Chapter Notes›Class 12 Chemistry

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Class 12 ChemistryChapter Notes

10 chapters · Definitions, key points, formulas & exam tips · Updated 2025-26

Ch 1 Ch 2 Ch 3 Ch 4 Ch 5 Ch 6 Ch 7 Ch 8 Ch 9 Ch 10

Ch 1

Solutions

Key Definitions

Colligative Property: Property that depends only on number of solute particles, not their nature: VP lowering, BP elevation, FP depression, osmotic pressure.

Raoult's Law: Vapour pressure of a solution P = x_solvent × P°_solvent. Ideal solutions obey Raoult's law at all concentrations.

Van't Hoff Factor (i): Ratio of actual particles in solution to formula units dissolved. i > 1 for dissociation, i < 1 for association.

Key Points to Remember

→Vapour pressure lowering: ΔP/P° = x_solute (mole fraction of solute).
→Elevation in boiling point: ΔTb = Kb × m (Kb = ebullioscopic constant).
→Depression in freezing point: ΔTf = Kf × m (Kf = cryoscopic constant).
→Osmotic pressure: π = iMRT (van't Hoff equation).
→Abnormal molecular mass: electrolytes dissociate (i > 1); acetic acid in benzene associates (i < 1).
→Henry's law: solubility of gas ∝ partial pressure above solution. Used to explain carbonated beverages.

Formulas & Equations

ΔTb = Kb × m

ΔTf = Kf × m

π = iMRT

M = Kf × w/(ΔTf × W) (molecular mass by FP depression)

Exam Tips

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Osmotic pressure most accurate method for finding molecular mass of high molecular weight compounds (polymers, proteins).

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Non-volatile solute always raises BP and lowers FP.

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Van't Hoff factor for NaCl = 2, MgCl₂ = 3, K₂SO₄ = 3 (complete dissociation).

Ch 2

Electrochemistry

Key Definitions

Standard Electrode Potential: EMF measured at 1 M concentration, 1 atm pressure, 25°C relative to Standard Hydrogen Electrode (SHE = 0 V).

Molar Conductivity: Λm = κ × 1000/M. Conductivity per unit molar concentration. Unit: S·cm²·mol⁻¹.

Faraday's Laws: Mass deposited in electrolysis ∝ charge passed (1st law); equivalent masses deposited by same charge (2nd law).

Key Points to Remember

→Cell EMF: E°cell = E°cathode − E°anode (reduction potential of cathode minus anode).
→Nernst equation: E = E° − (RT/nF)lnQ = E° − 0.0592/n × log Q (at 25°C).
→Kohlrausch's law: Λ°m = λ°_cation + λ°_anion (infinite dilution, independent migration).
→Strong electrolytes: Λm increases with dilution (√c relationship).
→Weak electrolytes: degree of dissociation α = Λm/Λ°m.
→Lead storage battery: 12V, PbO₂ (+) and Pb (−) in H₂SO₄. Rechargeable.

Formulas & Equations

E°cell = E°cathode − E°anode

ΔG° = −nFE°cell

E = E° − 0.0592/n × log Q (at 25°C)

m = ZIt (Faraday's electrolysis law)

Exam Tips

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More positive E° = better oxidising agent (reduction favoured).

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At equilibrium: E = 0, so log K = nE°/0.0592.

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Electrolytic cell: electrical energy → chemical energy. Galvanic cell: chemical energy → electrical energy.

Ch 3

Chemical Kinetics

Key Definitions

Rate of Reaction: Change in concentration of reactant or product per unit time. Rate = −d[A]/dt = d[B]/dt.

Rate Law: Rate = k[A]^m[B]^n. Order = m + n (experimentally determined, not from stoichiometry).

Activation Energy: Minimum energy reactants must possess to form products. Ea in Arrhenius equation: k = Ae^(−Ea/RT).

Key Points to Remember

→Zero order: rate = k. [A] = [A]₀ − kt. t₁/₂ = [A]₀/2k.
→First order: rate = k[A]. [A] = [A]₀e^(−kt). t₁/₂ = 0.693/k (constant, independent of concentration).
→Molecularity: number of molecules taking part in elementary step. Always a whole number.
→Order: determined experimentally from rate data. Can be fraction or zero.
→Temperature coefficient: rate doubles for every 10°C rise (approximate).
→Arrhenius equation: k = Ae^(−Ea/RT). log(k₂/k₁) = Ea/2.303R × (T₂−T₁)/(T₁T₂).

Formulas & Equations

First order: ln[A] = ln[A]₀ − kt

t₁/₂ = 0.693/k (first order)

k = Ae^(−Ea/RT) (Arrhenius)

log(k₂/k₁) = Ea/(2.303R) × (T₂−T₁)/(T₁T₂)

Exam Tips

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If t₁/₂ is independent of concentration → first order reaction.

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Rate constant units: zero order: mol/L/s; first order: s⁻¹; second order: L/mol/s.

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Catalyst lowers Ea but does not change ΔH (products and reactants unchanged).

Ch 4

d and f Block Elements

Key Definitions

Transition Elements: d-block elements with partially filled d-orbitals in their atoms or common ions. Period 4: Sc to Zn.

Lanthanoid Contraction: Gradual decrease in atomic/ionic radius across lanthanoids due to poor shielding by 4f electrons. Consequence: Hf and Zr have same size.

Key Points to Remember

→Variable oxidation states: due to small energy difference between (n−1)d and ns electrons.
→Coloured ions: due to d-d transitions in presence of ligands.
→Catalytic activity: due to variable oxidation states and ability to adsorb reactants on surface.
→Magnetic properties: due to unpaired d electrons.
→K₂Cr₂O₇ (orange): oxidising agent in acidic medium. Cr₂O₇²⁻ → 2Cr³⁺.
→KMnO₄ (purple): strong oxidising agent. MnO₄⁻ → Mn²⁺ (acidic), Mn⁴⁺ (neutral), Mn⁶⁺ (basic).

Exam Tips

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Zn is NOT a transition element (d¹⁰ configuration, no d-d transitions, no variable oxidation state in common compounds).

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KMnO₄: used in acidic, neutral, and basic medium — products differ in each.

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Lanthanoids: +3 is most common oxidation state. Ce shows +4, Eu and Yb show +2.

Ch 5

Coordination Compounds

Key Definitions

Ligand: Molecule or ion that donates electron pairs to the central metal ion. Monodentate: 1 pair; bidentate: 2 pairs; polydentate: multiple pairs.

Coordination Number: Total number of ligand donor atoms directly bonded to the central metal atom/ion.

Crystal Field Theory: Electrostatic model explaining colour and magnetic properties. d-orbitals split into two sets in presence of ligands: t₂g and eg.

Key Points to Remember

→IUPAC naming: ligands (alphabetical) before metal; anion complex adds -ate; oxidation state of metal in Roman numerals.
→Effective Atomic Number (EAN): central atom achieves noble gas configuration.
→Werner's theory: primary valence (ionic) and secondary valence (coordinate bonds).
→VBT: hybridisation determines geometry. e.g., [Ni(CN)₄]²⁻: dsp², square planar.
→Optical isomerism: non-superimposable mirror images. Common in tris(bidentate) complexes.
→Chelates: ring-shaped complexes with polydentate ligands. More stable than monodentate complexes.

Formulas & Equations

CFSE for octahedral: t₂g = −0.4Δ₀, eg = +0.6Δ₀

Spectrochemical series: I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < en < CN⁻

Exam Tips

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Strong field ligands (CN⁻, en, CO): large Δ → low spin, inner orbital complex.

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Weak field ligands (Cl⁻, Br⁻): small Δ → high spin, outer orbital complex.

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IUPAC name order: anionic ligands first (with -o suffix), then neutral, then cationic.

Ch 6

Haloalkanes and Haloarenes

Key Definitions

SN2 Reaction: Bimolecular nucleophilic substitution. One-step, backside attack. Rate depends on both substrate and nucleophile. Inversion of configuration.

SN1 Reaction: Unimolecular nucleophilic substitution. Two-step: first carbocation formation. Rate depends only on substrate. Racemisation.

Optical Activity: Ability to rotate plane of polarised light. Due to chiral carbon (4 different groups). Enantiomers rotate light equally but in opposite directions.

Key Points to Remember

→SN2 favoured: primary alkyl halides, strong nucleophile, polar aprotic solvent.
→SN1 favoured: tertiary alkyl halides, weak nucleophile/polar protic solvent (stabilises carbocation).
→Reactivity order in SN2: CH₃X > 1° > 2° > 3°.
→Reactivity order in SN1: 3° > 2° > 1° > CH₃X (carbocation stability).
→Haloarenes: C-X bond has partial double bond character (resonance). Less reactive in substitution than haloalkanes.
→Uses: DDT (insecticide), freons (refrigerants — deplete ozone), iodoform (antiseptic).

Exam Tips

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Markovnikov's rule not applicable here — specific mechanisms (SN1/SN2) determine product.

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Grignard reagent: RMgX — reacts with aldehydes, ketones, esters to give alcohols.

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Freons: CCl₂F₂ and similar — release Cl atoms in stratosphere, catalyse ozone destruction.

Ch 7

Alcohols, Phenols and Ethers

Key Definitions

Lucas Test: Test to distinguish 1°, 2°, 3° alcohols using ZnCl₂/HCl. 3° reacts immediately (turbidity), 2° after 5 min, 1° no immediate reaction.

Dehydration: Removal of water from alcohol (using conc. H₂SO₄ at 170°C) to give alkene. At 140°C: gives ether. Follows Zaitsev rule.

Key Points to Remember

→Alcohols: hydrogen bonding gives higher BP than alkanes of similar MW.
→Acidity order: phenol > water > alcohol. Phenol more acidic due to resonance stabilisation of phenoxide ion.
→Esterification: alcohol + carboxylic acid ⇌ ester + water (Fischer esterification, reversible).
→Victor Meyer test: also distinguishes 1°, 2°, 3° alcohols by colour of nitrosoamine.
→Ethers: relatively unreactive. Cleavage with HI (stronger acid needed).
→Industrial ethanol: fermentation of sugars; cannot be used as fuel without denaturing.

Exam Tips

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Primary alcohol → aldehyde (mild oxidation) → carboxylic acid (strong oxidation).

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Secondary alcohol → ketone on oxidation.

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Tertiary alcohol: resistant to oxidation (no H on C-OH carbon).

Ch 8

Aldehydes, Ketones and Carboxylic Acids

Key Definitions

Nucleophilic Addition: Addition to C=O by a nucleophile attacking the electrophilic carbon. Aldehydes more reactive than ketones (less steric hindrance, less electron donation).

Aldol Condensation: Reaction of aldehyde/ketone with α-H with another carbonyl compound (same or different) in base to give β-hydroxy aldehyde/ketone, which dehydrates to give α,β-unsaturated compound.

Cannizzaro Reaction: Disproportionation of aldehydes with no α-H in presence of strong base. Half oxidised to acid, half reduced to alcohol.

Key Points to Remember

→Distinguishing test: Tollens' (silver mirror) and Fehling's — positive only for aldehydes, not ketones.
→Iodoform test: CH₃COR and CH₃OH give iodoform (CHI₃, yellow ppt) with I₂/NaOH.
→Acidic strength: carboxylic acid > carbonic acid > phenol > water.
→Electron-withdrawing groups increase acidity of carboxylic acids (e.g., Cl−CH₂COOH > CH₃COOH).
→Hell-Volhard-Zelinsky reaction: HVZ — halogenation of RCOOH at α-carbon using X₂/PX₃.
→Nucleophilic addition reactions: HCN, RMgX, NaHSO₃ (with aldehydes and methyl ketones).

Formulas & Equations

Tollen's test: Ag[(NH₃)₂]⁺ → Ag mirror (with aldehyde)

Iodoform: RCOCH₃ + I₂/NaOH → CHI₃ + RCOONa

Exam Tips

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Ketones give iodoform test only if they have CH₃CO− group (methyl ketones).

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Aldol condensation: requires α-hydrogen. Formaldehyde, benzaldehyde (no α-H) cannot self-condense.

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Carboxylic acid derivatives reactivity: acid chloride > anhydride > ester > amide.

Ch 9

Amines

Key Definitions

Basicity: Amines are Lewis bases — lone pair on N donates to proton. Aliphatic amines more basic than ammonia; arylamines less basic (lone pair delocalised into ring).

Diazonium Salt: R−N₂⁺X⁻. Prepared from primary amine + NaNO₂/HCl at 0–5°C. Important intermediate in aromatic synthesis.

Key Points to Remember

→Basicity order (in water): 2° aliphatic > 1° > NH₃ > 3° (in gas phase: 3° > 2° > 1°).
→Arylamines are weaker bases than alkylamines due to resonance of lone pair with aromatic ring.
→Hofmann bromamide reaction: RCONH₂ + Br₂/NaOH → RNH₂ (primary amine, 1 carbon less).
→Gabriel synthesis: gives only primary amines.
→Diazonium coupling: ArN₂⁺ + Ar'OH → azo dye (used in dyeing).
→Sandmeyer reaction: ArN₂⁺ + CuCl/CuBr → aryl chloride/bromide. Useful for introducing halogens.

Exam Tips

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Amine identification: primary reacts with HNO₂ to give nitrogen gas; secondary gives oil (N-nitrosamine); tertiary gives salt.

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Hofmann bromamide: product has one less carbon than starting amide.

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Diazonium salts: unstable, kept at 0–5°C. React with phenol, amine (coupling) and Cu salts (Sandmeyer).

Ch 10

Biomolecules

Key Definitions

Monosaccharide: Simplest carbohydrate that cannot be hydrolysed further. Glucose (aldohexose) and fructose (ketohexose) are most important.

Peptide Bond: CO-NH linkage formed between amino group of one amino acid and carboxyl group of another with elimination of water.

Nucleotide: Building block of nucleic acids. Contains: phosphate group + pentose sugar + nitrogenous base. DNA uses deoxyribose; RNA uses ribose.

Key Points to Remember

→Glucose: open chain formula C₆H₁₂O₆, aldohexose. Reacts with Fehling's, Tollens', bromine water.
→Sucrose (non-reducing): glucose + fructose. Lactose and maltose are reducing sugars.
→Proteins: primary (amino acid sequence), secondary (α-helix/β-sheet), tertiary (3D folding), quaternary (multi-subunit).
→Denaturation: protein loses its 3D structure due to pH, heat, or chemicals. Does not change primary structure.
→DNA: double helix (Watson-Crick). A pairs with T (2 H-bonds); G pairs with C (3 H-bonds).
→Vitamins: A, D, E, K (fat-soluble); B-complex and C (water-soluble).

Exam Tips

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Reducing sugar test: Tollens'/Fehling's test — glucose, maltose, lactose are reducing; sucrose is not.

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Enzyme specificity: one enzyme, one substrate — lock and key model.

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Amino acids have both NH₂ and COOH groups — can act as both acid and base (amphoteric).

Frequently Asked Questions

Are these notes based on 2025-26 CBSE syllabus for Class 12 Chemistry?

Yes. All chapter notes here are based on the latest 2025-26 CBSE syllabus for Class 12 Chemistry. Deleted topics are clearly marked so you focus only on what will be tested in your board exam.

How to study Class 12 Chemistry notes effectively for board exams?

Read each chapter's notes once to build understanding. Then close the notes and try to recall every key point, definition, and formula from memory. Anything you miss is your weak area — revisit only those points. This active recall method takes less time and retains far more than re-reading.

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Chapter notes contain detailed definitions, key terms, formulas, and concept breakdowns — they're for learning and understanding. Chapter summaries are shorter paragraph-style overviews — they're for quick revision. Use notes when you're studying a chapter for the first time; use summaries the night before the exam.

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