The p-Block Giants: Group 15 & Group 16 Elements — A Complete Deep-Dive
Based on Inorganic Chemistry by Housecroft & Sharpe (4th Edition) | JEE Adv NEET GATE IIT-JAM CSIR-NET BITSAT · TGT · PGT
📋 Contents at a Glance
- Group 15 — The Pnictogens: Introduction & General Trends
- Occurrence, Extraction & Uses
- Physical Properties & Bonding Concepts
- Allotropy — Nitrogen, Phosphorus, As/Sb/Bi
- Hydrides: NH₃, PH₃, N₂H₄, HN₃ and Relatives
- Nitrides, Phosphides & Heavier Pnictides
- Halides & Oxohalides
- Oxides of Nitrogen
- Oxoacids: Nitrous, Nitric, Phosphoric & Beyond
- Phosphazenes — Inorganic Rubber
- Group 16 — The Chalcogens: Introduction & General Trends
- Occurrence, Extraction & Uses
- Physical Properties & Bonding
- The Elements: Allotropy of Sulfur, O₂, O₃, Se, Te
- Hydrides: H₂O₂, H₂S, H₂Se, H₂Te
- Halides — Sulfur Fluorides, Chlorides, SeF₄, TeF₆
- Oxides: SO₂, SO₃, SeO₂, TeO₂
- Oxoacids of Sulfur — H₂SO₄ and the Contact Process
- Exam Tips, Tricks & Quick Recall
Part 1 — Group 15: The Pnictogens
1.1 Introduction and General Character
The group 15 family — nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi) — is collectively called the pnictogens (from the Greek pniktos, "suffocating," owing to the choking properties of N₂ in enclosed spaces). The group spans an enormous range: from the lightest, most electronegative N, a non-metal whose triple bond in N₂ is among the strongest single bonds in chemistry, down to Bi, which behaves almost entirely as a soft heavy metal forming Bi³⁺ in solution.
A defining feature of this group is that simple monatomic ions are rare. The N³⁻ and P³⁻ ions are electrostatically improbable (highly endothermic to form) and most chemistry is covalent. The only confirmed simple cation in solution is Bi³⁺. Nitrogen alone spans oxidation states from −3 (NH₃) to +5 (HNO₃), producing an extraordinary diversity of compounds.
1.2 Electronic Configuration and Oxidation States
The ground state configurations follow ns²np³ (half-filled p subshell), giving special stability to the trivalent state:
Key properties of Group 15 elements. Note the small IE₁ drop from P→As (similar to the Al→Ga anomaly, due to poor d-orbital shielding).
A crucial trend: ionisation energies show a sharp jump after removal of the three 2p electrons in N (or equivalent np electrons in heavier members), reflecting the stability of the filled 2s (or ns) shell. The fourth IE is dramatically higher than the third, explaining why the +5 state is harder to achieve as you go down the group (Bi⁵⁺ is strongly oxidising and requires highly electronegative fluorine to stabilise it as BiF₅).
Part 2 — Occurrence, Extraction and Uses
2.1 Nitrogen — The Atmospheric Giant
Dinitrogen N₂ comprises 78% by volume of Earth's dry atmosphere. Its extraordinary triple bond (bond enthalpy 946 kJ mol⁻¹) makes it kinetically inert to most reactions at room temperature — a remarkable paradox, since its thermodynamic reactivity is substantial. Industrial nitrogen is obtained by fractional distillation of liquid air (bp 77 K). The key application of liquid N₂ as a low-temperature coolant is exploited in hundreds of laboratory procedures, including NMR, biological sample preservation and cryogenic surgery.
Nitrogen fixation — the conversion of N₂ into usable nitrogen compounds — is one of the most important chemical processes on Earth. It occurs biologically (nitrogenase enzymes in root-nodule bacteria of leguminous plants), naturally via lightning, and industrially through the Haber-Bosch process. The only natural non-atmospheric nitrogen source of industrial significance is sodium nitrate (NaNO₃, Chile saltpetre) from South American desert deposits.
2.2 Phosphorus — In Bones, DNA and Matches
Phosphorus does not occur free in nature. It is found as apatites Ca₅X(PO₄)₃ (where X = F in fluorapatite, Cl in chlorapatite, OH in hydroxyapatite). Hydroxyapatite is the principal mineral of bones and teeth. The global mining of phosphate rock (>160 million tonnes per year) primarily feeds fertiliser production.
Industrial extraction of elemental phosphorus uses a high-temperature electric furnace reduction:
Phosphorus vapour condenses under water to yield white phosphorus (P₄). Sand (SiO₂) acts as a slag-former by combining with calcium from the phosphate, while carbon (coke) acts as the reductant. The phosphorus vapour is collected and condensed below water to prevent spontaneous oxidation.
2.3 Arsenic, Antimony and Bismuth
The principal sources and extraction methods are:
| Element | Main Ore | Extraction Route |
|---|---|---|
| As | Mispickel (FeAsS), realgar (As₄S₄), orpiment (As₂S₃) | FeAsS →heat FeS + As↑ (sublimate condensed) |
| Sb | Stibnite (Sb₂S₃) | Sb₂S₃ + 3 Fe → 2 Sb + 3 FeS |
| Bi | Bismuthinite (Bi₂S₃), bismite (Bi₂O₃), native Bi | Reduction with C; also by-product of Pb, Cu, Sn refining |
Arsenic is notoriously toxic — a lethal dose is about 130 mg — and interacts with sulfur-containing proteins (thiol groups in enzymes). GaAs is an important III-V semiconductor used in transistors, LEDs and solar cells because of its high electron mobility (8500 cm² V⁻¹ s⁻¹, vs. 1500 for Si).
Part 3 — Physical Properties and Bonding
3.1 Bond Enthalpy Anomalies
The bond enthalpies in Table 15.3 of Housecroft & Sharpe reveal two critical anomalies that all competitive exam students must master:
Fig. 1 — Comparing triple vs. single bond enthalpies in N, P and C. The N≡N triple bond is anomalously strong; the N–N single bond is anomalously weak.
Two crucial observations:
Why is the N–N single bond (160 kJ mol⁻¹) so weak? Each nitrogen in N₂H₄ carries a lone pair. When two N atoms are singly bonded (as in hydrazine), the lone pairs on adjacent atoms experience strong repulsion because the atoms are small and the lone pairs are close together. This inter-lone-pair repulsion (illustrated in Housecroft Fig. 15.2) dramatically weakens the N–N, O–O and F–F single bonds compared to their heavier analogues (P–P, S–S, Cl–Cl) where the larger atoms hold lone pairs further apart.
Why does phosphorus form P₄ molecules rather than P₂? The N≡N triple bond is thermodynamically stable (positive ΔfH° for N₂H₄), meaning N₂ is more stable than P₄ analogues. For phosphorus, the strong σ-bond in P–P outweighs the weaker pπ–pπ overlap (because P uses 3p orbitals which spread out and overlap less effectively than compact 2p orbitals of N). So P₄ (with four P–P single bonds in a tetrahedral cage) is thermodynamically more stable than P₂.
3.2 Hypervalency: Why NF₅ Does Not Exist
Phosphorus forms PCl₅ and [PF₆]⁻ while nitrogen does not form NF₅ or [NF₆]⁻. The historical explanation invoked d-orbital participation (sp³d hybridisation) for P but not N. The modern view is more nuanced: it is simply that the N atom is too small to accommodate five substituents around it. Even with electronegative fluorine that would lower d-orbital energy, the d-orbitals play no significant role in hypervalent bonding. The correct description uses a 3-centre 4-electron (3c-4e) bond model, or recognises that the bonding in hypervalent molecules can be explained within a valence set of ns and np orbitals alone (as described in Housecroft Chapter 5).
Part 4 — The Elements: Allotropy and Reactivity
4.1 Nitrogen — Industrial Separation and Lab Preparations
Industrial N₂ comes from fractional distillation of liquid air (containing ~1% Ar and traces of O₂). Oxygen is removed by adding a little H₂ and passing over a Pt catalyst. An increasingly used alternative is gas-permeable membrane separation — cheaper but gives less pure N₂ (0.5–5% O₂). This is suited for inert blanket atmospheres for fruit and vegetable storage or electronics manufacture.
Laboratory preparations of small amounts of N₂:
N₂ is generally unreactive. It reacts with Li at room temperature, with heated metals (Ca, Ba, Al, Mg, many d-block metals), and with Ca₂C in the industrial production of calcium cyanamide. Atomic nitrogen (from electric discharge) is far more reactive than molecular N₂ and reacts even with S and Hg. The reaction of N₂ with H₂ (Haber process) and its coordination chemistry (N₂ complexes of d-block metals) are discussed below.
4.2 Phosphorus — The Most Complex Allotropy in the p-Block
Fig. 2 — Three key allotropes of phosphorus: (a) tetrahedral P₄ in white phosphorus; (b) infinite chain structure in red phosphorus; (c) layered puckered ring structure of black phosphorus (isostructural with As, Sb, Bi).
White phosphorus (P₄): Metastable standard state (ΔfH° = 0 by convention). Contains P₄ tetrahedra with strained 60° bond angles. Highly toxic; stored under water; ignites above 323 K. Soluble in CS₂ and benzene. The chemiluminescent oxidation in moist air ("Will-o'-the-Wisp") involves production of P₄O₈ as an intermediate.
Red phosphorus: Made by heating white P in inert atmosphere at ~540 K. Hittorf's (violet) and fibrous forms both contain infinite chains with P–P single bonds at 222 pm. Non-toxic, insoluble in organic solvents, ignites above 520 K. Used in safety matches (combined with antimony sulfide and glass powder on the strike surface).
Black phosphorus: Most stable form thermodynamically. Made from white P under high pressure. Puckered layered structure like graphite; conducts electricity; extremely kinetically inert (doesn't ignite even at 670 K).
4.3 Arsenic, Antimony and Bismuth
At room temperature all three exist as grey metallic-looking solids with the layered puckered hexagonal ring structure of black phosphorus (Fig. 15.3c). Going down from As to Bi, intralayer bond distances increase as expected, but interlayer spacing does not increase proportionally, effectively changing the coordination number from 3 to 6 as interlayer contacts become comparable to covalent bond distances. This structural feature correlates with increasing metallic character. All three burn in excess O₂:
They react with concentrated HNO₃ to give H₃AsO₄, hydrated Sb₂O₅ and Bi(NO₃)₃ respectively. Concentrated H₂SO₄ gives As₄O₆, Sb₂(SO₄)₃ and Bi₂(SO₄)₃.
Part 5 — Hydrides of Group 15
5.1 The Trihydrides EH₃
Each group 15 element forms a trihydride EH₃. The most important physical trend is the anomalously high boiling point of NH₃ compared with the trend PH₃ → AsH₃ → SbH₃, providing direct evidence for hydrogen bonding in liquid ammonia. Ammonia's boiling point of 240 K is roughly 55 K higher than expected from the trend.
Fig. 3 — Boiling points of Group 15 trihydrides. NH₃'s bp (~240 K) is ~57 K higher than extrapolated from PH₃-SbH₃ trend, providing direct evidence for N–H···N hydrogen bonding.
5.2 Ammonia — Industrial Giant and Chemical Reagent
Ammonia (NH₃) is a trigonal pyramidal molecule with H–N–H = 107° and N–H = 101 pm. It has a very low barrier to inversion (~24 kJ mol⁻¹) — the lone pair flips through the plane of the three H atoms — which is the basis of the umbrella inversion motion used as a model in spectroscopy.
The Haber process is one of the most important industrial chemical reactions:
The reaction is exothermic, so Le Chatelier's principle dictates that lower temperatures favour higher yield. However, lower temperatures slow the reaction to impractical rates. Industrial conditions compromise at T ≈ 723 K, P ≈ 20,000 kPa, with Fe₃O₄ catalyst (promoted by K₂O, SiO₂, Al₂O₃). The active phase is α-Fe formed by reduction of Fe₃O₄. The rate-determining step is adsorption and dissociation of N₂ onto the iron surface.
The entropy term: since 4 moles of gas → 2 moles, ΔrS° is negative (~−99 J K⁻¹ mol⁻¹), confirming that at very high temperatures the free energy becomes positive and yield drops sharply.
5.3 Phosphane (PH₃) and Heavier Trihydrides
The H–E–H angles in PH₃ (93.3°), AsH₃ (92.1°) and SbH₃ (91.6°) are close to 90°, suggesting the bonding uses almost pure p orbitals, with negligible s-character compared to NH₃'s 107° angle. PH₃ is extremely toxic, colourless and odourless. Unlike NH₃, aqueous PH₃ solutions are neutral (the P–H bond is not polar enough for hydrogen bonding). PH₃ in liquid NH₃ acts as an acid (eq. 15.30).
The spontaneous flammability sometimes observed in technical PH₃ is due to traces of diphosphane (P₂H₄). The Marsh test (historically used in forensic toxicology) uses the thermal instability of AsH₃ and SbH₃:
5.4 Hydrazine (N₂H₄) — Rocket Fuel and Reducing Agent
Hydrazine is a colourless liquid (mp 275 K, bp 386 K) miscible with water. Although ΔfH°(N₂H₄, l) = +50.6 kJ mol⁻¹ (endothermic), it is kinetically stable. The gas-phase structure adopts a gauche conformation to minimise lone-pair repulsions.
Industrial synthesis — the Raschig process:
NH₂Cl + NH₃ + NaOH → N₂H₄ + NaCl + H₂O (slow, rate-determining)
Glue or gelatin is added to inhibit the competing side-reaction where N₂H₄ is oxidised by NH₂Cl (reaction 15.37). Methyl derivatives of hydrazine (e.g., 1,1-dimethylhydrazine, UDMH) were used as rocket fuels in the Apollo missions.
In aqueous solution, N₂H₄ is a weaker base than NH₃: pK_b(1) = 6.05 (K_b = 8.9 × 10⁻⁷). It is also a reducing agent used to remove dissolved O₂ from industrial water boilers, preventing corrosion:
5.5 Hydrogen Azide (HN₃) and Azide Salts
Sodium azide NaN₃ is prepared from molten NaNH₂ + N₂O at 460 K. The azide ion [N₃]⁻ is isoelectronic with CO₂ — both are linear, 16-electron species. The symmetrical [N₃]⁻ ion has degenerate N–N bond lengths and is best described by resonance. HN₃ (hydrogen azide) is a dangerously explosive liquid (ΔfH° = +264 kJ mol⁻¹) and a weak acid (pKa = 4.75). Heavy metal azides (Pb(N₃)₂, AgN₃) are insoluble and explosive, used as initiating agents in detonators.
The azide group behaves as a pseudo-halogen — it forms salts analogous to halide salts, acts as a ligand in coordination chemistry, and exists as a dimer ([N₃]⁻ analogous to X₂ would be N₆, which has not been isolated).
Part 6 — Nitrides, Phosphides and Heavier Pnictides
6.1 Classification of Nitrides
Nitrides fall into four categories:
| Type | Examples | Character | Notes |
|---|---|---|---|
| Saline | Li₃N, Mg₃N₂, Ca₃N₂, AlN | Ionic (N³⁻) | Hydrolysed by water → NH₃ |
| Covalent (p-block) | BN, Si₃N₄, Sn₃N₄, S₄N₄ | Covalent network | Often ceramic-like; very hard |
| Interstitial (d-block) | TiN, ZrN, HfN, VN | Metallic | Hard, high-mp; cutting tool coatings |
| Pernitrides | BaN₂, SrN₂, PtN₂ | Contain [N₂]²⁻ | Made under extreme pressure |
The d-block interstitial nitrides (TiN is golden-coloured) are produced by thin-film chemical vapour deposition (CVD) and protect high-speed cutting tools. Silicon nitride (Si₃N₄) is remarkable for its high-temperature strength, thermal shock resistance and the ability of Si₃N₄ glow plugs in diesel engines to heat a combustion chamber to 600°C in a few seconds.
6.2 Phosphides and the Zintl Ions
Metal phosphides show enormous structural variety. Group 1 and 2 phosphides (M₃P, M₃P₂) are ionic and hydrolyse to give PH₃. Alkali metal phosphides with more complex stoichiometries contain cage or ring phosphide polyanions: [P₇]³⁻ (norbornane-like cage, structure 15.18), [P₁₁]³⁻ and [P₄]²⁻ (aromatic, 6π-electron ring — a phosphorus analogue of cyclobutadiene dianion).
Gallium arsenide (GaAs) — a III-V semiconductor in the zinc-blende structure — has a band gap of 1.42 eV at 298 K and electron mobility of 8500 cm² V⁻¹ s⁻¹. It is widely used in solid-state lasers, LEDs (infrared), photovoltaic cells and high-speed transistors. Advantages over Si: better optoelectronic properties, higher electron mobility. Disadvantages: more expensive, more brittle, lower thermal conductivity.
Bismuth cluster cations: [Bi₅]³⁺ (trigonal bipyramidal), [Bi₈]²⁺ (square antiprismatic), [Bi₉]⁵⁺ and the remarkable [Pd@Bi₁₀]⁴⁺ (Pd-centred pentagonal antiprismatic) are examples of homoatomic and heteroatomic Zintl ion clusters that obey Wade's rules.
Part 7 — Halides and Oxohalides
7.1 Nitrogen Halides
Nitrogen forms only trihalides (NX₃) — no NX₅ exists because N is too small to accommodate five substituents (not due to lack of d-orbitals). The key contrast is between NF₃ and NCl₃:
| Property | NF₃ | NCl₃ |
|---|---|---|
| ΔfH° (kJ mol⁻¹) | −132.1 (stable) | +230.0 (explosive) |
| Dipole moment (D) | 0.24 (very small) | 0.39 |
| Lewis basicity | None (no donor properties) | Moderate |
| Hydrolysis by alkali | Resistant | 2 NCl₃ + 6 OH⁻ → N₂ + 3 [OCl]⁻ + 3 Cl⁻ + 3 H₂O |
Why is NF₃ non-basic despite the lone pair? The highly electronegative F atoms withdraw electron density from N, making the lone pair unavailable for donation. The dipole moment in NF₃ is also surprisingly small: the N–F bond dipoles point from N to F (away from N), opposing the lone pair contribution to the dipole — nearly cancelling each other out.
7.2 Phosphorus Halides
Phosphorus forms both PX₃ and PX₅ (X = F, Cl, Br). The key contrast:
P₄ + 10 Cl₂ (excess) → 4 PCl₅ (excess Cl₂)
PCl₅ structure anomaly: In the gas phase, PCl₅ has a molecular trigonal bipyramidal structure (P–Clax = 214 pm, P–Cleq = 202 pm). However, in the solid state, it ionises into [PCl₄]⁺ (tetrahedral) and [PCl₆]⁻ (octahedral), crystallising with a CsCl-type structure. PBr₅ crystallises as [PBr₄]⁺Br⁻. This is a classic structure–bonding exam question.
PF₅ Berry pseudo-rotation: In solution, all five F atoms in PF₅ are equivalent in the ¹⁹F NMR spectrum (one doublet, coupling to ³¹P). This fluxional behaviour occurs via Berry pseudo-rotation — an interchange of axial and equatorial positions through a square-pyramidal intermediate. This is a key concept in Group 15 and Group 16 NMR.
POCl₃ (phosphoryl trichloride) is an important synthetic reagent — a phosphorylating and chlorinating agent used to make phosphate esters and anhydrous metal chlorides.
Part 8 — Oxides of Nitrogen
Nitrogen forms the most extensive series of oxides of any element, spanning all formal oxidation states from +1 (N₂O) to +5 (N₂O₅). All are endothermic or only weakly exothermic — their formation is thermodynamically unfavoured at 298 K. This has profound environmental consequences (NOₓ chemistry).
| Oxide | N Ox. State | Structure | ΔfH° (kJ mol⁻¹) | Key Properties |
|---|---|---|---|---|
| N₂O | +1 | Linear, bent: N–N–O (NNO) | +82.1 | Anaesthetic ("laughing gas"); supports combustion; weak oxidant |
| NO | +2 | Diatomic radical (bond order 2.5) | +90.2 | Paramagnetic; biological signalling molecule; brown ring test for NO₃⁻ |
| N₂O₃ | +3 | Planar, very long N–N (186 pm) | +83.7 (g) | Blue solid; acid anhydride of HNO₂; unstable above 195 K |
| NO₂ | +4 | Bent radical; N–O = 120 pm | +33.2 | Paramagnetic brown gas; in equilibrium with N₂O₄ |
| N₂O₄ | +4 | Planar; very long N–N (175 pm) | −19.5 (l) | Colourless diamagnetic solid; powerful oxidant |
| N₂O₅ | +5 | Gas: planar molecule; Solid: [NO₂]⁺[NO₃]⁻ | −43.1 (s) | Acid anhydride of HNO₃; made by P₄O₁₀ + 2 HNO₃ |
8.1 Nitrogen Monoxide (NO) — Radical, Ligand and Biological Messenger
NO has a bond order of 2.5 (one unpaired electron in a π* antibonding MO). Despite being thermodynamically unstable with respect to its elements (positive ΔfH°), it does not decompose below 1270 K. Industrial preparation: catalytic oxidation of NH₃ (the first step of the Ostwald process for HNO₃ manufacture):
The brown ring test for nitrate is based on reduction of [NO₃]⁻ by Fe²⁺ in dilute H₂SO₄ to give NO, which then forms the brown complex [Fe(NO)(OH₂)₅]²⁺ at the interface between the test solution and added cold concentrated H₂SO₄. The brown colour is consistent with an [NO]⁻ ligand bound to Fe(III) (confirmed by IR: ν(NO) = 1810 cm⁻¹).
In biological systems, NO is synthesised from L-arginine by NO synthase (NOS) enzymes and acts as a vasodilator (relaxes blood vessel walls), neurotransmitter and cytotoxic agent in immune responses. The 1998 Nobel Prize in Physiology/Medicine was awarded for the discovery of NO as a signalling molecule.
8.2 The NO₂/N₂O₄ Equilibrium
Laboratory preparation of NO₂/N₂O₄:
8.3 NOₓ and Environmental Chemistry
NOₓ (NO + NO₂) from vehicle exhausts and power plants contributes to photochemical smog and acid rain. The catalytic converter simultaneously oxidises CO, hydrocarbons and reduces NOₓ using Rh metal:
2 NO + 2 H₂ → N₂ + 2 H₂O
In the troposphere, NO reacts with HO• radicals and O₂ in a chain reaction that ultimately produces tropospheric ozone — a pollutant that damages lung tissue, in contrast to the beneficial stratospheric ozone layer.
Part 9 — Oxoacids of Nitrogen and Phosphorus
9.1 Nitrous Acid (HNO₂)
HNO₂ exists only in solution and vapour. It is a weak acid (pKa = 3.37) and unstable with respect to disproportionation:
Its reduction products depend on the reducing agent used — giving NO (with I⁻ or Fe²⁺), N₂O (with Sn²⁺), NH₂OH (with SO₂) or NH₃ (with Zn in alkali). An important kinetic point: in dilute solution, HNO₂ (not HNO₃) oxidises I⁻ to I₂, even though the standard potentials E°([NO₃]⁻/HNO₂) = +0.93 V and E°(HNO₂/NO) = +0.98 V are similar. HNO₂ is a faster, not more powerful, oxidising agent.
9.2 Nitric Acid (HNO₃) — The Ostwald Process
The industrial process involves three stages:
Step 2: 2 NO + O₂ ⇌ 2 NO₂
Step 3: 3 NO₂ + H₂O → 2 HNO₃ + NO (the NO is recycled)
This produces ~60% HNO₃ by weight. The azeotrope contains 68% HNO₃ (bp 393 K). "Fuming" nitric acid contains excess NO₂ (hence the orange-brown colour).
Nitric acid reacts with most metals (except Au, Pt-group, and passivated Fe/Cr). The reducing product depends on acid concentration:
Cu + 4 HNO₃(conc) → Cu(NO₃)₂ + 2 H₂O + 2 NO₂↑
Aqua regia (3:1 conc. HCl:conc. HNO₃) dissolves Au and Pt by combining the oxidising power of HNO₃ with the complexing ability of Cl⁻:
Fig. 4 — Structure of HNO₃ and [NO₃]⁻. In [NO₃]⁻, all three N–O bonds are equivalent (124 pm) due to delocalized π-bonding over all four atoms, rationalized by both VB resonance and MO theory (see Housecroft Fig. 5.25).
9.3 Oxoacids of Phosphorus — A System Based on Tetrahedral PO₄
All phosphorus(V) oxygen chemistry is built on the tetrahedral PO₄ unit. The key acids differ in how many OH groups are present (basicity) and whether P–H bonds exist (non-ionisable hydrogens):
| Acid | Formula | Basicity (OH groups) | P oxidation state | pKa values |
|---|---|---|---|---|
| Phosphinic acid (hypophosphorous) | H₃PO₂ | Monobasic (1 OH) | +1 | 1.24 |
| Phosphonic acid (phosphorous) | H₃PO₃ | Dibasic (2 OH) | +3 | 2.00, 6.59 |
| Phosphoric acid (orthophosphoric) | H₃PO₄ | Tribasic (3 OH) | +5 | 2.21, 7.21, 12.67 |
| Diphosphoric (pyrophosphoric) | H₄P₂O₇ | Tetrabasic | +5 | 0.85, 1.49, 5.77, 8.22 |
| Triphosphoric | H₅P₃O₁₀ | Pentabasic | +5 | ≈0, 0.89, 4.09, 6.98, 9.93 |
Condensed phosphates form by dehydration (P–O–P bridge formation): 2 H₃PO₄ → H₄P₂O₇ + H₂O. Tri-sodium polyphosphate Na₅P₃O₁₀ (manufactured by heating Na₂HPO₄ + NaH₂PO₄ at 550-650 K) was widely used in detergents as a water softener, but has been replaced by zeolites owing to eutrophication concerns (phosphate runoff causes algal blooms in lakes).
The biological importance of phosphates is immense: DNA and RNA backbones are phosphate esters, ATP/ADP interconversion is the cell's primary energy currency, and hydroxyapatite Ca₅(OH)(PO₄)₃ is the structural mineral of bones and teeth.
Part 10 — Phosphazenes
Phosphazenes are P(V)/N(III) compounds featuring alternating P and N in chains or rings, oligomers of the hypothetical unit N≡PR₂. The principal synthetic route:
The cyclic trimer (NPCl₂)₃ has a planar P₃N₃ ring with equal P–N bond lengths (~158 pm) — shorter than expected for a single P–N bond (~177 pm) but not double bond length either. All P–N bonds are equivalent despite the alternating double/single bond drawn in the Kekulé-like resonance structure. The modern bonding picture invokes negative hyperconjugation: donation of the N lone pair into σ*(P–Cl) antibonding orbitals, delocalising electron density and equalising bond lengths. Ionic (charge-separated) resonance structures also contribute.
The Cl substituents are replaced by nucleophiles (F, NR₂, OR, etc.) to give phosphazene polymers (polyphosphazenes) with diverse properties: fire resistance, elasticity, water repellence, and even suitability for artificial blood vessels (when R = –CH₂CF₃).
Part 11 — Group 16: The Chalcogens
11.1 Introduction and General Trends
The chalcogens — oxygen (O), sulfur (S), selenium (Se), tellurium (Te) and polonium (Po) — take their name from the Greek chalkos (copper ore), reflecting the prevalence of sulfide and selenide minerals. The group spans from the most electronegative of all chalcogens (O, χ = 3.4) to semi-metallic Te and metallic Po.
Unlike Group 15, where the highest oxidation state is +5, Group 16 has a maximum of +6 (as in SF₆, H₂SO₄). The commonest oxidation states are −2, 0, +2, +4 and +6. Oxygen is unique in forming no +4 or +6 compounds with lighter elements (OF₂ is O in the formal +2 state) because its small size prevents hypervalency.
Part 12 — Occurrence, Extraction and Uses (Group 16)
12.1 Oxygen — the Ubiquitous Oxidant
Dioxygen makes up 21% of Earth's atmosphere and 47% of Earth's crust (by mass, as oxygen-containing compounds). It is obtained industrially by fractional distillation of liquid air. The major use is in oxyacetylene and hydrogen flames, steel manufacture (basic oxygen process) and as a respiratory support. Laboratory preparations include decomposition of H₂O₂ or KClO₃:
2 KClO₃ Δ, MnO₂→ 3 O₂ + 2 KCl
12.2 Sulfur — From Mining to Sulfuric Acid
Native sulfur occurs near volcanoes. The traditional Frasch process (superheated water + compressed air to force molten S to the surface) is now largely replaced by sulfur recovery from natural gas and petroleum refining. The dominant sulfur source now involves the Claus process:
Over 80% of sulfur (mostly as H₂SO₄) goes into phosphate fertiliser production; it is also used in vulcanisation of rubber, making of paper, pharmaceuticals and explosives.
Part 13 — Physical Properties and Bonding (Group 16)
13.1 Bond Enthalpy Trends
The anomalously weak O–O (146 kJ mol⁻¹) and O–F (190 kJ mol⁻¹) bonds are again explained by inter-lone-pair repulsions between adjacent small electronegative atoms (exactly as for N–N and F–F in Group 15). In contrast, S–S (266 kJ mol⁻¹) is much stronger, which underlies sulfur's strong tendency to catenate — form chains and rings. This is exploited in the rich allotropy of sulfur (S₆, S₇, S₈, S₁₂, S₂₀ rings, catena-S∞).
The key contrast between O and S in terms of multiple bonding parallels N vs. P: O=O (498 kJ mol⁻¹) is very strong, making O₂ stable. S=S (427 kJ mol⁻¹) is much weaker relative to S–S, so S₈ rings (four S–S single bonds per ring) are more stable than S₂ dimers (with S=S double bonds), just as P₄ is more stable than P₂.
The absence of stable sulfur analogues of CO and NO reflects this: while C≡O (1076 kJ mol⁻¹) and N≡O (very strong) benefit from powerful pπ–pπ overlap, C≡S and N≡S bonds would require 3p–2p interaction, which is weaker and leads to less stable species.
Part 14 — The Elements: Allotropy and Reactivity
14.1 Dioxygen — Triplet and Singlet States
The ground state of O₂ is ³Σ_g⁻ (triplet) — two unpaired electrons occupying degenerate π*g MOs with the same spin. This paramagnetic character makes liquid O₂ attracted to magnets. Despite the high thermodynamic driving force for oxidation, the kinetic barrier is often high because most organic molecules are in a singlet state (paired electrons) and direct reaction with triplet O₂ is spin-forbidden.
Singlet oxygen (¹Δ_g), lying 94.7 kJ mol⁻¹ above the ground state, is far more reactive. It is generated photochemically (dye sensitisation), or chemically (H₂O₂ + NaOCl). In singlet O₂, the two electrons occupy one π*g orbital (paired), making it more electrophilic and able to react with alkenes via Diels-Alder-type pathways.
O₂ can be oxidised by PtF₆ (a very powerful oxidiser) to give [O₂]⁺[PtF₆]⁻. This reaction by Neil Bartlett in 1962 also led to the first xenon compound.
14.2 Ozone (O₃)
Ozone is endothermic (ΔfH° = +142.7 kJ mol⁻¹) and its liquid is dangerously explosive. It has a bent structure (bond angle 117°, O–O = 128 pm), with a bond order of 1.5 (two resonance structures). The O₃ molecule is a very powerful oxidising agent (E° = +2.07 V at pH 0), used in water purification as a substitute for chlorine. In the upper atmosphere, ozone absorbs UV radiation in the 200-300 nm range, protecting Earth from UV-B damage.
Reaction with [I]⁻ in alkaline solution or with PbS are used to detect and measure ozone:
4 O₃ + PbS → 4 O₂ + PbSO₄
14.3 Sulfur Allotropy — The Crown of Complexity
Fig. 5 — Key sulfur allotropes. The α-S₈ crown ring (orthorhombic, standard state) is the most stable at room temperature. Plastic sulfur (catena-S∞) is formed by quenching molten S in cold water.
When α-sulfur melts at 388 K, it initially gives a mobile yellow liquid (S₈ rings). Above 433 K, S₈ rings cleave homolytically into diradicals that polymerise rapidly, forming chains of up to 10⁶ atoms — the liquid becomes highly viscous and dark brown. Above ~473 K, the viscosity decreases as chains fragment. This distinctive viscosity–temperature behaviour is a characteristic exam topic.
Part 15 — Hydrides of Group 16
15.1 Hydrogen Peroxide (H₂O₂)
H₂O₂ is a colourless liquid (mp 272.6 K, bp 425 K, slightly higher than water due to extensive hydrogen bonding). The gas-phase structure shows an internal dihedral angle of 111° between the two O–H planes. The O–O bond distance (147.5 pm) is longer than in O₂ (121 pm), consistent with a single bond weakened by lone pair repulsions.
Industrial manufacture uses the anthraquinol oxidation cycle (Fig. 16.9): an organic anthraquinol is oxidised by O₂ to H₂O₂, which is extracted into water, then the organic product is re-reduced by H₂ over Pd catalyst. This avoids the hazards of direct synthesis from elements.
H₂O₂ is a powerful oxidising agent (E° = +1.78 V at pH 0) but also acts as a reducing agent towards strong oxidants like [MnO₄]⁻:
Cl₂ + H₂O₂ → 2 HCl + O₂ (H₂O₂ reduces Cl₂)
Applications include bleaching paper pulp (replacing chlorine), antiseptic, rocket propellant (mixed with hydrazine), and synthesis of sodium peroxoborate (used in laundry bleaches at lower temperatures than percarbonate).
15.2 H₂S, H₂Se and H₂Te
The H–E–H bond angles are ~92° (H₂S), 91° (H₂Se) and 90° (H₂Te) — close to 90° suggesting nearly pure p-orbital bonding with little s-character mixed in, unlike water (104.5°). Thermal stabilities decrease down the group: ΔfH° = −20.6 (H₂S), +29.7 (H₂Se), +99.6 kJ mol⁻¹ (H₂Te). All are toxic and have offensive odours; H₂S smells of rotten eggs.
Acid strength increases down the group: pKa(1) = 7.04 (H₂S), 4.0 (H₂Se), 3.0 (H₂Te). This is the opposite of the O–H group trend (HF < HCl < HBr < HI) and reflects the decreasing bond enthalpy of E–H bonds as atomic radius increases, making deprotonation easier. The second dissociation constant of H₂S is extremely small (~10⁻¹⁹), so S²⁻ barely exists in aqueous solution.
15.3 Polysulfanes H₂Sₓ and Polysulfide Ions [Sₓ]²⁻
Polysulfanes H₂Sₓ (x = 2–6) are obtained by acidification of polysulfide solutions. H₂S₂ has a structure like H₂O₂ with an 90.5° dihedral angle. All are thermodynamically unstable with respect to H₂S + S, but kinetically isolable. Polysulfide ions [Sₓ]²⁻ (x = 2–6) form when S dissolves in alkali metal sulfide solutions; higher polysulfides develop helical chains that can be chiral.
Part 16 — Halides of Group 16
16.1 Oxygen Fluorides — Unusual Oxidation States
OF₂ is the formal anhydride of hypofluorous acid HOF. In OF₂, oxygen is in the unusual +2 oxidation state (more electronegative F "wins"). It is highly toxic and reacts slowly with water:
H₂O + OF₂ → O₂ + 2 HF
O₂F₂ (dioxygen difluoride) is even more remarkable — an extremely powerful oxidising/fluorinating agent that inflames S at 93 K and reacts with UF₄ at ambient temperatures (reactions that normally require F₂ and high temperatures). O₂F₂ has a structure like H₂O₂ with long O–F bonds (157.5 pm) and a short O–O bond (122 pm), suggesting formal O–O multiple bond character contributed from resonance structures where O bears a positive charge.
16.2 Sulfur Fluorides — A Rich Chemistry
Sulfur forms a wider range of fluorides than any other element in the group:
Fig. 6 — VSEPR structures of SF₄ (see-saw, C₂ᵥ, lone pair in equatorial position of a trigonal bipyramid) and SF₆ (perfect octahedron, O_h symmetry). Note SF₄ undergoes rapid Berry pseudo-rotation in solution.
SF₄ is a selective fluorinating agent (converts C=O to CF₂ without destroying double bonds). It has a see-saw structure with the lone pair in an equatorial position of a trigonal bipyramid (consistent with VSEPR: lone pairs prefer equatorial sites).
SF₆ is the paragon of inertness — unattacked by acids, alkalis, water or most reagents up to 300°C. The thermodynamic driving force for hydrolysis is large (ΔrG° = −221 kJ mol⁻¹), but kinetic protection is almost absolute due to the crowded coordination environment around S. SF₆ is a greenhouse gas regulated under the Kyoto Protocol. Its primary use is as an electrical insulator in high-voltage equipment.
S₂F₁₀ (disulfur decafluoride) is extremely toxic and disproportionates on heating: S₂F₁₀ → SF₄ + SF₆.
16.3 Sulfur Chlorides and Oxochlorides
S₂Cl₂ (fuming orange liquid) and SCl₂ (dark red liquid) are manufactured by reacting Cl₂ with molten S. Both are hydrolysed by water to complex mixtures. Key reactions:
SOCl₂ + H₂O → SO₂ + 2 HCl (hence fuming in moist air)
SOCl₂ + RCO₂H → RCOCl + SO₂ + HCl (synthesis of acyl chlorides)
Sulfuryl chloride: SO₂Cl₂ — from SO₂ + Cl₂ with activated charcoal catalyst
Part 17 — Oxides of Group 16
17.1 Sulfur Dioxide (SO₂)
SO₂ is a colourless, dense gas with a pungent smell (bp 263 K). It is a bent molecule (O=S=O angle 119.5°, S–O = 143 pm). Resonance structures analogous to SO₃²⁻ (structure 16.44 in the textbook) show S with an octet and two equivalent S–O bonds with partial double bond character.
Environmental significance: SO₂ is the primary precursor of acid rain (H₂SO₃/H₂SO₄ formation in the atmosphere). The Contact process — the industrial synthesis of H₂SO₄ — begins with oxidation of SO₂:
The equilibrium is exothermic and favoured by lower temperatures (Le Chatelier), but rate requires higher temperature. The V₂O₅ catalyst makes the process viable at ~750 K (98% conversion). Increasing pressure shifts equilibrium right (fewer moles of gas), but the yield is already >98% without high pressure.
17.2 Sulfur Trioxide (SO₃)
SO₃ is a versatile electrophile. In the gas phase, planar monomeric molecules (D₃h, S–O = 142 pm) are in equilibrium with a cyclic trimer. The solid is polymorphic:
| Polymorph | Structure | Formation |
|---|---|---|
| γ-SO₃ | Trimeric units (ice-like crystals) | Rapid condensation |
| β-SO₃ | Polymeric chains | Traces of water catalyse conversion |
| α-SO₃ | Polymeric chains in layers | Most stable form |
Industrially, SO₃ is absorbed into concentrated H₂SO₄ to give oleum (H₂S₂O₇, disulfuric acid), then diluted with water. Direct absorption into water is not done because SO₃ + H₂O is violently exothermic, forming a dense acid mist that is difficult to handle.
Part 18 — Oxoacids of Sulfur and the Contact Process
18.1 Sulfuric Acid (H₂SO₄) — The Industrial Chemical
The Contact process:
Step 2: 2 SO₂ + O₂ ⇌ 2 SO₃ (V₂O₅, ~750 K, >98% conversion)
Step 3: SO₃ + H₂SO₄(conc) → H₂S₂O₇ (oleum)
Step 4: H₂S₂O₇ + H₂O → 2 H₂SO₄
Concentrated H₂SO₄ (98%) is a dense, viscous, strongly hydrogen-bonded liquid with an extremely high boiling point (610 K). It is a strong acid (pKa₁ ≈ −3, pKa₂ = 1.92), a powerful dehydrating agent, an oxidising agent (concentrated), and a sulfonating/nitrating agent. Dilute H₂SO₄ reacts with most metals; concentrated H₂SO₄ passivates Fe and oxidises Cu, C and S:
C + 2 H₂SO₄(conc, hot) → CO₂ + 2 SO₂ + 2 H₂O
18.2 Other Important Sulfur Oxoacids
| Acid | Formula | Key Feature |
|---|---|---|
| Sulfurous acid | H₂SO₃ | Exists only in solution; equilibrium SO₂ + H₂O ⇌ H₂SO₃. Reducing agent. |
| Dithionous acid | H₂S₂O₄ | Only salts known; [S₂O₄]²⁻ is a powerful reductant (E° = −1.12 V) |
| Thiosulfuric acid | H₂S₂O₃ | Only salts (thiosulfates, e.g. Na₂S₂O₃·5H₂O) stable; formed by SO₃²⁻ + S; key in photography (sodium thiosulfate as "hypo" — fixes silver halides) |
| Peroxydisulfuric acid | H₂S₂O₈ | Contains O–O bond; made by electrolytic oxidation; very powerful oxidant; hydrolyses to H₂O₂ + 2 HSO₄⁻ |
| Disulfuric acid (oleum, pyrosulfuric acid) | H₂S₂O₇ | Intermediate in Contact process; S–O–S bridge connects two SO₄ tetrahedra |
S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2 I⁻ (the tetrathionate product S₄O₆²⁻ has an S–S–S–S chain)
This is a standard exam and lab reaction. Thiosulfate (average S oxidation state = +2) is oxidised to tetrathionate (+2.5 average). Starch indicator turns blue-black with I₂ and becomes colourless when I₂ is consumed at the endpoint.
Part 19 — Exam Tips, Tricks & Common Pitfalls
High-Frequency JEE/NEET Questions & How to Handle Them
- Inert pair effect: Bi³⁺ more stable than Bi⁵⁺; BiF₅ is strongly oxidising; Bi₂O₃ is basic, not acidic. Same pattern in Tl, Pb.
- N vs P anomaly: NF₅ unknown (N too small), PCl₅ known. N forms strong π-bonds (N₂, NO, HNO₃), P does not (no stable P₂ at room temperature).
- Bond angle trend in EH₃: NH₃ 107° (sp³ character) → PH₃ 93° → AsH₃ 92° → SbH₃ 91.6° (nearly pure p-orbitals). As angle → 90°, the bonds use more p character.
- Basicity of EH₃: NH₃ >> PH₃ > AsH₃ in water. NH₃ is basic (pKb = 4.74), PH₃ is neutral, AsH₃/SbH₃ are acidic in liquid NH₃.
- Acid strength in H₂E: H₂O < H₂S < H₂Se < H₂Te. Decreasing E–H bond enthalpy dominates over other effects as you go down the group.
- Counting P-acid protons: Only O–H hydrogens are acidic. H₃PO₂ is monobasic (2 P–H, 1 O–H); H₃PO₃ dibasic (1 P–H, 2 O–H); H₃PO₄ tribasic (3 O–H).
- Reducing power of P-oxoacids: H₃PO₂ > H₃PO₃ > H₃PO₄. Lower oxidation state of P → stronger reductant. H₃PO₂ and H₃PO₃ disproportionate on heating.
- PCl₅ structure: Gas phase = trigonal bipyramidal molecule. Solid state = [PCl₄]⁺[PCl₆]⁻ ionic. PBr₅ solid = [PBr₄]⁺Br⁻. This is due to the ionic lattice energy compensating for the charge separation.
- Contact process conditions: ~750 K (compromise — high T for rate, low T for yield), 1–2 atm, V₂O₅ catalyst. SO₃ absorbed in H₂SO₄ (not water directly — mist formation).
- Ostwald process: NH₃ → NO (Pt/Rh, 1300 K) → NO₂ → HNO₃ (absorbed in water). ~60% HNO₃ produced; concentrated to 98% using H₂SO₄ dehydration.
- HONC-1234 rule: H forms 1 bond, O forms 2, N forms 3, C forms 4 in typical organic compounds. For counting oxidation states of N in nitrogenous compounds, trace bonds systematically.
- "Black Phosphorus is the Most Stable": Stability order: Black > Red > White phosphorus (thermodynamic stability increases from white to black).
- SF₄ VSEPR: Lone pair goes equatorial in a TBP → see-saw shape (C₂ᵥ). Confirmed by two types of F in low-temperature ¹⁹F NMR (axial and equatorial), but fluxional at room temperature.
- For N oxides: N₂O (dental anaesthetic/laughing gas), NO (brown ring test, biological signalling), NO₂ (brown gas, paramagnetic), N₂O₄ (rocket oxidiser), N₂O₅ (acid anhydride of HNO₃).
- "Strong acids at +5, amphoteric at +3": P₄O₁₀ and H₃PO₄ are acidic; As₂O₃/As₄O₆ and Sb₄O₆ are amphoteric; Bi₂O₃ is basic. Descending group 15 → oxides become more basic at the same oxidation state.
- Azide ≡ CO₂ isostere: [N₃]⁻ is isoelectronic with CO₂, N₂O, [NO₂]⁺ and CS₂. Drawing their Lewis structures and MO diagrams is a common exam task.
Commonly Confused Pairs and How to Distinguish Them
| Pair | Key Distinction |
|---|---|
| PCl₃ vs PCl₅ in solid state | PCl₃ is molecular; PCl₅ is ionic [PCl₄]⁺[PCl₆]⁻ |
| HNO₂ vs HNO₃ as oxidising agents | HNO₂ is kinetically faster; HNO₃ is thermodynamically more powerful. HNO₂ oxidises I⁻ in dilute solution; HNO₃ does not (kinetically controlled) |
| H₃PO₃ vs H₃PO₂ | H₃PO₃ dibasic (2 O–H, 1 P–H); H₃PO₂ monobasic (2 P–H, 1 O–H) |
| H₂S₂O₃ vs H₂S₂O₄ | Thiosulfate [S₂O₃]²⁻ = S(VI) + S(−2); Dithionite [S₂O₄]²⁻ has very long S–S bond (eclipsed); both are reducing agents |
| SO₂ vs SO₃ structural | SO₂: bent (C₂ᵥ, 119.5°, dipole ≠ 0); SO₃: planar trigonal (D₃h, dipole = 0) |
| White P vs Red P reactivity | White ignites above 323 K, very reactive; Red ignites above 520 K, less reactive. Both produce P₄O₁₀ (or P₄O₆) on burning in air |
SVG — Oxidation State Summary for Group 15 Nitrogen Species
Fig. 7 — Oxidation states of nitrogen from −3 to +5 with representative species. Each state is experimentally accessible — nitrogen chemistry spans the widest redox range of any Group 15 element.
Multi-Concept Worked Example (JEE-style)
For N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g): ΔrH° = −92 kJ mol⁻¹, ΔrG° = −33 kJ mol⁻¹ (298 K).
Step 1: Calculate ΔrS°. Use ΔrG° = ΔrH° − TΔrS°: −33 = −92 − 298 × ΔrS° → ΔrS° = (−92 + 33)/298 = −0.1980/298 ≈ −198 J K⁻¹ mol⁻¹.
Step 2: Confirm sign. The reaction converts 4 mol gas → 2 mol gas, so ΔrS° must be negative. ✓
Step 3: Calculate ln K at 298 K. ΔrG° = −RT ln K → ln K = −ΔrG°/(RT) = 33000/(8.314 × 298) = 13.3; K = e¹³·³ ≈ 6 × 10⁵. Equilibrium very favourable at 298 K.
Step 4: At 700 K, ΔrH° ≈ −52 kJ mol⁻¹ and ΔrG° = +27 kJ mol⁻¹. Now ln K = −27000/(8.314 × 700) = −4.64, K = e⁻⁴·⁶⁴ ≈ 0.01. Equilibrium highly unfavourable!
Conclusion: This is why Haber process uses a catalyst — the thermodynamics at 700 K are terrible, but the catalyst is essential to achieve any practical rate. The temperature is a compromise between kinetics (fast at high T) and thermodynamics (favourable at low T).
Reaction Summary Table: Key Group 15 & 16 Reactions
| Reaction | Equation | Exam Tag |
|---|---|---|
| Haber Process | N₂ + 3 H₂ ⇌ 2 NH₃ (Fe₃O₄, 723 K, 200 atm) | JEENEET |
| Ostwald Process Step 1 | 4 NH₃ + 5 O₂ → 4 NO + 6 H₂O (Pt/Rh, 1300 K) | JEENEET |
| P₄ extraction | 2 Ca₃(PO₄)₂ + 6 SiO₂ + 10 C → P₄ + 6 CaSiO₃ + 10 CO | GATEJAM |
| Aqua regia dissolves Au | Au + HNO₃ + 4 HCl → HAuCl₄ + NO + 2 H₂O | JEE |
| Raschig process (hydrazine) | NH₃ + NaOCl → NH₂Cl + NaOH; NH₂Cl + NH₃ + NaOH → N₂H₄ | CSIR |
| Contact process (SO₃) | 2 SO₂ + O₂ ⇌ 2 SO₃ (V₂O₅, ~750 K) | JEENEET |
| Dilute HNO₃ + Cu | 3 Cu + 8 HNO₃(dil) → 3 Cu(NO₃)₂ + 4 H₂O + 2 NO↑ | JEE |
| Conc. HNO₃ + Cu | Cu + 4 HNO₃(conc) → Cu(NO₃)₂ + 2 H₂O + 2 NO₂↑ | JEE |
| NaN₃ in airbag | 2 NaN₃(s) → 2 Na + 3 N₂ (electrical impulse) | NEET |
| P₄ + NaOH | P₄ + 3 NaOH + 3 H₂O → 3 NaH₂PO₂ + PH₃ | GATE |
| Thiosulfate + I₂ (iodometry) | 2 S₂O₃²⁻ + I₂ → S₄O₆²⁻ + 2 I⁻ | NEETGATE |
| H₂SO₄ (conc) + Cu | Cu + 2 H₂SO₄(conc, hot) → CuSO₄ + SO₂ + 2 H₂O | JEE |
- NH₃ — trigonal pyramidal, H–N–H = 107°
- N₂H₄ — gauche conformation (gas phase), two NH₂ groups
- HNO₃ — planar (Cs), two different N–O bonds (96 pm for N–O–H; 121 and 119 pm for the two N=O/N–O)
- [NO₃]⁻ — trigonal planar D₃h, all N–O = 124 pm
- PCl₃ — trigonal pyramidal (like NH₃)
- PCl₅ — gas: trigonal bipyramidal; solid: [PCl₄]⁺[PCl₆]⁻
- SF₄ — see-saw (C₂ᵥ); SF₆ — octahedral (Oh); SOCl₂ — pyramidal
- SO₂ — bent 119.5°; SO₃ — planar D₃h; H₂SO₄ — tetrahedral S
- H₂O₂ — gauche dihedral 111° (gas); peroxide O–O = 147.5 pm
- P₄ tetrahedron with 60° P–P–P angle (strain)
© Article content based on original synthesis and study of: Housecroft & Sharpe, Inorganic Chemistry (4th ed.), Chapters 15 & 16. All reactions and molecular data are validated against IUPAC nomenclature and primary literature. For educational use — JEE Advanced, NEET, IIT-JAM, GATE, CSIR-NET, BITSAT, TGT/PGT.
