Tamil Nadu HSE Second Year Chemistry cheat sheet

Concentration: froth flotation suits sulphide ores (pine oil frother, collectors make ore hydrophobic); gravity/levigation for oxide ores; magnetic separation for magnetic gangue like FeWO4.

Roasting = heating ore strongly in excess air (sulphides to oxides, e.g. 2ZnS + 3O2 → 2ZnO + 2SO2). Calcination = heating in limited air or absence of air (carbonates/hydrated oxides, e.g. CaCO3 → CaO + CO2).

Ellingham diagrams plot ΔG° vs T for oxide formation; the curve lowest on the diagram belongs to the most stable oxide. A metal can reduce the oxide of any metal whose curve lies above its own.

At the point where two curves cross, ΔG° = 0 for the coupled reduction, so both metal and oxide coexist at that temperature (e.g. Mg/MgO crossing Al/Al2O3 near 1350 °C).

Carbon reduction is used below ~710 °C; above ~710 °C carbon's curve lies above CO-forming 2C + O2 → 2CO, so CO becomes the effective reducing agent in the blast furnace.

Flux + gangue = slag. In iron extraction CaCO3 decomposes to CaO, which combines with SiO2 gangue to give CaSiO3 (slag), floating on molten iron.

Aluminium: Hall–Héroult electrolysis of Al2O3 dissolved in molten cryolite (Na3AlF6) with CaF2; cryolite lowers the melting point and raises conductivity; carbon anodes are consumed.

Ultrapure metals: van Arkel–de Boer process (Mond-like iodide route) and zone refining exploit the greater solubility of impurities in the melt than in the solid.

Froth flotation, cyanide leaching (macArthur–Forrest for Au/Ag, forming soluble complex), and Mond process (Ni + 4CO → Ni(CO)4 at 50–60 °C, decomposed near 180 °C) are classic separation/refining routes.

Group 15: N, P, As, Sb, Bi. Nitrogen is a diatomic gas (N≡N, 941 kJ/mol, very inert); the rest are solids. Basicity of hydrides falls down the group: NH3 > PH3 > AsH3 > SbH3 > BiH3.

Nitrogen: NH3 is pyramidal (sp3, lone pair), a Lewis base; HNO3 is a strong oxidising acid; N2O (nitrous oxide) is a mild anaesthetic, NO is paramagnetic, NO2 is brown and dimerises to N2O4.

Phosphorus allotropes: white P (P4 tetrahedra, discrete, reactive, soluble in CS2), red P (polymeric, less reactive), black P (thermodynamically most stable). PCl5 is trigonal bipyramidal but ionic [PCl4]+[PCl6]− in the solid state.

Group 16: O, S, Se, Te, Po. Oxygen is a gas; sulphur forms S8 crowns; acidic strength of hydrides increases down the group (H2O < H2S < H2Se < H2Te).

SO2 is bent (sp2, lone pair), SO3 is planar trigonal; H2SO4 is a strong dibasic acid and strong oxidiser. Contact process: S → SO2 → SO3 (V2O5 catalyst) → oleum → H2SO4.

Ozone is bent (bond angle ~116.8°), diamagnetic, and oxidises iodide to iodine (used in quantitative estimation); it is thermodynamically unstable relative to O2.

Group 17: F2, Cl2 (gases), Br2 (liquid), I2 (solid). Oxidising power: F2 > Cl2 > Br2 > I2; electron gain enthalpy of F is less negative than Cl (small size, electron–electron repulsion).

Halogen hydrides: acidic strength HF < HCl < HBr < HI; HF is weak due to strong H-bonding and high bond enthalpy. Interhalogens (ClF3, BrF5, IF7) form when the heavier halogen is central.

Group 18: noble gases. Xe forms fluorides (XeF2 linear, XeF4 square planar, XeF6 distorted octahedral); XeF4 + 2H2O → Xe + O2 + 4HF; helium has the lowest boiling point of any element.

A transition element has an incomplete d subshell in its atom or in a common oxidation state (so Zn, Cd, Hg are excluded). General configuration (n−1)d^(1–10) ns^(1–2).

High melting/boiling points, hardness, and density arise from strong metallic bonding involving delocalised d electrons; Cr has the highest melting point of 3d series (excepting W in 5d).

Variable oxidation states occur because 3d and 4s energies are close; Mn shows +2 to +7; Cr shows +2 to +6. Highest state in 3d: Ti +4, V +5, Cr +6, Mn +7.

Colour arises from d–d transitions (partially filled d); Sc3+, Ti4+, Zn2+ are colourless. Paramagnetism is from unpaired d electrons (spin-only moment μ = √(n(n+2)) BM).

Mn3+ and Co3+ are strong oxidisers; Cr2+ is a strong reducing agent; Fe3+/Fe2+ interconvert readily. Lanthanoid contraction: ionic radii decrease from La to Lu due to poor 4f shielding.

Lanthanoid +3 is the most common state; Ce4+ and Eu2+ are notable exceptions (Ce4+ is a good oxidiser). Lanthanoids are used in phosphors, catalysts, and magnets.

Actinoids (5f) show more variable oxidation states than lanthanoids because 5f, 6d, 7s energies are close; all actinoids are radioactive.

Interstitial compounds (e.g. TiC, Fe3C) are hard, high-melting, retain metallic conductivity. Alloy formation is favoured by similar atomic radii (e.g. Cr, Mo, W).

Potassium dichromate: orange, K2Cr2O7; in acid it is a strong oxidiser (Cr2O72− + 14H+ + 6e− → 2Cr3+ + 7H2O, E° = 1.33 V). Potassium permanganate: purple, MnO4−; in acid MnO4− + 8H+ + 5e− → Mn2+ + 4H2O.

Coordination number = number of donor atoms bonded to the central metal. Common geometries: CN 2 linear, CN 4 tetrahedral or square planar, CN 6 octahedral.

Ligand types: monodentate (NH3, Cl−), bidentate (en, oxalate C2O42−), polydentate/chelating (EDTA4− is hexadentate), ambidentate (NO2− as nitro or nitrito, SCN− as thiocyanato or isothiocyanato).

Nomenclature: ligands alphabetically before the metal (ignoring prefixes); anionic complex ends in -ate (e.g. [Cu(CN)4]3− = tetracyanocuprate(II)); prefixes di-, tri-, tetra-, penta-, hexa- for repeated ligands.

Werner's theory: primary valency = ionisable, satisfied by anions; secondary valency = non-ionisable, fixed by coordination number and directional (gives geometry).

EAN (effective atomic number) rule: EAN = Z − oxidation state + 2 × CN; stable when EAN = next noble gas (e.g. [Co(NH3)6]3+: 27 − 3 + 12 = 36).

Isomerism: ionisation, hydrate, linkage, coordination, and stereo (geometric cis/trans, optical). [Co(NH3)4Cl2]+ exists as cis (violet) and trans (green) isomers.

CFT: in an octahedral field d orbitals split into t2g (lower) and eg (higher) with Δo; tetrahedral splitting Δt ≈ (4/9)Δo. Strong-field ligands (CN−, CO, NO2−) give low-spin; weak-field (I−, Br−, F−) give high-spin.

Spectrochemical series (increasing field): I− < Br− < S2− < SCN− < Cl− < F− < OH− < H2O < NH3 < en < NO2− < CN− < CO.

Magnetic moment (spin-only): μ = √(n(n+2)) BM; [Fe(CN)6]4− (low-spin d6, n=0) is diamagnetic while [Fe(H2O)6]2+ (high-spin d6, n=4) is paramagnetic with μ ≈ 4.9 BM.

Unit cells: simple cubic (1 atom, r = a/2), body-centred cubic (2 atoms, r = √3a/4), face-centred cubic (4 atoms, r = √2a/4). Packing efficiency: SC 52.4%, BCC 68%, FCC/HCP 74%.

Coordination numbers: SC 6, BCC 8, FCC/HCP 12. For FCC, a = 2√2 r; for BCC, a = 4r/√3.

Density relation: ρ = (Z × M) / (a³ × NA), where Z = atoms per unit cell, a in cm, M in g/mol, NA = 6.022 × 10^23.

Defects: Schottky (missing cation+anion pair, lowers density, favoured by high CN and similar ion sizes, e.g. NaCl); Frenkel (ion displaced to interstitial, density unchanged, e.g. AgCl).

Metal excess defect: anion vacancy with trapped electron gives F-centre colour (NaCl yellow, KCl violet). Metal deficiency arises from variable valency (FeO with Fe3+).

Bragg's law: nλ = 2d sinθ; X-ray diffraction determines interplanar spacing d and hence lattice parameters.

Packing: HCP = ABAB, FCC/CCP = ABCABC. Both have 74% efficiency; the difference lies in layer stacking sequence.

Radius ratio rules: r+/r− 0.225–0.414 tetrahedral (CN 4), 0.414–0.732 octahedral (CN 6), 0.732–1.0 cubic (CN 8).

Point defects change physical properties; ionic solids conduct only via defects (Schottky/Frenkel) and increase with temperature. Semiconductors: n-type (P, As donors), p-type (B, Al acceptors).

First law: ΔU = q + w; with work done on the system w = −PΔV. At constant volume qv = ΔU; at constant pressure qp = ΔH = ΔU + ΔnRT.

Enthalpy of formation is defined per mole of compound formed from elements in their standard states; ΔH°f of an element in its standard state is zero.

Hess's law: ΔH of a reaction is independent of path; use ΔH°rxn = Σ ΔH°f(products) − Σ ΔH°f(reactants).

Bond enthalpy relation: ΔHrxn = Σ bond enthalpies broken − Σ bond enthalpies formed (positive means endothermic).

Entropy: ΔS = qrev/T. ΔS°rxn = Σ S°(products) − Σ S°(reactants). Second law: ΔS_universe = ΔS_system + ΔS_surroundings ≥ 0.

Gibbs free energy: ΔG = ΔH − TΔS. Spontaneous when ΔG < 0; at equilibrium ΔG = 0, so ΔH = TΔS. ΔG° = −RT ln K = −2.303 RT log K.

Standard free energy of formation: ΔG°f = 0 for elements in their standard states. ΔG°rxn = Σ ΔG°f(products) − Σ ΔG°f(reactants).

Third law: entropy of a perfectly crystalline substance is zero at 0 K; absolute entropies are always positive for real substances.

Internal energy is a state function; heat and work are path functions. ΔU = qv and ΔH = qp are the calorimetric working equations.

Rate = −d[A]/dt for reactants, +d[B]/dt for products; average rate over a time interval, instantaneous rate from the tangent to the concentration–time curve.

Rate law: rate = k[A]^m[B]^n; order = m + n (experimental), molecularity is a theoretical count for an elementary step. Order can be zero, fractional, or negative; molecularity cannot.

Zero order: [A] = [A]0 − kt; t1/2 = [A]0/2k. First order: ln([A]0/[A]) = kt; t1/2 = 0.693/k (independent of initial concentration).

Second order (single reactant): 1/[A] − 1/[A]0 = kt; t1/2 = 1/(k[A]0).

Arrhenius equation: k = A e^(−Ea/RT); ln(k2/k1) = (Ea/R)(1/T1 − 1/T2); a plot of ln k vs 1/T gives slope −Ea/R and intercept ln A.

Collision theory: rate depends on collision frequency, orientation, and fraction with energy ≥ Ea. Catalyst lowers Ea without changing ΔH or equilibrium position.

Pseudo-first-order example: acid hydrolysis of ester with water in large excess; order in ester is one, overall pseudo-first-order.

Rate-determining step is the slowest elementary step; the overall rate law reflects it. For a multistep mechanism, sum of steps gives the stoichiometric equation.

Half-life for first order is constant; for zero and second order it depends on initial concentration. Units of k: zero order mol L−1 s−1; first order s−1; second order L mol−1 s−1.

Ostwald dilution law for weak electrolyte: Ka = Cα²/(1 − α); for small α, α = √(Ka/C). Degree of dissociation increases on dilution.

pH = −log[H3O+]; pOH = −log[OH−]; pH + pOH = 14 at 25 °C (Kw = 1 × 10^−14).

Weak acid: [H+] = √(Ka·C); weak base: [OH−] = √(Kb·C). pH of a salt of weak acid and strong base = 7 + ½(pKa + log C).

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

Solubility product: Ksp = [A+]^m[B−]^n; for a sparingly soluble salt, molar solubility s relates to Ksp (AB: Ksp = s²; AB2: Ksp = 4s³).

Common ion effect: adding a common ion decreases solubility; precipitation occurs when the ionic product exceeds Ksp.

Lewis acids accept electron pairs (BF3, AlCl3); Lewis bases donate them (NH3, F−). Bronsted acids donate protons; conjugate base is the species left after proton loss.

Buffer range is roughly pKa ± 1; a mixture of weak acid and its salt resists pH change on adding small amounts of strong acid or base.

Salt hydrolysis: strong acid–strong base salt is neutral; strong acid–weak base salt is acidic; weak acid–strong base salt is basic; weak acid–weak base salt depends on relative Ka and Kb.

Oxidation occurs at the anode, reduction at the cathode. In a galvanic cell the anode is negative (electrons leave); in electrolysis the anode is positive.

Standard cell potential: E°cell = E°cathode − E°anode (reduction potentials). A positive E°cell means a spontaneous reaction under standard conditions.

Nernst equation at 25 °C: E = E° − (0.0591/n) log Q; for a concentration cell with identical electrodes E° = 0 and E depends only on the ion concentration ratio.

Relation to free energy: ΔG° = −nFE°cell; and ΔG° = −RT ln K, so E°cell = (0.0591/n) log K at 25 °C.

Kohlrausch's law: limiting molar conductivity of an electrolyte is the sum of the limiting molar conductivities of its ions; used to find Λ°m of weak electrolytes.

Conductivity: κ = (1/R)(l/A); molar conductivity Λm = κ × 1000 / M (M in mol/L). Λm increases with dilution; for strong electrolytes it follows Λm = Λ°m − A√C.

Faraday's laws: mass deposited m = (Q/F)(M/z) = (It/F)(M/z). One faraday = 96,485 C/mol of electrons.

Lead–acid cell: anode Pb, cathode PbO2, electrolyte dilute H2SO4; discharge consumes H2SO4 and forms PbSO4 at both electrodes; it is rechargeable.

Corrosion (rusting) is an electrochemical process; iron acts as anode (Fe → Fe2+ + 2e−), oxygen reduction at cathodic regions; galvanising with Zn protects by sacrificial action.

Physisorption: weak van der Waals forces, low enthalpy (~20–40 kJ/mol), reversible, multilayer possible. Chemisorption: chemical bonds, high enthalpy (~80–240 kJ/mol), irreversible, monolayer only.

Freundlich isotherm: x/m = k P^(1/n) (0 < 1/n < 1); log(x/m) vs log P is linear with slope 1/n and intercept log k. Applies to physical adsorption.

Langmuir isotherm assumes monolayer adsorption on identical sites; x/m = aP/(1 + bP).

Catalysis: a catalyst lowers activation energy, increases rate, does not change equilibrium constant or ΔH. Homogeneous (same phase) vs heterogeneous (different phase, surface reaction).

Enzyme catalysis: lock-and-key / induced-fit; enzymes are highly specific and sensitive to pH and temperature. Example: invertase hydrolyses sucrose.

Colloids: particle size 1–1000 nm. Lyophilic (solvent-loving, reversible, e.g. starch) vs lyophobic (solvent-repelling, irreversible, e.g. gold sol).

Tyndall effect = scattering of light by colloidal particles; Brownian motion = random motion; electrophoresis = movement of charged particles in an electric field.

Hardy–Schulze rule: coagulating power of an ion increases with its charge; Al3+ > Ca2+ > Na+ for negatively charged sols.

Emulsions: oil-in-water (milk) and water-in-oil (butter); emulsifiers stabilise them. Gold number = minimum mg of protective colloid needed to prevent coagulation of 10 mL gold sol by 1 mL of 10% NaCl.

Nomenclature: haloalkane = alkyl halide (CH3CH2Cl = chloroethane); haloarene = halogen on an aromatic ring (C6H5Cl = chlorobenzene). Classify as primary, secondary, tertiary by the carbon bearing the halogen.

SN1: two steps via carbocation, rate = k[substrate], favoured by tertiary substrate, polar protic solvent, good leaving group; gives racemisation (with inversion). SN2: one step, rate = k[substrate][nucleophile], favoured by primary substrate, strong nucleophile, inversion of configuration.

Leaving-group ability: I− > Br− > Cl− > F− (weaker C–X bond and greater polarisability of larger halides). C–F bond is strongest, so fluorides react slowly.

Haloarenes resist nucleophilic substitution because of partial double-bond character (resonance) of C–X and sp2 carbon; they undergo substitution only under drastic conditions (Dow process: chlorobenzene + NaOH at 623 K, 300 atm → phenol).

Reactivity toward SN: alkyl > allyl/benzyl > vinyl/aryl. Allyl and benzyl halides are reactive because the resulting carbocation is resonance-stabilised.

Elimination (dehydrohalogenation): alcoholic KOH gives alkene; Saytzeff rule - the more substituted (more stable) alkene is the major product.

Grignard reagent: RX + Mg in dry ether → RMgX; reacts with water to give alkane, with aldehydes to give secondary alcohols, with ketones to give tertiary alcohols.

Haloarenes undergo electrophilic substitution (nitration, halogenation) because halogen is ortho/para-directing and mildly deactivating. Wurtz reaction: 2RX + 2Na → R–R + 2NaX.

Chloroform (CHCl3) + KOH gives isocyanide with primary amines; stored with ethanol to prevent phosgene formation; used as an anaesthetic historically.

Amines: primary RNH2, secondary R2NH, tertiary R3N. Basicity in water: secondary > primary > tertiary for alkylamines (solvation and steric effects); ammonia is weaker than alkylamines.

Aromatic amines are weaker bases than aliphatic amines because the lone pair on N is delocalised into the ring (aniline pKb ≈ 9.4).

Amine preparation: reduction of nitro compounds (Sn/HCl), Hofmann bromamide degradation (RCONH2 + Br2 + 4KOH → RNH2, one carbon fewer), Gabriel phthalimide synthesis (gives pure primary amines).

Amines react with nitrous acid: primary aliphatic → alcohol + N2; primary aromatic → diazonium salt (0–5 °C); secondary → N-nitrosamine (yellow oil); tertiary → no reaction (aromatic tertiary gives nitrosamine).

Diazonium salts: ArN2+Cl− prepared at 273–278 K; coupling with phenol gives azo dye (orange-red); Sandmeyer reaction with CuCl/HCl gives chlorobenzene; with H2O gives phenol.

Carbylamine test: primary amine + CHCl3 + KOH → isocyanide (foul smell). Hinsberg test uses benzenesulphonyl chloride to distinguish primary, secondary, tertiary amines.

Carbohydrates: monosaccharides (glucose, fructose), disaccharides (sucrose, maltose, lactose), polysaccharides (starch, cellulose, glycogen). Sucrose is non-reducing; maltose and lactose are reducing.

Glucose: open chain has aldehyde at C1 and hydroxyls; exists mainly as cyclic hemiacetal (α/β pyranose). Oxidation with Br2 water gives gluconic acid; with HNO3 gives saccharic acid.

Proteins: amino acids linked by peptide bonds; primary, secondary (α-helix, β-pleated sheet), tertiary, quaternary structure. Denaturation destroys secondary and tertiary structure (e.g. boiling egg white).

Nucleic acids: DNA (deoxyribose, bases A, T, G, C) and RNA (ribose, U replaces T). Nucleotide = base + sugar + phosphate; the double helix is held by hydrogen bonds (A–T two, G–C three).