Maharashtra HSC Chemistry cheat sheet

Raoult's law for volatile solutes: p_total = p₁°x₁ + p₂°x₂; for a non-volatile solute p = p₁°x₁.

Relative lowering of vapour pressure: (p₁° − p)/p₁° = x₂ = n₂/(n₁+n₂), a colligative property independent of solute identity.

Elevation of boiling point ΔTb = Kb·m and depression of freezing point ΔTf = Kf·m, with m = molality (mol kg⁻¹ solvent).

Osmotic pressure π = CRT = (n/V)RT; for electrolytes multiply by van't Hoff factor i: π = iCRT.

van't Hoff factor i = observed colligative value / calculated value; i > 1 for dissociation (NaCl → 2), i < 1 for association (benzoic acid dimer in benzene → 0.5).

Henry's law: p = KH·x; solubility of gas falls as temperature rises and rises with pressure; KH increases with temperature.

Colligative property order for equal molality: osmotic pressure ≈ ΔTf > ΔTb > lowering of vapour pressure in practical sensitivity.

Ideal solution: ΔH_mix = 0, ΔV_mix = 0, obeys Raoult over whole range; positive deviation → weaker A–B forces, negative deviation → stronger A–B forces.

Molarity depends on temperature (volume changes); molality and mole fraction do not.

Ostwald dilution law for weak acid: Ka = Cα²/(1−α) ≈ Cα² when α is small; α = √(Ka/C).

pH = −log[H₃O⁺]; pOH = −log[OH⁻]; at 298 K pH + pOH = 14 and Kw = [H₃O⁺][OH⁻] = 1×10⁻¹⁴.

Buffer (Henderson–Hasselbalch): pH = pKa + log([salt]/[acid]); for basic buffer pOH = pKb + log([salt]/[base]).

Salt hydrolysis: strong acid–strong base salt pH = 7; weak acid–strong base pH = 7 + ½(pKa + log C); weak base–strong acid pH = 7 − ½(pKb + log C).

Solubility product Ksp: for AₓBᵧ, Ksp = xˣyʸs^(x+y); AgCl 1:1 gives Ksp = s².

Precipitation occurs when the ionic product exceeds Ksp; common-ion addition lowers solubility.

Acid–base theory: Arrhenius (H⁺/OH⁻ in water), Brønsted–Lowry (proton donor/acceptor), Lewis (electron-pair acceptor/donor, e.g. BF₃).

Buffer capacity is maximum when [salt]/[acid] = 1, i.e. pH = pKa; useful range pKa ± 1.

First law: ΔU = q + w; with pressure–volume work only, w = −p_extΔV; at constant volume qV = ΔU, at constant pressure qp = ΔH.

Enthalpy relation: ΔH = ΔU + Δn_gRT, where Δn_g = moles gaseous products − reactants.

Standard enthalpy of formation of an element in its standard state is zero; Hess's law allows ΔH to be summed over steps.

Entropy: ΔS = q_rev/T; ΔS_universe = ΔS_system + ΔS_surroundings > 0 for a spontaneous process.

Gibbs energy: ΔG = ΔH − TΔS; ΔG < 0 spontaneous, ΔG = 0 equilibrium, ΔG > 0 non-spontaneous (at constant T, p).

ΔG° = −RT ln K = −nFE°_cell; large negative ΔG° means large K.

Bond enthalpy: ΔH_reaction = Σ(bonds broken) − Σ(bonds formed); exothermic when bonds formed are stronger.

Calorimetry: q = mcΔT; bomb calorimeter gives ΔU (constant volume), coffee-cup gives ΔH.

State functions: U, H, S, G; path functions: q and w.

Nernst equation at 298 K: E = E° − (0.0592/n)·log Q; for a cell E_cell = E°_cell − (0.0592/n)·log([products]/[reactants]).

Standard cell potential: E°_cell = E°_cathode − E°_anode; positive E°_cell means spontaneous galvanic cell.

Relation to Gibbs energy: ΔG° = −nFE°; F = 96485 C mol⁻¹.

Conductance: G = 1/R; conductivity κ = G·(l/A); molar conductivity Λm = κ×1000/M (M in mol L⁻¹).

Kohlrausch law: Λ°m = ν₊λ°₊ + ν₋λ°₋; used to find Λ° of weak electrolytes and degree of dissociation α = Λm/Λ°m.

Faraday's laws: mass deposited m = (Q/F)(M/z) = (I·t·M)/(z·F).

Galvanic cell: anode oxidation (negative electrode), cathode reduction (positive electrode); electrons flow anode → cathode in external circuit.

Rusting requires both oxygen and moisture; prevention by galvanising, sacrificial anodes (Zn, Mg) and cathodic protection.

Strong electrolytes show linear Λm vs √C (Debye–Hückel–Onsager); weak electrolytes rise steeply at low concentration.

Rate = −Δ[reactant]/Δt = Δ[product]/Δt; average vs instantaneous rate; rate law determined experimentally, not from stoichiometry.

Order of reaction = sum of exponents in rate law; molecularity is a theoretical count of species in an elementary step.

First-order integrated rate law: k = (2.303/t)·log(a/(a−x)); half-life t½ = 0.693/k, independent of initial concentration.

Zero-order: [A] = [A]₀ − kt, t½ = [A]₀/2k; second-order: 1/[A] = 1/[A]₀ + kt, t½ = 1/(k[A]₀).

Arrhenius equation: k = A·e^(−Ea/RT); log(k₂/k₁) = (Ea/2.303R)(1/T₁ − 1/T₂).

Collision theory: rate depends on collision frequency, orientation, and fraction of molecules with energy ≥ Ea; catalyst lowers Ea.

Temperature coefficient ≈ 2–3 per 10 K rise for many reactions near room temperature.

Rate-determining step is the slowest elementary step; intermediates do not appear in the overall rate law.

Units of k: zero-order mol L⁻¹ s⁻¹; first-order s⁻¹; second-order L mol⁻¹ s⁻¹.

Group 16 (O, S, Se, Te, Po): oxygen shows −2 and −1 (peroxides); sulphur shows +4, +6; metallic character and acidic strength increase down the group.

Group 17 (F, Cl, Br, I, At): F only −1 (no d-orbitals); Cl/Br/I show +1, +3, +5, +7 in oxoacids.

Acid strength of oxoacids: HClO < HClO₂ < HClO₃ < HClO₄; more oxygen atoms increase acidity by stabilising the conjugate base.

Interhalogen compounds (ClF₃, BrF₅, ICl) form when halogen of lower electronegativity is central; larger halogens form higher fluorides.

Bleaching powder: Ca(OCl)Cl or Ca(OCl)₂·CaCl₂·Ca(OH)₂·2H₂O, made by Cl₂ + dry slaked lime; gives Cl₂ with dilute acid.

Group 18: monatomic gases, very low boiling points, chemically inert; xenon forms fluorides XeF₂, XeF₄, XeF₆ and oxyfluorides.

Anomalies of oxygen: small size, high electronegativity, absence of d-orbitals, hydrogen bonding in water, paramagnetism of O₂ (two unpaired electrons).

Sulphuric acid: contact process (S → SO₂ → SO₃ with V₂O₅ → oleum → H₂SO₄); dehydrating and oxidising agent.

Chlorine preparation in lab: MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O.

Transition element: partially filled d-orbitals in atom or common ions; Zn, Cd, Hg are d-block but not typical transition metals.

General properties: variable oxidation states, coloured ions (d–d transitions), paramagnetism, catalytic activity, complex formation, alloy formation.

Across a period: atomic radius decreases, density and ionisation enthalpy increase; ionic radii of M²⁺ decrease.

Lanthanoid contraction: steady decrease in size from La to Lu because 4f electrons poorly shield nuclear charge; causes similarity of Y³⁺ and lanthanoid ions.

Actinoids show larger oxidation-state range (up to +7 in Np, Pu) and greater radioactivity; 5f orbitals less shielded than 4f.

Interstitial compounds (TiC, Fe₃H, Mn₄N) are hard, high melting, retain metallic conductivity.

KMnO₄ in acidic medium (H₂SO₄) → Mn²⁺ (pink to colourless, 5-electron change); in neutral/alkaline → MnO₂ (brown).

K₂Cr₂O₇ in acidic medium → Cr³⁺ (green); used in volumetric analysis with diphenylamine indicator.

Coordination number = number of donor atoms bonded to the central metal; ligands classified by denticity (mono, bi, poly, chelating).

IUPAC naming: ligands alphabetically before metal; anionic complex ends in '-ate'; oxidation state of metal in Roman numerals in parentheses.

Werner theory: primary valence (ionisable) and secondary valence (non-ionisable, coordination number).

EAN rule: stable when effective atomic number = atomic number of next noble gas; e.g. [Co(NH₃)₆]³⁺ EAN 36.

Crystal field splitting: Δo (octahedral) > Δt (tetrahedral); strong-field ligands (CN⁻, CO, NO₂⁻) give low spin, weak-field (I⁻, Br⁻, F⁻, H₂O) give high spin.

Spectrochemical series: I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < NO₂⁻ < CN⁻ < CO.

Colour arises from d–d transition; high-spin octahedral d¹–d⁹ complexes are coloured; d⁰ and d¹⁰ are colourless.

Isomerism: ionisation, hydrate, linkage (NO₂⁻/ONO⁻), coordination, geometrical (cis/trans) and optical (d and l).

Crystal field stabilisation energy and magnetic moment: μ = √(n(n+2)) BM for n unpaired electrons.

Classification: alkyl (primary/secondary/tertiary), allylic, vinylic, aryl, benzylic halides; vinylic and aryl halides are unreactive to SN1/SN2 due to partial double-bond character.

SN1: two steps, carbocation intermediate, first-order, favoured by tertiary halides and polar protic solvents; racemisation.

SN2: one-step backside attack, second-order, favoured by primary halides and strong nucleophiles; inversion of configuration (Walden).

Reactivity order toward SN1: 3° > 2° > 1° > methyl (carbocation stability); toward SN2: methyl > 1° > 2° > 3° (steric hindrance).

Preparation: alcohol + HX/ZnCl₂ (Lucas reagent), PCl₅, SOCl₂ (gives pure alkyl chloride, by-product SO₂ and HCl gas).

Haloalkanes with alcoholic KOH give alkenes (dehydrohalogenation); with aqueous KOH give alcohols; with alcoholic KCN give nitriles.

Wurtz reaction: 2R–X + 2Na → R–R + 2NaX in dry ether; Finkelstein gives alkyl iodides using NaI in acetone.

Freons (CCl₂F₂) damage the ozone layer; chloramphenicol is an antibiotic containing halogen; DDT is a persistent organochlorine pesticide.

Grignard reagent RMgX: prepared in dry ether; reacts with water to give alkane, with CO₂ to give carboxylic acid, with carbonyl to give alcohol.

Classification: mono/di/trihydric alcohols; phenols have –OH directly on aromatic ring; ethers R–O–R′.

Lucas test: tertiary alcohol turns cloudy immediately, secondary in ~5 min, primary not at room temperature (with conc. HCl/ZnCl₂).

Dehydration of alcohols with conc. H₂SO₄ at 443 K gives alkenes; order of ease 3° > 2° > 1°; Saytzeff rule gives more substituted alkene.

Phenol is acidic because phenoxide ion is resonance-stabilised; pKa ≈ 10, weaker than carboxylic acids but stronger than alcohols.

Phenol reactions: with Br₂ water gives 2,4,6-tribromophenol (white ppt), with FeCl₃ gives violet colour, with NaOH gives sodium phenoxide.

Reimer–Tiemann: phenol + CHCl₃ + NaOH → salicylaldehyde; Kolbe: phenol + CO₂/NaOH → salicylic acid.

Williamson synthesis: R–ONa + R′–X → R–O–R′; primary halides give good yields; tertiary give alkenes (elimination).

Ethers are relatively inert, cleaved by HI (Ziesel) giving alcohol + alkyl iodide; lower ethers are flammable and form peroxides on storage.

Boiling point order for equal molecular mass: alcohol > ether > alkane (hydrogen bonding in alcohols).

Aldehydes are easily oxidised; Tollens' reagent gives silver mirror with aldehydes only; Fehling's gives red Cu₂O with aliphatic aldehydes.

2,4-DNP (Brady's reagent) gives orange/yellow precipitate with aldehydes and ketones; iodoform test (I₂/NaOH) is positive for CH₃CO– and CH₃CH(OH)– groups.

Aldol condensation needs α-hydrogen; Cannizzaro reaction (disproportionation) occurs with aldehydes lacking α-hydrogen, e.g. HCHO, benzaldehyde.

Clemmensen (Zn-Hg/HCl) and Wolff–Kishner (NH₂NH₂/KOH) reduce C=O to CH₂; NaBH₄ reduces only C=O, not C=C.

Nucleophilic addition order: HCHO > RCHO > R₂CO (steric and electronic factors); HCN gives cyanohydrin.

Carboxylic acids: acidity order HCOOH > CH₃COOH > CH₃CH₂COOH; electron-withdrawing groups (Cl, NO₂) increase acidity.

Acid derivatives reactivity: acid chloride > anhydride > ester > amide toward nucleophilic substitution; esters hydrolyse to acid + alcohol.

HVZ reaction: carboxylic acid with α-H + Br₂/P gives α-bromocarboxylic acid; used in amino-acid synthesis.

Ketones resist oxidation; Baeyer's reagent (alkaline KMnO₄) decolourises with aldehydes and unsaturation.

Classification: primary, secondary, tertiary (by number of H replaced), and quaternary ammonium salts; aromatic amines have NH₂ on ring.

Basicity order in aqueous solution: 2° > 1° > 3° > NH₃ for aliphatic amines; aromatic amines are much weaker bases (lone pair delocalised).

Hinsberg test: benzenesulphonyl chloride distinguishes 1° (soluble product), 2° (insoluble), 3° (no reaction).

Carbylamine test: CHCl₃ + KOH gives foul-smelling isocyanide with primary amines only.

Hofmann bromamide degradation: RCONH₂ + Br₂/NaOH → RNH₂ with one carbon less; used to prepare pure primary amines.

Gabriel phthalimide synthesis gives pure primary amines and fails for aromatic amines.

Diazonium salts: ArN₂⁺Cl⁻ from ArNH₂ + NaNO₂ + HCl at 273–278 K; coupling with phenol gives azo dyes.

Sandmeyer reaction: ArN₂⁺ + CuCl/HCl → ArCl; also gives ArBr (CuBr), ArCN (CuCN/KCN), ArOH (H₂O, warm).

Amines are less acidic than alcohols; nitration of aniline gives meta product because protonation of –NH₂ forms meta-directing –NH₃⁺.

Carbohydrates: monosaccharides (glucose, fructose), disaccharides (sucrose, lactose, maltose), polysaccharides (starch, cellulose, glycogen).

Glucose is an aldohexose; fructose is a ketohexose; sucrose is non-reducing (no free anomeric OH), maltose and lactose are reducing.

Proteins are polymers of α-amino acids joined by peptide bonds; primary, secondary (α-helix, β-sheet), tertiary, quaternary structure.

Denaturation destroys secondary/tertiary structure without breaking peptide bonds; caused by heat, pH change, urea.

Enzymes are biological catalysts, mostly proteins, highly specific; lock-and-key and induced-fit models.

Vitamins: fat-soluble A, D, E, K; water-soluble B-complex and C; deficiency diseases: scurvy (C), beri-beri (B₁), night blindness (A), rickets (D).

Nucleic acids: DNA deoxyribose, A-T, G-C, double helix; RNA ribose, A-U; DNA replication and protein synthesis.

Hormones: insulin (polypeptide) lowers blood glucose; deficiency causes diabetes mellitus.

Addition polymers: polyethene, PVC, polystyrene, Teflon (PTFE), from monomers with C=C; no by-product.

Condensation polymers: nylon-6,6 (adipic acid + hexamethylenediamine), terylene/PET (ethylene glycol + terephthalic acid), bakelite (phenol + formaldehyde).

Natural polymers: cellulose, starch, proteins, natural rubber (cis-polyisoprene); synthetic: PVC, nylon, terylene.

Vulcanisation: natural rubber + sulphur + heat gives cross-links, increasing strength and elasticity.

Copolymerisation: two or more different monomers, e.g. Buna-S (butadiene + styrene), Buna-N (butadiene + acrylonitrile).

Thermoplastic polymers soften on heating (PVC, polyethene); thermosetting polymers do not (bakelite, melamine).

Biodegradable polymers: PHBV, Nylon-2-nylon-6, used to reduce plastic pollution.

Degree of polymerisation = number of repeating units; molecular mass of polymer = DP × monomer mass.

Green chemistry aims to design chemical products and processes that reduce or eliminate hazardous substances.

Twelve principles include atom economy, prevention of waste, safer solvents, energy efficiency, use of renewable feedstocks, catalysis, and design for degradation.

Atom economy = (mass of desired product / total mass of reactants) × 100%; higher atom economy means less waste.

Green alternatives: supercritical CO₂ as solvent, H₂O₂ as oxidant, solid acid catalysts, solvent-free reactions, microwave-assisted synthesis.

Nanochemistry: study of materials at 1–100 nm scale, where surface area to volume ratio becomes very large.

Nanoparticles show size-dependent properties: gold nanoparticles appear red, TiO₂ becomes more reactive for photocatalysis, carbon nanotubes are strong and conductive.

Synthesis routes: sol–gel, chemical vapour deposition, laser ablation, and green synthesis using plant extracts.

Crystal systems: 7 crystal systems and 14 Bravais lattices; unit cell described by a, b, c and α, β, γ.

Schottky defect is common in ionic crystals with similar cation/anion size (NaCl); Frenkel defect in those with large size difference (AgCl, ZnS).

Frenkel defect keeps density unchanged; Schottky lowers density.

Amorphous solids: glass, rubber, plastics, short-range order, isotropic, no sharp melting point.

Crystalline solids: long-range order, anisotropic, sharp melting point.

Coordination number of NaCl (rock salt) is 6:6; CsCl is 8:8; ZnS (zinc blende) is 4:4; CaF₂ is 8:4.

Number of octahedral voids per atom = 1; tetrahedral voids per atom = 2.

Magnetic properties: paramagnetic (unpaired electrons), diamagnetic (paired), ferromagnetic (Fe, Co, Ni), antiferromagnetic (MnO), ferrimagnetic (Fe₃O₄).

Concentration units: molarity M = moles solute/L solution; molality m = moles solute/kg solvent; mole fraction x; ppm; normality.

For dilute aqueous solutions at 25 °C, molarity ≈ molality (density ≈ 1 g mL⁻¹).

Kb and Kf are solvent constants; for water Kb = 0.512 K kg mol⁻¹, Kf = 1.86 K kg mol⁻¹.

Osmotic pressure is preferred for macromolecules (proteins) because it is measurable at low concentration.

Isotonic solutions have equal osmotic pressure; 0.9% NaCl (normal saline) is isotonic with blood.

Reverse osmosis desalinates seawater by applying pressure greater than osmotic pressure.

Solubility of solid in liquid usually increases with temperature; solubility of gas decreases with temperature.

Positive deviation examples: ethanol + acetone, ethanol + water; negative deviation: chloroform + acetone, HNO₃ + water.

Azeotropes: minimum boiling (positive deviation, e.g. ethanol-water 95%) and maximum boiling (negative deviation, e.g. HNO₃-water 68%).

Degree of dissociation α = (Λm / Λ°m) for weak electrolytes; for weak acid α = √(Ka/C).

pH of very dilute strong acid needs water contribution; [H⁺]² − Ka[H⁺] − Kw = 0 for weak acid.

Buffer examples: CH₃COOH/CH₃COONa (acidic, pH ~4.7), NH₃/NH₄Cl (basic, pH ~9.25).

Salt of weak acid and weak base: pH = 7 + ½(pKa − pKb).

Amphiprotic salt (NaHCO₃): pH = ½(pKa₁ + pKa₂).

Ksp of AgCl ≈ 1.8×10⁻¹⁰ at 298 K; molar solubility s ≈ 1.34×10⁻⁵ M.

Lewis acids: BF₃, AlCl₃, FeCl₃, SO₃; Lewis bases: NH₃, H₂O, F⁻.

The pH of pure water at 298 K is 7; at higher temperature Kw increases and pH falls below 7 while remaining neutral.

System types: open (matter and energy exchange), closed (energy only), isolated (neither).

Intensive properties (T, p, density) independent of amount; extensive (U, H, S, G, V) depend on amount.

Standard conditions: 298 K, 1 bar; standard enthalpy of combustion of graphite = −393.5 kJ mol⁻¹.

Reversible isothermal expansion of ideal gas: w = −nRT ln(V₂/V₁).

Adiabatic process: q = 0, so ΔU = w; for ideal gas reversible adiabatic, TV^(γ−1) = constant.

Second law: ΔS_total ≥ 0; for reversible process ΔS_total = 0.

Third law: entropy of a perfect crystal at 0 K is zero.

Trouton's rule: ΔH_vap/Tb ≈ 85–88 J K⁻¹ mol⁻¹ for many liquids.

Hess's law application: ΔH_f°(reaction) = Σ ΔH_f°(products) − Σ ΔH_f°(reactants).

Standard hydrogen electrode (SHE) assigned E° = 0.00 V at 298 K, 1 bar H₂, 1 M H⁺.

Electrochemical series: Li⁺/Li most negative (−3.04 V); F₂/F⁻ most positive (+2.87 V).

A more negative E° means stronger reducing agent; a more positive E° means stronger oxidising agent.

Electrolytic cell: electrical energy → chemical energy; anode is positive, cathode is negative.

Galvanic cell: chemical energy → electrical energy; anode is negative, cathode is positive.

Lead storage battery: anode Pb, cathode PbO₂, electrolyte dilute H₂SO₄; ~2 V per cell.

Fuel cell: H₂–O₂ cell gives ~1.23 V, water is the only product, high efficiency.

Conductivity decreases with dilution for strong electrolytes; molar conductivity increases with dilution.

Corrosion of iron is an electrochemical process; anodic reaction Fe → Fe²⁺ + 2e⁻.

Pseudo-first-order reaction: one reactant in large excess, e.g. acid hydrolysis of ester; rate depends on ester only.

Rate law for 2A → products, second order: rate = k[A]²; integrated 1/[A] = 1/[A]₀ + kt.

For a first-order reaction, 99.9% completion takes ~10 half-lives.

Fraction remaining after n half-lives = (1/2)ⁿ; fraction reacted = 1 − (1/2)ⁿ.

Graphical tests: zero-order [A] vs t linear; first-order ln[A] vs t linear; second-order 1/[A] vs t linear.

Rate constant k depends on temperature (Arrhenius), not on concentration.

Activation energy is the minimum energy needed for a collision to lead to reaction; catalysts lower it.

Photochemical reactions: rate depends on intensity and wavelength of light; H₂ + Cl₂ → 2HCl is a chain reaction.

Order can be zero, fractional, or negative; molecularity is always a positive integer.

Ozone: bent molecule, diamagnetic, powerful oxidant; prepared by Siemen's ozoniser (silent electric discharge in O₂).

Ozone depletion by CFCs: Cl radical catalyses O₃ → O₂; one Cl atom destroys many O₃ molecules.

Allotropes of sulphur: rhombic (α, stable below 369 K) and monoclinic (β, stable above 369 K).

Oxides of sulphur: SO₂ (acid rain precursor, reducing), SO₃ (strong acid anhydride, oxidising).

Halogen colour and state at room temp: F₂ pale yellow gas, Cl₂ greenish-yellow gas, Br₂ red-brown liquid, I₂ violet solid.

Bromine and iodine are extracted from sea water/brine; iodine gives blue colour with starch.

Noble gases: He used in balloons and cryogenics, Ne in signs, Ar in welding, Kr in flash lamps, Xe in anaesthesia.

Clathrate compounds: noble gases trapped in cavities of host lattices, e.g. Xe in ice.

Halogen acids: HF weak (H-bonding), HCl, HBr, HI strong; acid strength HF < HCl < HBr < HI.

Electronic configuration: Cr 3d⁵4s¹ and Cu 3d¹⁰4s¹ are exceptions due to extra stability of half-filled/full d.

Oxidation states: Mn shows +2 to +7; Cr +2 to +6; Fe +2, +3; Cu +1, +2.

Colour: Sc³⁺, Ti⁴⁺, Zn²⁺ are colourless (d⁰ or d¹⁰); Cu²⁺ blue, Fe³⁺ yellow, Ni²⁺ green.

Catalytic uses: Fe (Haber), V₂O₅ (contact), Ni (hydrogenation), Pt/Rh (catalytic converters).

Alloy examples: brass (Cu-Zn), bronze (Cu-Sn), steel (Fe-C), stainless steel (Fe-Cr-Ni).

Lanthanoid oxidation state: mostly +3; Ce⁴⁺ and Eu²⁺ are notable exceptions.

Ce⁴⁺/Ce³⁺ redox couple is used in analytical chemistry (cerimetry).

KMnO₄ preparation: pyrolusite MnO₂ fused with KOH and KNO₃ to give K₂MnO₄, then electrolytic oxidation to KMnO₄.

K₂Cr₂O₇ preparation: chromite FeCr₂O₄ fused with Na₂CO₃ and roasted to Na₂CrO₄, then acidified to Na₂Cr₂O₇.

Common coordination numbers: 2 (linear, Ag⁺), 4 (tetrahedral or square planar, Pt²⁺), 6 (octahedral, Co³⁺, Fe³⁺).

Geometrical shapes: CN 2 linear, CN 4 tetrahedral or square planar, CN 6 octahedral.

Chelate effect: polydentate ligands (EDTA, en) form more stable complexes than monodentate analogues.

EDTA is hexadentate; used in complexometric titration and to treat heavy-metal poisoning.

Cisplatin, cis-[Pt(NH₃)₂Cl₂], is an anticancer drug; trans isomer is inactive.

Haemoglobin: Fe²⁺ in porphyrin ring binds O₂; chlorophyll: Mg²⁺ porphyrin complex.

Vitamin B₁₂ is a Co(III) complex (cyanocobalamin).

Amminetric titration uses [Ag(NH₃)₂]⁺; silver halides dissolve in excess NH₃ (AgCl fully, AgBr partially, AgI not).

Hybridisation from VBT: CN 4 tetrahedral sp³; CN 4 square planar dsp²; CN 6 octahedral d²sp³ or sp³d².

Alkyl halides are polar; boiling point order for same alkyl group: RI > RBr > RCl > RF (higher molar mass).

Boiling point for isomeric halides: 1° > 2° > 3° (branching lowers boiling point).

Nucleophilic substitution competes with elimination; strong bulky base favours elimination.

Elimination follows Saytzeff rule: more substituted alkene is major product.

Alkyl fluorides prepared by Swarts reaction (AgF, Hg₂F₂, CoF₂) or by treating alcohol with KF.

Chloroform (CHCl₃) is stored in dark brown bottles with ethanol to prevent phosgene formation.

Chloroform + conc. HNO₃ gives chloropicrin (CCl₃NO₂), used as fumigant.

DDT: p,p′-dichlorodiphenyltrichloroethane, insecticide, banned due to persistence and biomagnification.

Haloarenes give substitution (not nucleophilic substitution) because C–X bond has partial double-bond character.

Preparation of alcohols: hydration of alkenes, reduction of carbonyls (NaBH₄, LiAlH₄), Grignard addition, fermentation of sugars.

Primary alcohol + Cu at 573 K gives aldehyde; secondary gives ketone; tertiary gives alkene (dehydration).

Oxidation of primary alcohol with PCC stops at aldehyde; with acidic KMnO₄ goes to carboxylic acid.

Phenol preparation: cumene process (benzene + propene → cumene → phenol + acetone); Dow process (chlorobenzene + NaOH).

Phenol with dilute HNO₃ gives o- and p-nitrophenol; with conc. HNO₃ gives picric acid (2,4,6-trinitrophenol).

Ethers: Williamson synthesis uses alkoxide + primary alkyl halide; intramolecular version gives cyclic ethers.

Alcohols and phenols both react with Na to give H₂; only phenols react with NaOH and FeCl₃.

Ethers do not react with Na or NaOH; they are neutral to litmus.

Ether cleavage with HI: aryl-alkyl ethers give phenol + alkyl iodide (I attacks alkyl side).

Preparation of aldehydes: oxidation of primary alcohols (PCC), ozonolysis of alkenes, Rosenmund reduction (RCOCl + H₂/Pd-BaSO₄).

Preparation of ketones: oxidation of secondary alcohols, Friedel–Crafts acylation, hydration of alkynes.

Aldehydes: formaldehyde, acetaldehyde, benzaldehyde; ketones: acetone, acetophenone, benzophenone.

Formalin is 40% aqueous formaldehyde, used as preservative and disinfectant.

Acetone is a common laboratory solvent; miscible with water.

Iodoform test positive for acetaldehyde, acetone, ethanol, isopropanol.

Carboxylic acids: formic acid is both aldehyde and acid; acetic acid is in vinegar (4–6%).

Acid strength: trichloroacetic acid > dichloroacetic > monochloroacetic > acetic acid.

Benzoic acid is a food preservative; oxalic acid is used in cleaning; lactic acid in curd.

Preparation: reduction of nitro compounds (Sn/HCl or H₂/Pd), reduction of nitriles (LiAlH₄), Hofmann bromamide, Gabriel synthesis.

Amines are basic; basicity in gas phase: 3° > 2° > 1° > NH₃ (no solvation effect).

Aromatic amines: aniline is weaker base than cyclohexylamine; p-nitroaniline is weaker than aniline; p-toluidine is stronger.

Coupling reaction: diazonium salt + phenol (pH 9–10) or aniline (pH 4–5) gives azo dye.

Diazonium salt + H₃PO₂ gives arene (deamination); + H₂O warm gives phenol; + KI gives aryl iodide.

Acylation of amines with acid chloride gives amide; benzoylation with benzoyl chloride gives benzamide.

Reaction with nitrous acid: 1° aliphatic gives N₂ gas; 1° aromatic gives diazonium salt; 2° gives N-nitrosoamine; 3° gives nitrite salt.

Amines are less acidic than alcohols; they form salts with acids (e.g. anilinium chloride).

Quaternary ammonium salts are used as cationic surfactants and antiseptics (e.g. cetyltrimethylammonium bromide).

Glucose: C₆H₁₂O₆, six carbons, aldehyde group at C1, exists in α- and β-pyranose forms; mutarotation.

Glucose reactions: Tollens' positive, Fehling's positive, osazone formation, acetylation to pentaacetate.

Sucrose hydrolysis gives glucose + fructose (invert sugar); enzyme invertase.

Starch: amylose (linear α-1,4) and amylopectin (branched α-1,4 and α-1,6); iodine gives blue colour.

Cellulose: β-1,4-glycosidic bonds, not digested by humans; provides dietary fibre.

Proteins: amino acids, peptide bond, zwitterion at isoelectric point.

Enzyme examples: amylase (starch), pepsin (protein), lipase (fats), catalase (H₂O₂).

Vitamins: A (retinol), B₁ (thiamine), B₂ (riboflavin), B₆ (pyridoxine), B₁₂ (cobalamin), C (ascorbic acid), D (calciferol), E (tocopherol), K (phylloquinone).

Nucleotides: sugar + phosphate + nitrogenous base; DNA has deoxyribose, RNA has ribose.

Monomer vs repeating unit: monomer is the starting molecule; repeating unit is the part in the polymer chain derived from it.

Addition polymerisation of ethene requires initiator (peroxide) and high pressure; gives HDPE or LDPE depending on conditions.

LDPE: branched, soft, low density (carry bags); HDPE: linear, rigid, high density (bottles).

PVC: rigid, used in pipes; plasticised PVC is flexible (wires, upholstery).

Bakelite: phenol + formaldehyde, thermosetting, used in electrical switches and handles.

Melamine-formaldehyde resin: hard, heat resistant, used in dinnerware.

Natural rubber: cis-1,4-polyisoprene; gutta-percha: trans isomer.

Buna-S (SBR): butadiene + styrene, synthetic rubber for tyres; Buna-N: butadiene + acrylonitrile, oil resistant.

Teflon (PTFE): chemically inert, non-stick, high melting point.

Atom economy vs yield: a reaction can have high yield but low atom economy if by-products are large.

Example: addition reactions have 100% atom economy; substitution reactions have lower atom economy.

Green solvents: water, supercritical CO₂, ionic liquids, ethanol; avoid chlorinated solvents.

Catalysis is a key green principle: lower energy, higher selectivity, less waste.

Renewable feedstocks: biomass, glucose, vegetable oils; replacing petroleum-derived starting materials.

Design for degradation: biodegradable polymers (PHBV), avoid persistent organohalogens.

Nanoscale materials: fullerenes (C₆₀), carbon nanotubes, graphene, quantum dots, metal nanoparticles.

Characterisation of nanoparticles: TEM, SEM, XRD, UV-visible spectroscopy.

Applications: drug delivery, targeted therapy, water purification, sensors, catalysis, sunscreens (TiO₂, ZnO nanoparticles).