's rule explains why no ionic compound is ever truly 100% ionic — even a highly ionic salt like CsF carries some 9–10% covalent character. According to the rule, covalent character is favoured by: a small cation, a large anion, high charge on either ion, and a cation with an 18-electron outer shell rather than the noble-gas 8-electron configuration.
Covalent character ∝ anion size: LiI > LiBr > LiCl > LiF
Covalent character ∝ charge on cation/anion: SnCl₂ < SnCl₄; UF₄ < UF₆; SbCl₃ < SbCl₅; CaCl₂ > NaCl; SCl₃ > CaCl₂ > KCl
18-electron shell cations are more covalent: CuCl > NaCl, since Cu⁺ (3s²3p⁶3d¹⁰) has an 18-electron shell while Na⁺ has the ordinary 8-electron noble-gas shell
17. Oxides and Oxoacids of Nitrogen
Nitrogen gas exists as two stable isotopes, ¹⁴N and ¹⁵N, and can be generated in the lab from Ba(N₃)₂ or 2NaN₃, or through NH₄Cl + NaNO₂, or by heating ammonium dichromate, (NH₄)₂Cr₂O₇. Its reactivity, despite the strong N≡N triple bond, includes direct combination with reactive metals and with hydrogen and oxygen under forcing conditions:
N₂ + 3H₂ →(773K) 2NH₃ N₂ + O₂ →(Δ) 2NO
Oxides of Nitrogen
| Name | Formula | O.S. of N | Appearance / nature |
|---|---|---|---|
| Nitrogen(I) oxide, dinitrogen oxide | N₂O | +1 | Colourless gas, neutral |
| Nitrogen(II) oxide | NO | +2 | Colourless gas, neutral |
| Nitrogen(III) oxide, dinitrogen trioxide | N₂O₃ (planar) | +3 | Blue solid, acidic |
| Nitrogen(IV) oxide, nitrogen dioxide | NO₂ | +4 | Brown gas, acidic |
| Nitrogen(IV) oxide, dinitrogen tetroxide | N₂O₄ (planar) | +4 | Colourless solid/liquid, acidic |
| Nitrogen(V) oxide, dinitrogen pentoxide | N₂O₅ (planar) | +5 | Colourless solid, acidic |
Oxoacids of Nitrogen
Nitric acid (HNO₃) is both a strong mineral acid and a strong oxidising acid, with nitrogen at +5 in the nitrate ion NO₃⁻. Industrially it's made by the Ostwald process:
2NO + O₂ → 2NO₂ 3NO₂ + H₂O → 2HNO₃ + NO
It also forms from NaNO₃ + H₂SO₄. In water it behaves as a strong acid, fully ionising: HNO₃ + H₂O → H₃O⁺ + NO₃⁻. Gold and platinum resist it entirely — no reaction — while chromium and aluminium famously do not dissolve in HNO₃ either, because both form a passivating oxide layer on the surface. As an oxidiser, HNO₃ attacks non-metals directly:
C + 4HNO₃ → CO₂ + 2H₂O + 4NO₂
S₈ + 48HNO₃ → 8H₂SO₄ + 48NO₂ + 12H₂O
P₄ + 20HNO₃ → 4H₃PO₄ + 20NO₂ + 4H₂O (phosphoric acid)
Hyponitrous acid (H₂N₂O₂) has nitrogen at +1 and functions mainly as a reducing agent, forming the hyponitrite ion N₂O₂²⁻.
Nitrous acid (HNO₂) has nitrogen at +3, is weak and monobasic, and is unusual in being able to act as both an oxidising agent and a reducing agent depending on what it's paired with — this dual behaviour is a frequently asked conceptual point. Acidic strength across these three follows HNO₃ > HNO₂ > H₂N₂O₂.
18. Oxides and Oxoacids of Phosphorus, and Phosphorus Allotropes
Phosphorus, arsenic, and antimony each form two series of oxides at +3 and +5: P₄O₆ and P₄O₁₀, As₄O₆ and As₄O₁₀, Sb₄O₆ and Sb₄O₁₀. Bismuth, being the heaviest and most metallic, forms only Bi₂O₃ at +3 as its characteristic oxide. In P₄O₁₀'s cage structure, six oxygen atoms bridge the four phosphorus atoms while four terminal oxygens double-bond directly to each phosphorus. Acidity across the row falls down the group: N₂O₃ > P₄O₆ > As₄O₆ > Sb₄O₆ > Bi₂O₃.
Phosphorus allotropes
White phosphorus is chemiluminescent, reacts with hot concentrated NaOH to release phosphine, and burns in oxygen to white fumes of P₄O₁₀:
P₄ + 5O₂ → P₄O₁₀
Heating white phosphorus in an inert atmosphere at 573K converts it to red phosphorus, and red phosphorus converted in a closed tube at 803K gives black phosphorus. White phosphorus can also convert directly to β-black phosphorus under high pressure at 473K; black phosphorus itself exists as α- and β-forms.
Oxoacids of phosphorus
| Acid | Formula (O.S.) | Basicity / character |
|---|---|---|
| Orthophosphoric acid | H₃PO₄ (+5) | Tribasic; P₄O₁₀ + 6H₂O → 4H₃PO₄ |
| Hypophosphoric acid | H₄P₂O₆ (+4) | Tetrabasic, resists oxidation; salts are hypophosphates |
| Pyrophosphoric acid | H₄P₂O₇ (+5) | Tetrabasic |
| Metaphosphoric acid | (HPO₃)ₙ (+5) | Exists as a trimer, tribasic when n = 3 |
| Orthophosphorous acid | H₃PO₃ (+3) | Dibasic, one P–H bond acts as a strong reducing hydrogen; salts are phosphites |
| Metaphosphorous acid | (HPO₂)ₙ | May adopt a ring structure |
| Hypophosphorous acid | H₃PO₂ (+1) | Monobasic, strong reducing agent; salts are hypophosphites |
| Pyrophosphorous acid | H₄P₂O₅ | Dibasic, reducing |
19. Group 16 — Oxygen Family and Oxoacids of Sulphur
Boiling and melting points run O < S < Se < Te > Po — tellurium actually exceeds polonium, breaking the simple monotonic pattern you'd expect. Electron affinity runs S > Se > Te > Po > O, with oxygen's affinity being the smallest of the group despite being the lightest, because of strong electron-electron repulsion packed into oxygen's small 2p subshell. Hydride boiling points run H₂O > H₂Te > H₂Se > H₂S, with water elevated far above the trend by hydrogen bonding.
Oxidation states of +6, +4, and −2 all occur across the group, with oxygen itself being an exception — it shows −2 as by far its most common state, since oxygen (unlike its heavier congeners) essentially never reaches +4 or +6 in ordinary compounds; sulphur, selenium, and tellurium do reach +4 and +6, particularly in oxides and oxoacids, while polonium's chemistry is dominated by +2.
Among the halides, stability of the hexahalides is highest for fluorine — SF₆ is the only EF₆ that's genuinely stable, existing as an inert, kinetically stabilised gas, while SeF₆ and TeF₆ (liquid and solid respectively at comparable conditions) are progressively less robust. Monohalides of this group are dimeric in nature: S₂F₂, S₂Cl₂, S₂Br₂, Se₂Cl₂, Se₂Br₂.
Oxoacids of sulphur
| Series | Acid | O.S. | Property |
|---|---|---|---|
| Sulphurous | Sulphurous acid, H₂SO₃ | +4 | Basic, reducing agent |
| Pyrosulphurous acid, H₂S₂O₅ | +5 and +3 | Never isolated pure; its salts (disulphites) exist | |
| Dithionous acid, H₂S₂O₄ | +3 | Strong, facile reducing agent | |
| Sulphuric | Sulphuric acid, H₂SO₄ | +6 | Strong acid, undergoes autoprotolysis: 2H₂SO₄ ⇌ H₃SO₄⁺ + HSO₄⁻ |
| Thiosulphuric acid, H₂S₂O₃ | +6, −2 | Moderately strong reducing agent; the acid decomposes but its salts (thiosulphates) are stable | |
| Pyrosulphuric acid (oleum), H₂S₂O₇ | +6 | Formed in the contact process | |
| Thionic | Dithionic acid, H₂S₂O₆ | +5 | Resists both oxidation and reduction |
| Polythionic acid, H₂SₙO₆ (n = 1–12) | +5 and 0 | Can be prepared with up to six sulphur atoms in the chain | |
| Peroxo | Peroxomonosulphuric acid (Caro's acid), H₂SO₅ | +6 | Peroxodisulphate salts are strong oxidising agents |
| Peroxodisulphuric acid (Marshall's acid), H₂S₂O₈ | +6 | — |
The contact process for manufacturing H₂SO₄ runs through catalytic oxidation of SO₂ to SO₃ over V₂O₅, absorption into oleum, and dilution:
Concentrated H₂SO₄ acts as a strong dehydrating agent, pulling water directly out of organic compounds rather than merely absorbing free moisture — its action on sucrose is the classic demonstration: C₁₂H₂₂O₁₁ → 12C + 11H₂O. Hot concentrated H₂SO₄ is also a moderately strong oxidising agent; both metals and non-metals get oxidised by it, with the acid itself being reduced to SO₂ in the process:
S + 2H₂SO₄(conc.) → 3SO₂ + 2H₂O
C + 2H₂SO₄(conc.) → CO₂ + 2SO₂ + 2H₂O
20. Group 17 — Halogens
Electron affinity runs Cl₂ > F₂ > Br₂ > I₂ — chlorine, not fluorine, has the highest electron affinity in the group, because fluorine's very small atomic size packs the incoming electron into a crowded 2p subshell with significant electron-electron repulsion. Bond dissociation energy runs Cl₂ > Br₂ > F₂ > I₂, with the F–F bond unusually weak for the same small-size, high-repulsion reason.
Despite its lower bond dissociation energy, fluorine remains the most reactive and strongest oxidiser of the group: F₂ > Cl₂ > Br₂ > I₂, because the weak F–F bond breaks so easily that fluorine's overall reactivity outpaces the trend you'd expect from bond strength alone. Hydride melting points follow HI > HF > HBr > HCl, while boiling points follow HF > HI > HBr > HCl — in both cases HF sits well above where its molecular weight alone would place it, again because of hydrogen bonding.
Fluorine shows only the −1 oxidation state, since it's the most electronegative element and has no higher-oxidation-state chemistry available to it. Chlorine, bromine, and iodine show −1, +1, +3, +5, and +7, and chlorine and bromine additionally reach +4 and +6 in certain oxides and oxoacids.
Fluorine's own oxides, O₂F₂ and OF₂, are stable at 298K and both function as powerful fluorinating agents; O₂F₂ in particular is strong enough to oxidise plutonium metal all the way to PuF₆. Oxide stability across the halogens runs I > Cl > Br.
Manufacture and chemistry of chlorine
Deacon's process oxidises HCl gas with atmospheric oxygen over a CuCl₂ catalyst at 723K: 4HCl + O₂ →(CuCl₂) 2Cl₂ + 2H₂O. The electrolytic process liberates Cl₂ at the anode during electrolysis of concentrated brine, and chlorine is also recovered as a by-product in several industrial processes.
Chlorine reacts directly with metals and non-metals to give chlorides (2Al + 3Cl₂ → 2AlCl₃; P₄ + 6Cl₂ → 4PCl₃; S₈ + 4Cl₂ → 4S₂Cl₂), and it has a strong affinity for hydrogen, reacting with hydrogen-containing compounds to release HCl (H₂S + Cl₂ → 2HCl + S). With excess ammonia it gives nitrogen and ammonium chloride, but with excess chlorine the product switches to the explosive nitrogen trichloride:
NH₃ + 3Cl₂ → NCl₃ + 3HCl (excess Cl₂)
With cold, dilute NaOH, chlorine disproportionates to chloride and hypochlorite: Cl₂ + 2NaOH(cold, dilute) → NaCl + NaOCl + H₂O. With hot, concentrated alkali it goes further to chlorate: Cl₂ + 6NaOH(hot, conc.) → 5NaCl + NaClO₃ + 3H₂O. Chlorine water fades on standing because it slowly forms HCl and HOCl, and the HOCl releases nascent oxygen — the actual source of chlorine's oxidising and bleaching action.
Oxoacids of halogens
| Series | O.S. | Shape | F | Cl | Br | I |
|---|---|---|---|---|---|---|
| HOX (hypohalous) | +1 | Linear | HOF | HOCl | HOBr | HOI |
| HOXO (halous) | +2 | Angular | — | HOClO | — | — |
| HOXO₂ (halic) | +5 | Pyramidal | — | HOClO₂ | HOBrO₂ | HOIO₂ |
| HOXO₃ (perhalic) | +7 | Tetrahedral | — | HOClO₃ | HOBrO₃ | HOIO₃ |
Perchloric acid, HClO₄, is the strongest acid among all the oxoacids of the halogens. Acidic strength runs HOF > HOCl > HOBr > HOI within the +1 series, and HClO₄ > HBrO₄ > HIO₄ across the perhalic acids; within chlorine's own series it's HClO₄ (+7) > HClO₃ (+5) > HClO₂ (+3) > HClO (+1) — every oxoacid of the halogens is a good oxidising agent.
21. Interhalogen Compounds
Interhalogens follow four general types — XX′, XX′₃, XX′₅, XX′₇ — where X is the larger halogen and X′ the smaller one. Interhalogens are more reactive than the parent halogens themselves: ICl > IBr > Cl₂ in reactivity, and within the fluorides ClF < ClF₃ < ClF₅ < ClF₇ as more fluorines are packed on.
| Type | Examples | Structure |
|---|---|---|
| XX′ | ClF, BrF, IF (spectroscopic only), BrCl, ICl, IBr | — |
| XX′₃ | ClF₃, BrF₃, IF₃, ICl₃ (dimerises as I₂Cl₆) | Bent, T-shaped |
| XX′₅ | IF₅, BrF₅, ClF₅ | Square pyramidal |
| XX′₇ | IF₇ | Pentagonal bipyramidal |
These compounds are sparingly soluble in water and have unusually low melting and boiling points, since the only intermolecular interaction available to them is weak dispersion force. They also diffuse readily through common laboratory materials such as rubber, glass, and plastic — a property that shows up in questions about handling and storage precautions.
22. Group 18 — Noble Gases and Xenon Chemistry
Electron affinity across the noble gases runs Ne > Ar > Kr > Xe > Rn > He, and helium notably has the lowest boiling point of any known substance at 4.2K.
Noble gas chemistry effectively begins with Neil Bartlett's 1962 preparation of a xenon compound, achieved by reacting Xe with PtF₆: Xe + PtF₆ → Xe⁺[PtF₆]⁻ (a red-coloured solid). Among the other noble gases, krypton forms only KrF₂ and nothing else of note, radon's chemistry remains essentially unexplored, and argon, neon, and helium have no confirmed true compounds to date.
Xenon–fluorine compounds
Xenon forms three binary fluorides by direct reaction of the elements under controlled conditions:
Xe(g) + 2F₂(g) →(873K, 7 bar, 1:5 ratio) XeF₄(s)
Xe(g) + 3F₂(g) →(573K, 60–70 bar, 1:20 ratio) XeF₆(s)
XeF₆ can alternatively be prepared from XeF₄ and O₂F₂ at 143K: XeF₄ + O₂F₂ → XeF₆ + O₂. All three fluorides are colourless crystalline solids that sublime readily at 298K and act as powerful fluorinating agents. XeF₂ is linear, XeF₄ is square planar, and XeF₆ is a distorted octahedron.
Hydrolysis behaviour differs sharply between them. XeF₂ hydrolyses to elemental xenon: 2XeF₂ + 2H₂O → 2Xe + 4HF + O₂. XeF₄ or XeF₆ undergo complete hydrolysis to XeO₃, while XeF₆ alone can also undergo partial hydrolysis to give XeOF₄ and XeO₂F₂ instead of going all the way to XeO₃.
Xenon–oxygen compounds
| Formula | O.S. | Structure | e⁻ pairs | Lone pairs |
|---|---|---|---|---|
| XeO₃ | +6 | Pyramidal | 7 | 1 |
| XeOF₄ | +6 | Square pyramidal | 7 | 1 |
| XeO₄ | +8 | Tetrahedral | 8 | 0 |
| XeO₂F₂ | — | Trigonal bipyramidal | 7 | 1 |
| Ba₂[XeO₆]⁴⁻ (perxenate) | — | Octahedral | 8 | 0 |
XeO₃'s pyramidal shape comes from three Xe=O π bonds leaving four electron pairs to arrange tetrahedrally with one corner occupied by the lone pair. XeOF₄ has one π bond, leaving six pairs to form an octahedron with a lone pair in one position. XeO₄, with four π bonds and zero lone pairs, is a clean tetrahedron.
Complete hydrolysis of XeF₄ or XeF₆ gives XeO₃ + HF, while partial hydrolysis of XeF₆ specifically gives XeOF₄ + XeO₂F₂ + HF. XeO₃ is a powerful oxidising agent in acidic solution. In basic aqueous solution, the Xe(VI) oxoanion HXeO₄⁻ slowly decomposes and disproportionates:
23. Solved Examples
Solution: Moving across period 3, non-metallic character and hence acidity climbs steadily: Na₂O < MgO < Al₂O₃ < SiO₂ < P₄O₁₀ < SO₃ < Cl₂O₇.
Solution: Applying Fajan's rule, Be²⁺ is exceptionally small with a very high charge density, which polarises the chloride ion strongly enough to force significant electron sharing — that's covalent character. Ca²⁺ is much larger, polarises Cl⁻ far less, and CaCl₂ stays predominantly ionic.
Solution: Fluorine's small size allows efficient pπ–pπ back-donation from its filled p-orbital into boron's empty p-orbital, partly satisfying boron's electron deficiency before any external Lewis base arrives. Chlorine's larger, more diffuse orbitals back-bond far less efficiently, leaving BCl₃'s boron centre more electron-hungry and hence a stronger Lewis acid.
Solution: Complete hydrolysis replaces every fluorine with oxygen-containing groups, driving the reaction all the way to XeO₃. Partial hydrolysis — limited water — replaces only some of the fluorines, stopping at the intermediate oxyfluorides XeOF₄ and XeO₂F₂ before the reaction can go further.
24. Practice MCQs with Answers
1. Which of the following shows the correct order of increasing atomic radius?
(a) Ga < Al < In < Tl (b) B < Ga < Al < In < Tl (c) Al < Ga < In < Tl (d) B < Al < Ga < InAnswer: (b)
2. TlI₃ does not exist as a genuine Tl(III) compound because:
(a) Iodine cannot bond to thallium (b) The inert pair effect stabilises Tl⁺ over Tl³⁺, so it exists as Tl⁺I₃⁻ (c) TlI₃ is radioactive (d) Iodine is too electronegativeAnswer: (b)
3. Lewis acid strength of boron trihalides follows:
(a) BF₃ > BCl₃ > BBr₃ (b) BBr₃ > BCl₃ > BF₃ (c) BCl₃ > BF₃ > BBr₃ (d) All equalAnswer: (b)
4. The bridging bonds in diborane are best described as:
(a) Normal covalent bonds (b) Ionic bonds (c) Three-centre-two-electron bonds (d) Coordinate bonds onlyAnswer: (c)
5. Orthoboric acid acts as a Lewis acid because it:
(a) Donates a proton directly (b) Accepts OH⁻ from water, releasing H⁺ (c) Is a strong base (d) Has no acidic characterAnswer: (b)
6. Which does NOT exist?
(a) SnCl₄ (b) PbCl₄ (c) PbI₄ (d) SnI₄Answer: (c)
7. Catenation tendency in Group 14 follows:
(a) Pb > Sn > Ge > Si > C (b) C >> Si > Ge ≈ Sn (c) Si > C > Ge > Sn (d) All equalAnswer: (b)
8. Grey tin and white tin are related as:
(a) Isomers (b) Isotopes (c) Allotropes in equilibrium (d) Different compoundsAnswer: (c)
9. The multiple bonding in N₂ involves:
(a) dπ–dπ overlap only (b) One σ and two pπ–pπ bonds (c) Three σ bonds (d) One σ and two dπ–pπ bondsAnswer: (b)
10. Among P, As, Sb, and Bi, multiple bond character (where present) typically involves:
(a) pπ–pπ overlap (b) dπ–dπ overlap (c) No multiple bonding possible (d) sp³ overlapAnswer: (b)
11. Which trihalide of nitrogen family is ionic rather than covalent?
(a) NF₃ (b) PF₃ (c) BiF₃ (d) AsF₃Answer: (c)
12. Bond angle order in NH₃, PH₃, AsH₃, SbH₃ is:
(a) SbH₃ > AsH₃ > PH₃ > NH₃ (b) NH₃ > PH₃ > AsH₃ > SbH₃ (c) All equal (d) PH₃ > NH₃ > AsH₃ > SbH₃Answer: (b)
13. Bond angle is directly proportional to:
(a) p-character (b) s-character of the hybrid orbital (c) Lone pair count (d) Atomic massAnswer: (b)
14. Correct bond angle order: Cl₂O, H₂O, F₂O
(a) F₂O > H₂O > Cl₂O (b) Cl₂O > H₂O > F₂O (c) H₂O > Cl₂O > F₂O (d) All equalAnswer: (b)
15. According to Fajan's rule, covalent character increases with:
(a) Large cation, small anion (b) Small cation, large anion (c) Low charge on ions (d) Noble-gas type cation onlyAnswer: (b)
16. Which shows the highest covalent character?
(a) NaCl (b) CuCl (c) KCl (d) CsClAnswer: (b)
17. Oxidation state of nitrogen in N₂O₃ is:
(a) +2 (b) +3 (c) +4 (d) +5Answer: (b)
18. NO is an important intermediate in the industrial manufacture of:
(a) H₂SO₄ (b) HNO₃ (c) HCl (d) NH₃Answer: (b)
19. Which nitrogen oxoacid can act as both oxidising and reducing agent?
(a) HNO₃ (b) HNO₂ (c) H₂N₂O₂ (d) N₂O₅Answer: (b)
20. Gold and platinum do not dissolve in HNO₃ because:
(a) They are too reactive (b) They resist oxidation by HNO₃ under ordinary conditions (c) They form soluble nitrates instantly (d) HNO₃ is too dilute alwaysAnswer: (b)
21. Hypophosphorous acid, H₃PO₂, is:
(a) Tribasic (b) Dibasic (c) Monobasic and a strong reducing agent (d) Non-reducingAnswer: (c)
22. Reducing power of phosphorus oxoacids follows:
(a) H₃PO₄ > H₃PO₃ > H₃PO₂ (b) H₃PO₂ > H₃PO₃ > H₃PO₄ (c) All equal (d) H₃PO₃ > H₃PO₄ > H₃PO₂Answer: (b)
23. White phosphorus converts to red phosphorus on heating at:
(a) 473K under pressure (b) 573K in inert atmosphere (c) 803K in a closed tube (d) 298K in airAnswer: (b)
24. Which noble gas fluoride is NOT known?
(a) XeF₂ (b) XeF₄ (c) ArF₂ (d) XeF₆Answer: (c)
25. The shape of XeF₄ is:
(a) Tetrahedral (b) Square planar (c) Pyramidal (d) OctahedralAnswer: (b)
26. Complete hydrolysis of XeF₆ gives:
(a) XeO₂F₂ (b) XeOF₄ (c) XeO₃ (d) XeO₄Answer: (c)
27. Which pair is isostructural?
(a) XeF₄ and ICl₄⁻ (b) XeO₃ and ClO₃⁻ (c) XeF₂ and BrF₂⁻ (d) XeO₄ and IO₄⁻Answer: (a)
28. First noble gas compound was prepared by:
(a) Ramsay (b) Bartlett (c) Moissan (d) RutherfordAnswer: (b)
29. Electron affinity order of Group 17 is:
(a) F > Cl > Br > I (b) Cl > F > Br > I (c) I > Br > Cl > F (d) Br > Cl > F > IAnswer: (b)
30. Strongest oxidising halogen despite its weak bond dissociation energy is:
(a) Cl₂ (b) F₂ (c) Br₂ (d) I₂Answer: (b)
31. Interhalogens are generally __ reactive than the parent halogens.
(a) Less (b) More (c) Equally (d) UnpredictablyAnswer: (b)
32. Shape of ClF₃ is:
(a) Trigonal planar (b) Bent T-shaped (c) Tetrahedral (d) LinearAnswer: (b)
33. Which is the strongest acid among the oxoacids of halogens?
(a) HClO (b) HClO₃ (c) HClO₄ (d) HOIAnswer: (c)
34. Deacon's process converts HCl to Cl₂ using a catalyst of:
(a) V₂O₅ (b) CuCl₂ (c) MnO₂ (d) Fe₂O₃Answer: (b)
35. Cold dilute NaOH with Cl₂ gives:
(a) NaCl + NaClO₃ (b) NaCl + NaOCl (c) Only NaCl (d) NaClO₄Answer: (b)
36. Conc. H₂SO₄ acting on sugar demonstrates its role as a:
(a) Oxidising agent (b) Dehydrating agent (c) Reducing agent (d) CatalystAnswer: (b)
37. In the contact process, SO₂ is oxidised to SO₃ using catalyst:
(a) Pt (b) V₂O₅ (c) Fe (d) NiAnswer: (b)
38. Which sulphur oxoacid is known as Caro's acid?
(a) H₂S₂O₈ (b) H₂SO₅ (c) H₂S₂O₆ (d) H₂SO₃Answer: (b)
39. BeCl₂ in the vapour phase exists predominantly as:
(a) Polymer only (b) Monomer or dimer (c) Ionic lattice (d) TetramerAnswer: (b)
40. Solubility of alkaline earth sulphates (large anion case) follows:
(a) BeSO₄ > MgSO₄ > CaSO₄ > SrSO₄ > BaSO₄ (b) BaSO₄ > SrSO₄ > CaSO₄ > MgSO₄ > BeSO₄ (c) All equally soluble (d) CaSO₄ is most solubleAnswer: (a)
