CBSE Chemistry Study Guide

Overview

This guide covers the CBSE Class 11 and 12 Chemistry syllabus (NCERT). It is divided into three branches — Physical, Inorganic, and Organic Chemistry — with key concepts, reactions, and exam-focused advice.

The CBSE Class 12 Chemistry board exam carries 70 marks (theory) + 30 marks (practical). The theory paper is divided roughly into Physical Chemistry (23 marks), Inorganic Chemistry (19 marks), and Organic Chemistry (28 marks).

1. Physical Chemistry

1.1 Atomic Structure

Bohr model: Electrons orbit in fixed energy levels. $E_n = -13.6/n^2\;\text{eV}$ (hydrogen).

Quantum numbers:

Quantum number	Symbol	Values	Describes
Principal	$n$	$1, 2, 3, \ldots$	Energy level / shell
Azimuthal	$l$	$0, 1, \ldots, n-1$	Subshell (s, p, d, f)
Magnetic	$m_l$	$-l, \ldots, 0, \ldots, l$	Orbital orientation
Spin	$m_s$	$+1/2, -1/2$	Electron spin

Electronic configurations follow the Aufbau principle, Hund’s rule, and Pauli exclusion principle.

Exceptions: $\text{Cr} = [\text{Ar}]\,3d^5\,4s^1$ ; $\text{Cu} = [\text{Ar}]\,3d^{10}\,4s^1$ (half-filled and fully-filled $d$ -subshells are more stable).

1.2 Chemical Bonding

Ionic bond: Transfer of electrons; electrostatic attraction between cations and anions. Lattice energy $U \propto \frac{z_+ z_-}{r_+ + r_-}$ .

Covalent bond: Sharing of electron pairs.

VSEPR theory: Electron pair repulsion determines molecular geometry. Lone pairs repel more than bonding pairs.

Molecular orbital theory (MOT): Combines atomic orbitals to form bonding ( $\sigma$ , $\pi$ ) and antibonding ( $\sigma^*$ , $\pi^*$ ) orbitals. Bond order $= \frac{1}{2}(n_b - n_a)$ . Bond order > 0 means the molecule is stable.

Hybridisation:

Hybridisation	Geometry	Bond angle	Example
$sp$	Linear	$180°$	$\text{BeCl}_2$ , $\text{C}_2\text{H}_2$
$sp^2$	Trigonal planar	$120°$	$\text{BF}_3$ , $\text{C}_2\text{H}_4$
$sp^3$	Tetrahedral	$109.5°$	$\text{CH}_4$ , $\text{NH}_3$
$sp^3d$	Trigonal bipyramidal	$90°, 120°$	$\text{PCl}_5$
$sp^3d^2$	Octahedral	$90°$	$\text{SF}_6$

Hydrogen bonding: Strong dipole-dipole interaction involving H bonded to F, O, or N. Explains high boiling points of water, HF, NH $_3$ .

1.3 Thermodynamics

System types: Open (exchanges matter and energy), closed (energy only), isolated (neither).

Functions:

Internal energy ( $U$ ): Sum of all molecular energies. State function.
Enthalpy ( $H = U + PV$ ): Heat content at constant pressure.
Entropy ( $S$ ): Measure of disorder.
Gibbs free energy ( $G = H - TS$ ): Determines spontaneity.

Key relations:

$\Delta G = \Delta H - T\Delta S$

$\Delta G < 0$ : spontaneous
$\Delta G = 0$ : at equilibrium
$\Delta G > 0$ : non-spontaneous

Hess’s law: The total enthalpy change is the same regardless of the pathway.

1.4 Equilibrium

Law of mass action: For $aA + bB \rightleftharpoons cC + dD$ :

$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

Le Chatelier’s principle: If a system at equilibrium is disturbed, it shifts to oppose the change.

$K_p$ and $K_c$ : $K_p = K_c(RT)^{\Delta n}$ where $\Delta n = (\text{moles of gas products}) - (\text{moles of gas reactants})$ .

Acid-base:

pH $= -\log[\text{H}^+]$
$K_a$ (acid dissociation constant); $K_b$ (base dissociation constant)
$K_w = [\text{H}^+][\text{OH}^-] = 10^{-14}$ at 25°C
Buffer solution: resists pH change; Henderson-Hasselbalch: $\text{pH} = \text{p}K_a + \log\frac{[\text{A}^-]}{[\text{HA}]}$

Solubility product: $K_{sp} = [\text{A}^+]^m[\text{B}^-]^n$ for $A_mB_n \rightleftharpoons m\text{A}^+ + n\text{B}^-$ .

1.5 Electrochemistry

Standard electrode potential: $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$

Nernst equation: $E = E^\circ - \frac{RT}{nF}\ln Q$

At 298 K: $E = E^\circ - \frac{0.0591}{n}\log Q$

Faraday’s laws:

$m = Z \cdot Q = \frac{M}{nF} \cdot It$
Same charge liberates equivalent amounts of different substances.

Conductivity: $\kappa = \frac{1}{\rho}$ ; molar conductivity $\Lambda_m = \frac{\kappa}{c}$ .

1.6 Chemical Kinetics

Rate law: For $aA + bB \to \text{products}$ , rate $= k[A]^m[B]^n$ .

Order of reaction: Sum of exponents in the rate law.

Integrated rate laws (first order): $[A] = [A]_0 e^{-kt}$ or $k = \frac{2.303}{t}\log\frac{[A]_0}{[A]}$ .

Half-life (first order): $t_{1/2} = \frac{0.693}{k}$

Arrhenius equation: $k = A e^{-E_a/RT}$ ; $\ln\frac{k_2}{k_1} = \frac{E_a}{R}\left(\frac{1}{T_1} - \frac{1}{T_2}\right)$

Collision theory: Rate depends on collision frequency, orientation factor, and fraction of molecules with energy $\geq E_a$ .

1.7 Solutions

Raoult’s law: $P = x_A P_A^\circ$ for ideal solutions.

Colligative properties (depend on number of solute particles, not identity):

Relative lowering of vapour pressure: $\frac{\Delta P}{P^\circ} = x_2$ (mole fraction of solute)
Elevation in boiling point: $\Delta T_b = K_b \cdot m$
Depression in freezing point: $\Delta T_f = K_f \cdot m$
Osmotic pressure: $\Pi = cRT$ (where $c$ is molarity)

Van’t Hoff factor: $i = \frac{\text{experimental value}}{\text{theoretical value}}$ ; accounts for dissociation ( $i > 1$ ) or association ( $i < 1$ ).

1.8 Surface Chemistry

Adsorption: Accumulation on a surface. Physisorption (weak, van der Waals) vs. chemisorption (strong, chemical bond).

Catalysis: Homogeneous vs. heterogeneous. Catalyst lowers activation energy without being consumed.

Colloids: Particle size 1—1000 nm. Tyndall effect, Brownian motion, electrophoresis, coagulation.

2. Inorganic Chemistry

2.1 Periodic Table

Periodic trends:

Property	Across a period	Down a group
Atomic radius	Decreases	Increases
Ionisation energy	Increases	Decreases
Electron affinity	Generally increases	Generally decreases
Electronegativity	Increases	Decreases

Shielding effect: Inner electrons shield outer electrons from the nuclear charge.

Diagonal relationship: Li-Mg, Be-Al, B-Si show similar properties due to comparable charge density.

2.2 s-Block Elements

Group 1 (Alkali metals): Very reactive, $+1$ oxidation state, low ionisation energy, form ionic compounds. React vigorously with water. Flame test colours: Li (crimson), Na (yellow), K (lilac).

Group 2 (Alkaline earth metals): $+2$ oxidation state, less reactive than Group 1. Be is amphoteric.

Anomalous properties of Li and Be are due to their small size, high ionisation energy, and high polarising power.

2.3 p-Block Elements

Group 13 (B, Al, Ga, In, Tl): Boron is a metalloid; others are metals. $\text{Al}_2\text{O}_3$ is amphoteric. Important compounds: borax, boric acid, alum.

Group 14 (C, Si, Ge, Sn, Pb): Catenation (C > Si > Ge). CO $_2$ is linear and acidic; SiO $_2$ is a 3D network solid. Sn and Pb show inert pair effect (+2 oxidation state preferred over +4).

Group 15 (N, P, As, Sb, Bi): Nitrogen exists as N $_2$ (triple bond). Important oxides of nitrogen (NO, NO $_2$ , N $_2$ O $_4$ , N $_2$ O). Phosphorus allotropes: white, red, black. Oxyacids of phosphorus. Bi shows inert pair effect.

Group 16 (O, S, Se, Te, Po): Oxygen is diatomic (O $_2$ ) and highly electronegative. Sulfur exists as S $_8$ crowns. SO $_2$ and SO $_3$ are important. Oxyacids of sulfur: $\text{H}_2\text{SO}_4$ (contact process).

Group 17 (Halogens): F, Cl, Br, I. Oxidising agents, $-1$ oxidation state. Bleaching action of Cl $_2$ . Interhalogen compounds: XX’ $n$ where $n = 1, 3, 5, 7$ .

Group 18 (Noble gases): Complete octets, chemically inert under normal conditions. Xe forms compounds with F and O ( $\text{XeF}_2$ , $\text{XeF}_4$ , $\text{XeF}_6$ , $\text{XeOF}_4$ ).

2.4 d-Block and f-Block Elements

d-Block (transition metals): Partially filled $d$ -orbitals. Variable oxidation states, coloured compounds, paramagnetic behaviour, catalytic activity, complex formation.

Important transition metals:

Mn: shows all oxidation states from $+2$ to $+7$ ; $\text{KMnO}_4$ is a strong oxidising agent
Fe: $\text{Fe}^{2+}$ (green) and $\text{Fe}^{3+}$ (yellow/brown); $\text{Fe}_2\text{O}_3$ (thermite)
Cu: $\text{Cu}^{2+}$ (blue in solution); $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$
Zn, Cd, Hg: $d^{10}$ configuration; not true transition metals (no partially filled $d$ -orbitals in common oxidation states)

Lanthanide contraction: Gradual decrease in atomic/ionic radii from La to Lu due to poor shielding of 4f electrons.

f-Block elements: Lanthanides (4f) and actinides (5f). Many are radioactive.

2.5 Coordination Compounds

Werner’s theory: Central metal ion is bound to ligands via coordinate bonds (dative covalent bonds).

Ligand types: Monodentate (e.g., Cl $^-$ , NH $_3$ , H $_2$ O), bidentate (e.g., ethylenediamine), polydentate (EDTA).

IUPAC nomenclature: Name ligands alphabetically, then metal, then oxidation state in Roman numerals. Prefixes for number of ligands: di-, tri-, tetra-, penta-, hexa-.

Isomerism in coordination compounds:

Structural: linkage, ionisation, hydrate isomers
Stereoisomerism: geometric (cis-trans) and optical

Crystal field theory: Ligands split $d$ -orbitals. In octahedral field: $t_{2g}$ (lower) and $e_g$ (higher). Crystal field splitting energy $\Delta_o$ . Weak field (high spin) vs. strong field (low spin).

3. Organic Chemistry

3.1 IUPAC Nomenclature

Identify the longest carbon chain (parent hydrocarbon).
Number the chain to give substituents the lowest possible locants.
Name and number substituents alphabetically (ignoring prefixes like di-, tri-).
For functional groups, assign the principal functional group the lowest number; name others as prefixes.

Functional group priority (highest to lowest for suffix): $-$ COOH $>$ $-$ CHO $>$ $-$ C=O $>$ $-$ OH $>$ $-$ NH_2 $>$ $>$ C=C $>$ $-$ C≡C.

3.2 Hydrocarbons

Alkanes ( $\text{C}_n\text{H}_{2n+2}$ ): Single bonds, tetrahedral geometry. Reactions: combustion, halogenation (free radical substitution).

Alkenes ( $\text{C}_n\text{H}_{2n}$ ): C=C double bond, planar geometry. Reactions: addition (HX, X $_2$ , H $_2$ O), Markovnikov’s rule, anti-Markovnikov (peroxide effect/Kharasch effect), ozonolysis.

Alkynes ( $\text{C}_n\text{H}_{2n-2}$ ): C≡C triple bond, linear geometry. Reactions: addition (similar to alkenes but in two steps), acetylide formation.

Aromatic hydrocarbons: Benzene ring with delocalised $\pi$ -electrons (6 $\pi$ ). Electrophilic aromatic substitution: nitration, sulphonation, halogenation, Friedel-Crafts alkylation/acylation. Activating groups ( $-$ OH, $-$ NH $_2$ , $-$ CH $_3$ ) and deactivating groups ( $-$ NO $_2$ , $-$ COOH, $-$ CHO).

3.3 Halogen Derivatives

Classification: $1°$ , $2°$ , $3°$ based on the carbon bearing the halogen.

Nucleophilic substitution:

SN1: Two-step, rate $= k[\text{R-X}]$ , carbocation intermediate, racemisation, favoured for $3°$ halides
SN2: One-step, rate $= k[\text{R-X}][\text{Nu}^-]$ , Walden inversion, favoured for $1°$ halides

Elimination (E1, E2): Forms alkenes. Saytzeff rule: more substituted alkene is major product.

3.4 Oxygen-Containing Compounds

Alcohols ( $-$ OH): Hydrogen bonding, soluble in water. Oxidation: $1° \to$ aldehyde $\to$ carboxylic acid; $2° \to$ ketone; $3°$ resistant. Dehydration: forms alkenes (conc. H $_2$ SO $_4$ , 170°C).

Phenols: Aromatic $-$ OH. More acidic than alcohols due to resonance stabilisation of phenoxide ion. Reactions: electrophilic substitution (ortho-para directing), Kolbe’s reaction, Reimer-Tiemann reaction.

Ethers: $\text{R-O-R}'$ . Relatively inert. Williamson’s synthesis: $\text{R-X} + \text{NaOR}' \to \text{R-O-R}'$ .

Aldehydes and ketones: Polar C=O bond. Nucleophilic addition reactions (Grignard reagent, NaBH $_4$ , HCN, NH $_3$ derivatives). Aldol condensation, Cannizzaro reaction, Clemmensen reduction, Wolff-Kishner reduction.

Carboxylic acids: Hydrogen bonding, weak acids ( $\text{p}K_a \approx 4$ — $5$ ). Reactions: esterification (with alcohols), decarboxylation, Hell-Volhard-Zelinsky reaction, reaction with $\text{PCl}_3/\text{PCl}_5/\text{SOCl}_2$ .

3.5 Nitrogen-Containing Compounds

Amines ( $-$ NH $_2$ ): Classified as $1°$ , $2°$ , $3°$ (and quaternary ammonium salts). Basic nature due to lone pair on N. Order of basicity in gas phase: $3° > 2° > 1° > \text{NH}_3$ . In aqueous solution: $2° > 3° > 1° > \text{NH}_3$ (steric + solvation effects). Carbylamine reaction (isocyanide test) for $1°$ amines.

** Diazonium salts:** $\text{ArN}_2^+\text{Cl}^-$ formed by nitrous acid with aromatic $1°$ amines. Used for Sandmeyer reactions to introduce $-$ Cl, $-$ Br, $-$ CN, $-$ OH groups on the aromatic ring.

3.6 Polymers

Classification:

Addition (chain-growth): Polyethylene, PVC, polystyrene, Teflon
Condensation (step-growth): Nylon-6,6, Bakelite, Dacron
Biodegradable: PHBV, PLA

Key terms: Monomer, polymer, degree of polymerisation, tacticity (atactic, isotactic, syndiotactic).

3.7 Biomolecules

Carbohydrates: Polyhydroxy aldehydes/ketones. Classified as monosaccharides (glucose, fructose), disaccharides (sucrose, lactose, maltose), polysaccharides (starch, cellulose, glycogen).

Proteins: Polymers of amino acids joined by peptide bonds ( $-$ CO $-$ NH $-$ ). Primary, secondary ( $\alpha$ -helix, $\beta$ -pleated sheet), tertiary, and quaternary structure. Denaturation: loss of biological activity due to disruption of secondary/tertiary structure.

Lipids: Fats (saturated) and oils (unsaturated). Phospholipids in cell membranes. Steroids (cholesterol).

Nucleic acids: DNA (deoxyribonucleic acid) and RNA (ribonucleic acid). Bases: A, T/U, G, C. Watson-Crick base pairing: A—T (2 H-bonds), G—C (3 H-bonds).

Vitamins: Fat-soluble (A, D, E, K) and water-soluble (B-complex, C). Deficiency diseases.

4. Key Definitions

Term	Definition
Enthalpy of combustion	Heat released when 1 mol of a substance burns completely in O $_2$
Enthalpy of formation	Heat change when 1 mol of a compound forms from its elements in standard states
Entropy	Measure of randomness or disorder of a system
Electrode potential	Tendency of an electrode to lose or gain electrons
Molar conductivity	Conductance of a solution containing 1 mol of solute between electrodes 1 cm apart
Order of reaction	Sum of exponents in the rate law; experimentally determined
Activation energy	Minimum energy required for a reaction to proceed
Osmotic pressure	External pressure required to prevent osmosis across a semipermeable membrane
Coordination number	Number of donor atoms bonded to the central metal ion in a complex
Chelation	Formation of a ring structure by a multidentate ligand around a metal ion
Markovnikov’s rule	In HX addition to alkenes, H adds to the carbon with more H atoms
Enantiomers	Non-superimposable mirror image stereoisomers
Carbocation	Positively charged carbon with three bonds and an empty $p$ -orbital
Tautomerism	Structural isomerism involving the migration of a proton and a double bond
Inert pair effect	Reluctance of $ns^2$ electrons in heavier p-block elements to participate in bonding

5. Exam Tips

Balance every chemical equation. Unbalanced equations lose marks. Check atoms and charge on both sides, especially for redox reactions.
Use proper arrow-pushing mechanisms. In organic chemistry, curved arrows show electron movement. Incomplete mechanisms are heavily penalised.
Memorise named reactions. Know the reagents, conditions, and products for Cannizzaro, Aldol, Friedel-Crafts, Sandmeyer, Williamson’s, and Hell-Volhard-Zelinsky reactions.
Write the IUPAC name systematically. Number the chain, identify the principal functional group, and name substituents in alphabetical order. Even if the name is wrong, correct numbering earns partial credit.
Draw resonance structures where applicable. This is expected for benzene, phenol, carboxylate ions, and conjugated systems.
State the geometry of hybridised orbitals. Questions frequently ask for the geometry and bond angle of a specific molecule — know the hybridisation-table relationships.
Know the trends and exceptions. CBSE regularly tests lanthanide contraction, diagonal relationships, inert pair effect, and anomalous behaviour of Li, Be, and O.

Common Pitfalls

Confusing oxidation state and coordination number. Oxidation state is the charge on the central metal; coordination number is the number of ligand donor atoms bonded to it.
Incorrect arrow-pushing in organic mechanisms. Arrows must start from a lone pair or bond and point to an electron-deficient centre. Double-headed arrows for equilibrium, single for movement.
Wrong priority order in nomenclature. Not giving the principal functional group the lowest locant, or misordering substituents alphabetically.
Forgetting to account for dissociation in colligative properties. For electrolytes like NaCl ( $i = 2$ ), $\Delta T_b = iK_bm$ , not just $K_bm$ .
Mixing up SN1 and SN2 mechanisms. SN1 is two-step with a carbocation (favoured for $3°$ ); SN2 is one-step with inversion (favoured for $1°$ ). Substrates and conditions matter.
Incorrect electronic configuration of transition metals. Remember Cr ( $3d^5 4s^1$ ) and Cu ( $3d^{10} 4s^1$ ) exceptions. Also remember that $4s$ fills before $3d$ but is emptied first upon ionisation.
Not distinguishing between physisorption and chemisorption. Physisorption is weak, reversible, and non-specific; chemisorption is strong, often irreversible, and highly specific. Both may occur together at different temperature ranges.

Worked Examples

Example 1: Calculating pH and Buffer Capacity

Problem: A buffer solution is prepared by mixing 0.1 mol CH3COOH (pKa = 4.74) and 0.1 mol CH3COONa in 1 L of solution. Calculate the pH. Then find the pH after adding 0.01 mol HCl. Solution: Using Henderson-Hasselbalch: pH = pKa + log([A-]/[HA]) = 4.74 + log(0.1/0.1) = 4.74. After adding 0.01 mol HCl, it reacts with CH3COO-: new [A-] = 0.09, new [HA] = 0.11. pH = 4.74 + log(0.09/0.11) = 4.74 + log(0.818) = 4.74 - 0.087 = 4.65.

Example 2: Determining Oxidation States in Redox

Problem: Balance the redox reaction: Cr2O7^2- + Fe^2+ + H+ -> Cr^3+ + Fe^3+ + H2O. Solution: Cr goes from +6 to +3 (gain 3e- per Cr atom, 2 Cr atoms = 6e- gained). Fe goes from +2 to +3 (lose 1e- per Fe). LCM of electrons: 6. Multiply Fe^2+ by 6. Balancing atoms: Cr2O7^2- + 6Fe^2+ + 14H+ -> 2Cr^3+ + 6Fe^3+ + 7H2O. Check: charge: left = -2 + 12 + 14 = +24, right = 6 + 18 = +24. Balanced.

Example 3: First-Order Kinetics Half-Life

Problem: A first-order reaction has a rate constant k = 0.0693 min^-1. What fraction of the reactant remains after 20 minutes? Solution: t_1/2 = 0.693/0.0693 = 10 min. In 20 minutes, that is 2 half-lives. Fraction remaining = (1/2)^2 = 0.25, or 25%. Alternatively: [A] = [A]_0 e^{-kt} = [A]_0 e^{-0.0693 x 20} = [A]_0 e^{-1.386} = [A]_0 x 0.250.

Summary

CBSE Chemistry covers three branches. Physical chemistry includes atomic structure, bonding, thermodynamics (delta G = delta H - T delta S), equilibrium (Le Chatelier’s principle, Kc, Kp), electrochemistry (Nernst equation), kinetics (first-order rate laws, Arrhenius equation), and solutions (colligative properties). Inorganic chemistry covers periodic trends, s-block and p-block properties, transition metal chemistry (variable oxidation states, coordination compounds, crystal field theory). Organic chemistry covers IUPAC nomenclature, hydrocarbons, halogen derivatives (SN1, SN2, E1, E2), oxygen compounds, nitrogen compounds, polymers, and biomolecules. Key exam skills include balancing equations, arrow-pushing mechanisms, and applying named reactions.

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