Inorganic Chemistry-3 BCH4T07

Syllabus topics (official RTMNU)

Coordination compounds

• Distinction among simple salts, double salts and coordination compounds.
• Terminology and nomenclature of Coordination compounds.
• Types of ligands on the basis of denticity.
• Werner’s Coordination theory and its experimental verification.
• Sidgwicks electronic interpretation, EAN rule with examples.
• Valence Bond Theory of transition metal complexes.
• Determination of structures and magnetic properties of complexes on the basis of VBT.
• Chelates: Classification and their application.

Simple salts, double salts and coordination compounds

When two or more stable compounds combine in a fixed ratio, the product may behave in solution either as a mixture of its simple ions or as a new complex unit. This difference separates double salts from coordination compounds.

Terminology and IUPAC nomenclature

• Central metal atom/ion: the atom to which ligands are attached, e.g. Fe²⁺ in [Fe(CN)₆]⁴⁻. It acts as a Lewis acid.
• Ligand: ion or molecule that donates an electron pair to the metal (Lewis base), e.g. NH₃, H₂O, Cl⁻, CN⁻.
• Donor atom: the atom of the ligand that actually donates the pair (N in NH₃).
• Coordination number (CN): number of donor atoms bonded to the metal. CN of Fe in [Fe(CN)₆]⁴⁻ = 6.
• Coordination sphere: metal + ligands written inside square brackets; ions outside are counter ions.
• Oxidation number: charge on the metal after removing all ligands with their electron pairs.

Rules of naming

1. Name the cation first, then the anion (as in simple salts).
2. Inside the sphere, ligands are named first in alphabetical order, then the metal.
3. Anionic ligands end in "-o": Cl⁻ chlorido, CN⁻ cyanido, OH⁻ hydroxido, NO₂⁻ (N-bonded) nitrito-N. Neutral ligands keep their name except H₂O aqua, NH₃ ammine, CO carbonyl, NO nitrosyl.
4. Prefixes di, tri, tetra for simple ligands; bis, tris, tetrakis for ligands whose names already have a prefix (e.g. ethylenediamine → tris(ethylenediamine)). Prefixes are ignored when alphabetising.
5. Oxidation state of the metal in Roman numerals in brackets.
6. If the complex is an anion, the metal name ends in "-ate" (ferrate, cuprate, argentate, aurate, plumbate, stannate for Fe, Cu, Ag, Au, Pb, Sn).

• [Co(NH₃)₆]Cl₃ — hexaamminecobalt(III) chloride
• K₄[Fe(CN)₆] — potassium hexacyanidoferrate(II)
• [Co(NH₃)₅Cl]Cl₂ — pentaamminechloridocobalt(III) chloride
• [Pt(NH₃)₂Cl₂] — diamminedichloridoplatinum(II)

Types of ligands on the basis of denticity

Denticity = number of donor atoms through which one ligand binds to the metal.

Ambidentate ligands have two possible donor atoms but use one at a time: NO₂⁻ (N or O), SCN⁻ (S or N), CN⁻ (C or N).

Werner's coordination theory and its verification

Postulates

1. A metal shows two types of valency: primary (ionisable, = oxidation state, satisfied by negative ions) and secondary (non-ionisable, = coordination number, satisfied by negative ions or neutral molecules).
2. Every metal has a fixed number of secondary valencies (e.g. 6 for Co³⁺, 4 for Pt²⁺).
3. Secondary valencies are directed in space, giving a definite geometry (octahedral for CN 6; square planar or tetrahedral for CN 4).

Experimental verification (cobalt(III) chloride–ammonia complexes)

Molar conductance measurements matched the number of ions, and the existence of two isomers (green and violet) of CoCl₃·4NH₃ proved an octahedral arrangement.

Sidgwick's electronic interpretation and EAN rule

Sidgwick said each ligand donates a lone pair to the metal, forming a coordinate (dative) bond. The metal goes on accepting pairs until its total electrons equal those of the next noble gas.

• K₄[Fe(CN)₆]: Fe Z = 26, Fe²⁺ → 24, + 12 = 36 (Kr) — obeys.
• [Co(NH₃)₆]³⁺: 27 − 3 + 12 = 36 — obeys.
• [Ni(CO)₄]: 28 − 0 + 8 = 36 — obeys.
• K₃[Fe(CN)₆]: 26 − 3 + 12 = 35 — does not obey, yet stable. So EAN is a useful guide, not a strict law.

Valence Bond Theory (VBT) of complexes

1. The metal ion provides as many vacant hybrid orbitals as its coordination number.
2. Ligand lone pairs enter these orbitals, forming σ coordinate bonds.
3. Type of hybridisation fixes the shape. Strong-field ligands (CN⁻, CO) may pair up d-electrons, making inner d-orbitals free.

Structure and magnetic properties

• Magnetic moment μ = √[n(n+2)] BM, where n = number of unpaired electrons.
• [Ni(CN)₄]²⁻: Ni²⁺ is 3d⁸; CN⁻ pairs the electrons, one 3d orbital is free → dsp², square planar, n = 0, diamagnetic.
• [NiCl₄]²⁻: Cl⁻ is weak, no pairing, two unpaired electrons → sp³, tetrahedral, μ ≈ 2.83 BM, paramagnetic.
• [Co(NH₃)₆]³⁺: Co³⁺ 3d⁶, all paired → d²sp³, diamagnetic. [CoF₆]³⁻: 4 unpaired → sp³d², μ ≈ 4.90 BM.

Limitations of VBT: cannot explain colour, gives no energy picture, does not explain why some ligands are strong field.

Chelates: classification and applications

A chelate is a complex in which a polydentate ligand binds the metal through two or more donor atoms, forming a ring. Five- and six-membered rings are most stable. Chelates are more stable than similar complexes with monodentate ligands — the chelate effect (favourable entropy).

• By donor atoms: N,N (en, dimethylglyoxime); O,O (oxalate, acetylacetonate); N,O (glycinate, 8-hydroxyquinoline).
• By denticity of the ligand: bidentate, tridentate, ... hexadentate (EDTA).
• By charge: neutral inner-metallic chelates (Ni-DMG) or charged chelate ions.

Applications

• Estimation of Ni²⁺ as red Ni-dimethylglyoxime; Al³⁺/Mg²⁺ with oxine.
• Hardness of water: complexometric titration with EDTA (Ca²⁺, Mg²⁺).
• Medicine: EDTA removes lead in lead poisoning; cisplatin binds DNA.
• Biology: chlorophyll (Mg), haemoglobin (Fe), vitamin B₁₂ (Co) are natural chelates.

Likely exam questions

Long answer

1. State the postulates of Werner's coordination theory. How did Werner verify it using the cobalt chloride–ammonia complexes?
2. Explain the formation, structure and magnetic properties of [Ni(CN)₄]²⁻ and [NiCl₄]²⁻ on the basis of Valence Bond Theory.

Short answer

1. Distinguish between a double salt and a coordination compound with one example each.
2. What is the EAN rule? Calculate the EAN of iron in K₄[Fe(CN)₆].
3. What are chelates? Give two applications.

Syllabus topics (official RTMNU)

(A) Isomerism in coordination compounds

• Structural isomerism and Stereoisomerism (Geometrical and optical) in coordination compounds with four and six coordination numbers.

(B) Acids and bases

• Arrhenius theory and its limitations.
• Lowry-Bronsted concept, conjugate acid-base pair, relative strengths of acids and bases, mono and polyprotic acids and bases.
• Limitations of L-B concept.
• Solvent system concept and its limitations.
• Lux-Flood concept of metal oxides.
• Lewis theory, types of Lewis acids and bases, Strengths of Lewis acids and bases.
• Advantages and limitation of Lewis concept.

(A) Structural isomerism in coordination compounds

Isomers have the same molecular formula but a different arrangement of atoms. Structural isomers differ in which atoms are bonded to the metal.

• Ionisation isomerism: exchange of an ion between the sphere and outside. [Co(NH₃)₅Br]SO₄ (red-violet, gives SO₄²⁻ test with BaCl₂) and [Co(NH₃)₅SO₄]Br (red, gives Br⁻ test with AgNO₃).
• Hydrate (solvate) isomerism: water inside vs outside the sphere. [Cr(H₂O)₆]Cl₃ (violet), [Cr(H₂O)₅Cl]Cl₂·H₂O (grey-green), [Cr(H₂O)₄Cl₂]Cl·2H₂O (dark green).
• Linkage isomerism: an ambidentate ligand binds through different atoms. [Co(NH₃)₅(NO₂)]Cl₂ (nitrito-N, yellow) and [Co(NH₃)₅(ONO)]Cl₂ (nitrito-O, red).
• Coordination isomerism: in salts where both cation and anion are complexes, ligands are exchanged between the two metals. [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].
• Ligand isomerism: the ligands themselves are isomers (1,2-diaminopropane vs 1,3-diaminopropane).

Stereoisomerism (geometrical and optical)

Same bonds, different arrangement in space.

Coordination number 4

• Tetrahedral: all four positions are adjacent, so no geometrical isomers. Optical isomers are possible only with four different groups (rare).
• Square planar [Ma₂b₂]: cis (same groups side by side, 90°) and trans (opposite, 180°). Example: cis- and trans-[Pt(NH₃)₂Cl₂]. Cis-platin is an anticancer drug; the trans form is not active.
• [Mabcd] square planar gives three geometrical isomers. Square planar complexes are usually optically inactive because the plane itself is a plane of symmetry.

Coordination number 6 (octahedral)

• [Ma₄b₂]: cis and trans. Example: cis- (violet) and trans- (green) [Co(NH₃)₄Cl₂]⁺.
• [Ma₃b₃]: facial (fac — three same groups on one face) and meridional (mer — three same groups on a plane through the metal). Example: [Co(NH₃)₃(NO₂)₃].
• [M(AA)₂b₂]: cis and trans, e.g. [Co(en)₂Cl₂]⁺.

Optical isomerism

A complex with no plane or centre of symmetry is chiral and exists as two non-superimposable mirror images (d and l, or Δ and Λ) that rotate plane-polarised light in opposite directions.

• [Co(en)₃]³⁺ — d and l forms (propeller shape).
• cis-[Co(en)₂Cl₂]⁺ is optically active; trans-[Co(en)₂Cl₂]⁺ has a plane of symmetry and is optically inactive.

(B) Arrhenius theory and Lowry–Brønsted concept

Arrhenius theory

An acid gives H⁺ (H₃O⁺) in water; a base gives OH⁻ in water. Neutralisation: H⁺ + OH⁻ → H₂O.

Limitations: applies only to aqueous solutions; cannot explain basicity of NH₃ or Na₂CO₃ (no OH⁻ group), or acidity of CO₂, AlCl₃; free H⁺ does not exist in water.

Lowry–Brønsted concept

An acid is a proton donor; a base is a proton acceptor.

HCl + H₂O ⇌ H₃O⁺ + Cl⁻

• Conjugate pair: acid and base differing by one H⁺. HCl/Cl⁻ and H₃O⁺/H₂O above. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻: pairs NH₄⁺/NH₃ and H₂O/OH⁻.
• Relative strength: a strong acid has a weak conjugate base (HCl → Cl⁻); a weak acid has a strong conjugate base (CH₃COOH → CH₃COO⁻). The equilibrium lies on the side of the weaker acid and base.
• Monoprotic and polyprotic: HCl gives one proton; H₂SO₄ two; H₃PO₄ three in steps (Ka₁ > Ka₂ > Ka₃). Similarly polyprotic bases like CO₃²⁻ (accepts two protons) and PO₄³⁻.
• Amphiprotic: H₂O, HCO₃⁻ can donate or accept protons.

Limitations: cannot explain acidic nature of substances with no proton (BF₃, AlCl₃, SO₃) or acid–base reactions without proton transfer (CaO + SO₃ → CaSO₄).

Solvent system concept and Lux–Flood concept

Solvent system (Franklin, Cady–Elsey): a solvent self-ionises. An acid increases the concentration of the solvent's cation; a base increases the solvent's anion.

• 2H₂O ⇌ H₃O⁺ + OH⁻
• 2NH₃ (liq) ⇌ NH₄⁺ + NH₂⁻ — so NH₄Cl is an acid and NaNH₂ a base in liquid NH₃.
• 2BrF₃ ⇌ BrF₂⁺ + BrF₄⁻; N₂O₄ ⇌ NO⁺ + NO₃⁻.

Limitations: only for self-ionising solvents; ignores reactions in non-ionising solvents or with no solvent.

Lux–Flood concept (for oxide melts): a base is an oxide ion donor; an acid is an oxide ion acceptor.

CaO (base) + SiO₂ (acid) → CaSiO₃ i.e. O²⁻ + SiO₂ → SiO₃²⁻

Metal oxides (CaO, Na₂O) are bases; non-metal oxides (SiO₂, CO₂, SO₃, P₄O₁₀) are acids. Used in metallurgy (slag formation) and glass making.

Lewis theory

A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor. The product is an adduct with a coordinate bond: BF₃ + :NH₃ → F₃B←NH₃.

Types

• Lewis acids: molecules with incomplete octet (BF₃, AlCl₃); simple cations (Ag⁺, Fe³⁺, Cu²⁺); molecules with a multiple bond to an electronegative atom (CO₂, SO₃); atoms that can expand the octet (SiF₄, SnCl₄).
• Lewis bases: neutral molecules with a lone pair (NH₃, H₂O, ROH); anions (F⁻, OH⁻, CN⁻); molecules with π electrons (ethene, benzene).

Strength

• Acid strength rises with higher positive charge and smaller size of the cation (Al³⁺ > Mg²⁺ > Na⁺).
• Boron halides: BI₃ > BBr₃ > BCl₃ > BF₃ — BF₃ is weakest because of strong B–F back-bonding (pπ–pπ).
• Base strength depends on availability of the lone pair; steric crowding reduces it.

Advantages: most general theory; covers reactions without protons or solvents; explains complex formation and many organic reactions (AlCl₃ in Friedel–Crafts).

Limitations: no uniform scale of strength (order depends on the reference acid/base); covers almost every reaction, so it is too broad; does not explain the strength of protonic acids like HCl and H₂SO₄; acid–base reactions should be fast, but many coordinate-bond formations are slow.

Likely exam questions

Long answer

1. Explain geometrical and optical isomerism in coordination compounds with coordination numbers four and six, giving suitable examples.
2. Discuss the Lowry–Brønsted concept of acids and bases. Explain conjugate acid–base pairs and the relative strength of acids and bases. Give its limitations.

Short answer

1. What is linkage isomerism? Give one example.
2. Explain the Lux–Flood concept with an example.
3. What are Lewis acids? Give the types with one example each.

Syllabus topics (official RTMNU)

Oxidation and reduction

• Various definitions of oxidation and reduction, Single electrode potential, standard electrode potential, electrochemical series and its applications.
• Concept of overvoltage.
• Balancing of redox reaction by ion-electron method, Use of redox potential data: Analysis of Redox cycle, redox stability in water, Latimer diagram of Chlorine and Oxygen, Frost diagram of Nitrogen and Oxygen, and Porbaix diagrams of Iron.
• Predicting disproportionation and comproportionation phenomena.
• Deriving Frost diagram from Latimer diagram.

Definitions of oxidation and reduction

Example: Zn + Cu²⁺ → Zn²⁺ + Cu. Zn loses 2e⁻ (oxidised, reducing agent); Cu²⁺ gains 2e⁻ (reduced, oxidising agent). Oxidation and reduction always occur together (redox).

Single electrode potential, standard electrode potential and electrochemical series

• Single electrode potential: the potential difference set up between a metal and a solution of its ions at equilibrium (M ⇌ Mⁿ⁺ + ne⁻). It cannot be measured alone; it is measured against a reference electrode.
• Standard electrode potential (E°): the electrode potential at 298 K, 1 M ion concentration and 1 bar gas pressure, measured against the Standard Hydrogen Electrode (SHE), whose E° is taken as 0.00 V. By IUPAC convention E° values are written as reduction potentials.
• Nernst equation (for Mⁿ⁺ + ne⁻ → M): E = E° − (0.0591/n) log (1/[Mⁿ⁺]) at 298 K.

Electrochemical series

Electrodes arranged in order of their standard reduction potentials. Some values: Li⁺/Li −3.04 V, K⁺/K −2.93, Zn²⁺/Zn −0.76, Fe²⁺/Fe −0.44, H⁺/H₂ 0.00, Cu²⁺/Cu +0.34, Ag⁺/Ag +0.80, F₂/F⁻ +2.87 V.

Applications

1. Oxidising/reducing power: higher E° → stronger oxidising agent (F₂); lower E° → stronger reducing agent (Li).
2. Displacement: a metal with lower E° displaces one with higher E° from its salt (Zn displaces Cu).
3. H₂ from acids: metals with negative E° liberate H₂ from dilute acids; Cu, Ag do not.
4. EMF of a cell: E°cell = E°cathode − E°anode. For Daniell cell: 0.34 − (−0.76) = 1.10 V.
5. Feasibility: a positive E°cell means ΔG° = −nFE° is negative, so the reaction is spontaneous.

Concept of overvoltage

Overvoltage (overpotential) is the extra voltage, above the theoretical (reversible) decomposition potential, needed to make a gas or metal discharge at an electrode at a noticeable rate. It arises from slow electrode reactions, especially gas evolution. Hydrogen overvoltage is high on Hg, Pb, Zn and very low on platinised Pt. Because of high H₂ overvoltage on Hg, Na⁺ can be discharged at a mercury cathode in the chlor-alkali process, and Zn can be electroplated from aqueous solution.

Balancing redox reactions by the ion-electron method

Example: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (acidic medium).

1. Split into half-reactions: Fe²⁺ → Fe³⁺ and MnO₄⁻ → Mn²⁺.
2. Balance atoms other than O and H.
3. Balance O with H₂O and H with H⁺: MnO₄⁻ + 8H⁺ → Mn²⁺ + 4H₂O.
4. Balance charge with electrons: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O; Fe²⁺ → Fe³⁺ + e⁻.
5. Equalise electrons (multiply iron half by 5) and add.

In basic medium, balance as for acid and then add OH⁻ to both sides to cancel H⁺.

Use of redox potential data: redox cycle and redox stability in water

• Redox cycle: a series of linked steps in which a species is oxidised and then regenerated by reduction (or vice versa), so it acts catalytically — e.g. Fe³⁺/Fe²⁺ cycling. E° values tell whether each step is downhill.
• Oxidation by water: 2H₂O + 2e⁻ → H₂ + 2OH⁻ (or 2H⁺ + 2e⁻ → H₂, E° = 0). Metals with strongly negative E° (Na, K) are oxidised by water, releasing H₂.
• Reduction of water: O₂ + 4H⁺ + 4e⁻ → 2H₂O, E° = +1.23 V. Very strong oxidising agents (E° > 1.23 V, e.g. Co³⁺, F₂) oxidise water to O₂.
• Stability field of water: a species is stable in water if its E° lies between the two lines (roughly 0 to 1.23 V at pH 0, both falling by 0.059 V per pH unit). Practically, overvoltage widens this window by about 0.6 V.

Latimer, Frost and Pourbaix diagrams

Latimer diagram

Species of one element are written in decreasing oxidation state from left to right, with E° (V) for each step written over the arrow. Chlorine in acid:

ClO₄⁻ (+7) →(1.20) ClO₃⁻ (+5) →(1.18) HClO₂ (+3) →(1.67) HClO (+1) →(1.63) Cl₂ (0) →(1.36) Cl⁻ (−1)

Oxygen in acid: O₂ (0) →(+0.70) H₂O₂ (−1) →(+1.76) H₂O (−2); overall O₂/H₂O = +1.23 V.

Non-adjacent E°: E° = (n₁E°₁ + n₂E°₂)/(n₁ + n₂). For O₂ → H₂O: (2×0.70 + 2×1.76)/4 = 1.23 V.

Predicting disproportionation and comproportionation

• A species disproportionates if E° on its right is greater than E° on its left. H₂O₂: right 1.76 > left 0.70, so 2H₂O₂ → 2H₂O + O₂ is spontaneous. Similarly HClO₂ (1.67 > 1.18) disproportionates.
• Comproportionation is the reverse: two species of the same element in higher and lower states react to give the intermediate one, when the intermediate lies below the line joining them on a Frost diagram (e.g. Ag²⁺ + Ag → 2Ag⁺).

Frost diagram

Plot of nE° (volt-equivalent, which is proportional to free energy, since ΔG° = −nFE°) on the y-axis against oxidation number N on the x-axis. The element (N = 0) is at nE° = 0.

• Deriving from Latimer: for each species, nE° = (oxidation number) × (E° of the couple from that species to the element), i.e. sum the n×E° steps from 0 to that state. For oxygen: O₂ at (0, 0); H₂O₂ at (−1, −0.70); H₂O at (−2, −2.46).
• Lowest point = most stable state (H₂O for oxygen). Slope of a line between two points = E° of that couple; steeper slope = stronger oxidising agent.
• A point lying above the line joining its neighbours disproportionates (H₂O₂). A point below the line is stable to disproportionation.
• Nitrogen (acid): NH₄⁺ (−3) and N₂ (0) lie low; N₂ is very stable; HNO₃ (+5) lies high, so nitrate is a strong oxidising agent; NH₂OH lies above its neighbours and disproportionates.

Pourbaix diagram of iron

A plot of E (y-axis) vs pH (x-axis) showing which iron species is stable in water. Horizontal lines are pure electron-transfer couples (Fe³⁺/Fe²⁺, E° = 0.77 V, pH independent); vertical lines are pure acid–base/precipitation equilibria (Fe³⁺ → Fe(OH)₃ near pH 2–3); sloping lines involve both H⁺ and e⁻. Regions: Fe³⁺ at low pH, high E; Fe²⁺ at low pH, moderate E; Fe(OH)₃ at higher pH and oxidising conditions; Fe(OH)₂ at high pH, low E; metallic Fe at the bottom (very low E). It explains why iron rusts in aerated neutral water (Fe(OH)₃ / Fe₂O₃ zone) and why Fe²⁺ exists only in acidic or reducing water.

Likely exam questions

Long answer

1. What is the electrochemical series? Discuss its applications.
2. Explain Latimer and Frost diagrams with reference to oxygen. How is a Frost diagram derived from a Latimer diagram, and how is disproportionation predicted?

Short answer

1. Define standard electrode potential.
2. What is overvoltage? Give one application.
3. Balance by the ion-electron method: MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (acidic medium).

Syllabus topics (official RTMNU)

(A) Inorganic Polymers

• Silicones: Introduction, Nomenclature, preparation, properties and uses, General introduction to Silicon oils, Silicone Elastomers and Silicon Resins

(B) Phosphonitrilic halide polymers

• Introduction, Preparation, properties and uses.
• Structure and bonding in (NPCl2)3 and (NPCl2)4

(A) Inorganic polymers: introduction

Inorganic polymers are large molecules whose main chain (backbone) contains no carbon atoms, or contains atoms such as Si, P, N, S, B and O. They are generally more heat-resistant and less flammable than organic polymers. Important examples: silicones (Si–O backbone) and phosphonitrilic halides or phosphazenes (P=N backbone).

Silicones

Introduction

Silicones are organosilicon polymers with a repeating –R₂Si–O– unit (R = CH₃, C₂H₅, C₆H₅). The backbone is a chain of alternating Si and O atoms (siloxane linkage, Si–O–Si) with organic groups on silicon. Their empirical formula R₂SiO resembles a ketone R₂CO, so Kipping named them "silicones", but unlike ketones they do not have Si=O double bonds; they are polymers.

Nomenclature

• The basic unit –O–SiR₂– is called a siloxane unit; silicones are poly(siloxanes).
• The most common one, made from (CH₃)₂SiCl₂, is polydimethylsiloxane (PDMS), [–Si(CH₃)₂–O–]ₙ.
• Units are named by how many O atoms they share: M (mono-functional, R₃SiO–, chain end), D (di-functional, –R₂SiO–, chain unit), T (tri-functional, RSiO₃, cross-link) and Q (SiO₄).

Preparation

1. Making alkylchlorosilanes (Rochow direct process): Si + 2CH₃Cl → (CH₃)₂SiCl₂ (Cu catalyst, about 300 °C). CH₃SiCl₃ and (CH₃)₃SiCl also form and are separated by fractional distillation. Grignard method: SiCl₄ + 2CH₃MgCl → (CH₃)₂SiCl₂ + 2MgCl₂.
2. Hydrolysis: (CH₃)₂SiCl₂ + 2H₂O → (CH₃)₂Si(OH)₂ + 2HCl (dimethylsilanediol).
3. Condensation polymerisation: n (CH₃)₂Si(OH)₂ → [–Si(CH₃)₂–O–]ₙ + n H₂O.

• R₂SiCl₂ alone → long linear chains (or cyclic rings).
• Adding R₃SiCl → R₃SiOH caps the chain ends, controlling chain length (chain stopper).
• Adding RSiCl₃ → RSi(OH)₃ forms cross-links → 3-D network (resins).

Properties

• Thermally stable (Si–O bond is strong, about 452 kJ/mol); usable roughly from −50 °C to 250 °C.
• Water-repellent (hydrophobic) because the outer surface is covered with alkyl groups.
• Chemically inert; resist oxidation and most chemicals.
• Good electrical insulators; non-toxic; viscosity changes very little with temperature.

Uses

• Water-proofing fabrics, paper and masonry.
• Lubricants, greases, hydraulic fluids; insulators in electric motors.
• Sealants, gaskets, medical implants, cosmetics; antifoaming agents.

Silicone oils, elastomers and resins

(B) Phosphonitrilic halide polymers (phosphazenes)

Introduction

Phosphonitrilic halides (halophosphazenes) are compounds with repeating –N=PX₂– units, (NPX₂)ₙ, X = Cl, F, Br. They exist as cyclic trimers (n = 3), tetramers (n = 4), larger rings and long linear chains. Each P atom carries two halogens and is linked to two N atoms.

Preparation

Heating PCl₅ with ammonium chloride in a solvent such as sym-tetrachloroethane or chlorobenzene (about 120–130 °C):

n PCl₅ + n NH₄Cl → (NPCl₂)ₙ + 4n HCl

The product is a mixture, mainly the trimer (NPCl₂)₃ and tetramer (NPCl₂)₄, separated by fractional crystallisation or sublimation. On heating the trimer to about 250 °C, ring-opening polymerisation gives long chains of polydichlorophosphazene [NPCl₂]ₙ ("inorganic rubber").

Properties

• Trimer and tetramer are white crystalline solids, soluble in organic solvents.
• The P–Cl bonds are slowly hydrolysed by water; (NPCl₂)ₙ polymer turns brittle in moist air.
• Cl atoms can be replaced by nucleophiles: with NaOR → [NP(OR)₂]ₙ; with amines → [NP(NHR)₂]ₙ; with NaF → (NPF₂)ₙ. These substituted polymers are stable to water.
• Polydichlorophosphazene is elastic (rubber-like) and has a very low glass transition temperature.

Uses

• Flame-retardant materials and fire-resistant fibres (high P and N content).
• Low-temperature elastomers, gaskets and O-rings (fluoroalkoxy phosphazenes).
• Biomedical materials (biodegradable polymers for drug delivery).

Structure and bonding in (NPCl₂)₃ and (NPCl₂)₄

(NPCl₂)₃ — trimer

• Six-membered ring of alternating P and N atoms; the ring is almost planar.
• All P–N bond lengths are equal (about 158 pm), shorter than a P–N single bond (about 177 pm). So the π-electrons are delocalised over the ring, like benzene; it is drawn with alternate P=N bonds, but these are equivalent.
• Each P is roughly tetrahedral (sp³): two Cl atoms lie above and below the ring plane. Cl–P–Cl angle about 102°; N–P–N about 120°.

(NPCl₂)₄ — tetramer

• Eight-membered ring, non-planar: it adopts a puckered "tub" conformation.
• P–N bonds again equal and short (about 156 pm), showing delocalisation.

Bonding

1. Each P uses sp³ hybrids for σ bonds to two N and two Cl atoms. Each N uses sp² hybrids: two for σ bonds to P, one holds a lone pair in the ring plane.
2. The remaining unhybridised p orbital on each N (one electron) overlaps with d orbitals on P (one electron each) to form dπ–pπ bonds.
3. Because d orbitals can overlap with p orbitals on either side, the π system forms "islands" over three atoms (N–P–N) rather than a full ring current like benzene; therefore, the rings do not need to be planar and are not aromatic in the Hückel sense.
4. Modern view: the P–N bonding is largely ionic (P⁺–N⁻) plus negative hyperconjugation; but the dπ–pπ model is the one expected in B.Sc exams.

Likely exam questions

Long answer

1. What are silicones? Describe the preparation, properties and uses of silicones. Write a note on silicone oils, elastomers and resins.
2. How is phosphonitrilic chloride prepared? Discuss the structure and bonding in (NPCl₂)₃ and (NPCl₂)₄.

Short answer

1. What are inorganic polymers? Give two examples.
2. Why are silicones water-repellent and thermally stable?
3. Give any two uses of phosphazene polymers.