Notes for Class 11 chemistry

Unit 1: Some Basic Concepts of Chemistry 

This unit introduces the fundamentals of chemistry, emphasizing quantitative aspects and the nature of matter.

Key Subtopics and Detailed Explanation:

• Importance and Scope of Chemistry: Chemistry is the study of matter and its transformations. It overlaps with physics (e.g., quantum mechanics in atomic structure) and biology (e.g., biochemistry). Applications include pharmaceuticals, materials science, and environmental science.
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• Nature of Matter: Matter is anything with mass and volume. Classified as:
  • Mixtures: Homogeneous (e.g., saltwater) or heterogeneous (e.g., sand in water).
  • Pure Substances: Elements (e.g., oxygen) or compounds (e.g., water).
  • States: Solid, liquid, gas (plasma and Bose-Einstein condensate are advanced).
• Laws of Chemical Combination:
  • Law of Conservation of Mass (Lavoisier): Mass is neither created nor destroyed in a chemical reaction. Example: In combustion, mass of reactants = mass of products.
  • Law of Definite Proportions (Proust): A compound always contains elements in fixed ratios by mass. E.g., Water is always 11.11% H and 88.89% O.
  • Law of Multiple Proportions (Dalton): Elements combine in simple whole-number ratios to form multiple compounds. E.g., CO (12:16) and CO₂ (12:32).
  • Gay-Lussac's Law of Gaseous Volumes: Gases react in simple volume ratios.
  • Avogadro's Law: Equal volumes of gases contain equal molecules under same conditions.
• Dalton's Atomic Theory: Atoms are indivisible, indestructible particles. Postulates: All matter is made of atoms; atoms of same element are identical; compounds form by whole-number ratios. Limitations: Doesn't explain isotopes or subatomic particles.
• Atomic and Molecular Masses: Atomic mass unit (amu) = 1/12th mass of C-12. Molecular mass = sum of atomic masses.
• Mole Concept and Molar Mass: Mole = 6.022 × 10²³ particles (Avogadro's number, N_A). Molar mass = mass of 1 mole (in g/mol). Example: Molar mass of H₂O = 18 g/mol.
• Percentage Composition: % of element = (atomic mass × no. of atoms / molecular mass) × 100. Example: In CO₂, %C = (12/44) × 100 ≈ 27.27%.
• Empirical and Molecular Formula: Empirical = simplest ratio (e.g., CH for benzene). Molecular = n × empirical, where n = molecular mass / empirical mass.
• Chemical Reactions and Stoichiometry: Balancing equations. Calculations: Mass-mass, mass-volume, etc. Example: For 2H₂ + O₂ → 2H₂O, 4g H₂ reacts with 32g O₂ to give 36g H₂O.

Important Formulas:

• Number of moles (n) = mass / molar mass = volume (in L) / 22.4 (at STP for gases).
• Limiting reagent: Reactant that is completely consumed.

Practice Tips: Solve numerical problems on mole concept and stoichiometry.

Unit 2: Structure of Atom 

This unit covers atomic models and quantum mechanics basics.

Key Subtopics and Detailed Explanation:

• Discovery of Subatomic Particles: Electron (Thomson, 1897, cathode rays), Proton (Goldstein, anode rays), Neutron (Chadwick, 1932).
• Atomic Number, Isotopes, Isobars: Atomic number (Z) = protons. Mass number (A) = protons + neutrons. Isotopes: Same Z, different A (e.g., ¹²C, ¹⁴C). Isobars: Same A, different Z (e.g., ⁴⁰Ar, ⁴⁰Ca).
• Atomic Models:
  • Thomson's Plum Pudding: Atom as positive sphere with embedded electrons. Limitations: Doesn't explain scattering.
  • Rutherford's Model: Nucleus positive, electrons orbit. Alpha-scattering experiment: Most particles pass through, some deflect. Limitations: Doesn't explain stability (electrons should spiral in).
  • Bohr's Model: Electrons in fixed orbits with quantized energy. E = -13.6 / n² eV for H-atom. Limitations: Works only for H-like atoms, ignores wave nature.
• Dual Nature of Matter and Light: Light as wave (diffraction) and particle (photoelectric effect). de Broglie: λ = h / p (wavelength for matter waves).
• Heisenberg Uncertainty Principle: ∆x × ∆p ≥ h/4π (can't know position and momentum exactly).
• Quantum Numbers and Orbitals:
  • Principal (n): Energy level (1,2,3...).
  • Azimuthal (l): Subshell (0=s, 1=p, 2=d, 3=f).
  • Magnetic (m_l): Orientation (-l to +l).
  • Spin (m_s): +1/2 or -1/2.
  • Shapes: s (spherical), p (dumbbell), d (cloverleaf).
• Electron Configuration Rules: Aufbau (fill lower energy first), Pauli (no two electrons same quantum numbers), Hund (max unpaired in subshell).
• Stability: Half-filled (e.g., p³) and fully filled (p⁶) are stable, explaining exceptions like Cr (4s¹3d⁵).

Important Formulas:

• Bohr radius: a₀ = 0.529 Å.
• Energy: ∆E = hν.

Practice Tips: Draw orbital shapes; write configurations for elements 1-30.

Unit 3: Classification of Elements and Periodicity in Properties 

Focuses on the periodic table and trends.

Key Subtopics and Detailed Explanation:

• History of Periodic Table: Dobereiner's triads, Newlands' octaves, Mendeleev's (by atomic mass), Modern (by atomic number, Moseley).
• Modern Periodic Law: Properties are periodic functions of atomic number. 7 periods, 18 groups.
• Periodic Trends:
  • Atomic Radius: Decreases across period (increasing Z_eff), increases down group (more shells).
  • Ionic Radius: Cations smaller, anions larger than neutral.
  • Ionization Enthalpy (IE): Energy to remove electron. Increases across period, decreases down group. Exceptions: Group 2>13, 15>16 due to stable configurations.
  • Electron Gain Enthalpy (∆_egH): Energy change on adding electron. More negative = easier. Halogens most negative.
  • Electronegativity: Pauling scale; F highest (4.0). Increases across, decreases down.
  • Valency: Based on group (1-8 for s/p-block).
• Nomenclature for Z > 100: E.g., Ununbium (Uub) for 112 (now Copernicium).

Property	Across Period	Down Group
Atomic Radius	Decreases	Increases
IE	Increases	Decreases
Electronegativity	Increases	Decreases

Practice Tips: Explain trends with examples like why Na has lower IE than Mg.

Unit 4: Chemical Bonding and Molecular Structure 

Explains how atoms bond.

Key Subtopics and Detailed Explanation:

• Valence Electrons: Outermost electrons involved in bonding.
• Ionic Bond: Electron transfer (e.g., NaCl). Favored between metals/non-metals. Bond parameters: Lattice energy.
• Covalent Bond: Sharing (e.g., H₂). Polar (HCl) vs non-polar (Cl₂). Fajans' rule: Small cation, large anion → covalent character.
• Lewis Structures: Dot diagrams. Octet rule (exceptions: BF₃ incomplete, SF₆ expanded).
• Valence Bond Theory (VBT): Orbital overlap. Sigma (head-on), pi (sideways).
• Resonance: Delocalized electrons (e.g., benzene).
• VSEPR Theory: Electron pairs repel to minimize energy. Shapes: Linear (BeCl₂), Trigonal planar (BF₃), Tetrahedral (CH₄), etc. Include lone pairs (e.g., NH₃ pyramidal).
• Hybridization: Mixing orbitals. sp (linear), sp² (trigonal), sp³ (tetrahedral), sp³d (trigonal bipyramidal), sp³d² (octahedral).
• Molecular Orbital Theory (MOT): Bonding (lower energy) and antibonding orbitals. Bond order = (bonding e⁻ - antibonding e⁻)/2. For O₂, bond order 2, paramagnetic.
• Hydrogen Bond: Strong dipole-dipole (e.g., HF, H₂O). Explains high boiling point of water.

Important Formulas:

• Bond order in MOT.
• Formal charge = valence e⁻ - non-bonding e⁻ - 1/2 bonding e⁻.

Practice Tips: Draw Lewis structures and predict shapes for molecules like PCl₅, SF₆.

Unit 5: Chemical Thermodynamics 

Deals with energy changes in reactions.

Key Subtopics and Detailed Explanation:

• Systems and Surroundings: Open (exchange matter/energy), closed (energy only), isolated (none). State functions: Independent of path (e.g., U, H).
• First Law of Thermodynamics: ∆U = q + w. U = internal energy, q = heat, w = work (-P∆V for expansion).
• Enthalpy (H): H = U + PV. ∆H = ∆U + ∆(PV). For constant P, ∆H = q_p.
• Heat Capacity: C = q/∆T. Specific heat (per gram), molar (per mole).
• Hess's Law: ∆H total = sum of ∆H steps. Used for indirect calculations.
• Enthalpies: Bond dissociation (break bonds), formation (from elements), combustion (with O₂), etc.
• Second Law: Entropy (S) measures disorder. ∆S > 0 for spontaneous. Gibbs energy: G = H - TS. ∆G = ∆H - T∆S. ∆G < 0 spontaneous.
• Third Law: S = 0 at 0 K for perfect crystal.

Important Formulas:

• ∆G = -RT ln K (equilibrium).
• Hess's: E.g., C + O₂ → CO₂; use intermediate steps.

Examples: Exothermic (∆H < 0, heat release), endothermic (∆H > 0).

Practice Tips: Calculate ∆H using bond energies: ∆H = Σ bond broken - Σ bond formed.

Unit 6: Equilibrium 

Covers reversible reactions.

Key Subtopics and Detailed Explanation:

• Dynamic Equilibrium: Forward = reverse rate. Physical (e.g., vapor-liquid), chemical (e.g., N₂ + 3H₂ ⇌ 2NH₃).
• Law of Mass Action: K_c = [products]/[reactants] (powers = coefficients).
• Equilibrium Constant: K_c (concentration), K_p (pressure). Relation: K_p = K_c (RT)^∆n.
• Factors Affecting Equilibrium (Le Chatelier's Principle): Concentration (add reactant → shift right), temperature (endothermic → heat favors forward), pressure (for gases, high P favors fewer moles).
• Ionic Equilibrium: Weak acids/bases ionize partially. K_a = [H⁺][A⁻]/[HA].
• pH: pH = -log[H⁺]. Strong acid: pH low; weak: higher.
• Hydrolysis: Salts of weak acid/strong base → basic (e.g., Na₂CO₃).
• Buffer Solution: Resist pH change. Acidic: Weak acid + salt (HA + NaA). Henderson: pH = pK_a + log[salt]/[acid].
• Solubility Product (K_sp): For sparingly soluble (e.g., AgCl ⇌ Ag⁺ + Cl⁻, K_sp = [Ag⁺][Cl⁻]).
• Common Ion Effect: Adding common ion suppresses ionization.

Important Formulas:

• For buffer: pOH = pK_b + log[salt]/[base] for basic.
• Q vs K: Q < K → forward.

Practice Tips: Calculate K from concentrations; predict shifts.

Unit 7: Redox Reactions 

About oxidation-reduction.

Key Subtopics and Detailed Explanation:

• Oxidation and Reduction: Oxidation = loss e⁻/gain O/increase oxidation number (ON). Reduction = opposite.
• Oxidation Number: Rules: Elements 0, O -2 (except peroxides -1), H +1, etc. Example: In KMnO₄, Mn +7.
• Balancing Redox: Half-reaction method: Balance atoms, e⁻, add H₂O/H⁺, combine. Ion-electron for acidic/basic.
• Applications: Electrochemistry (batteries), corrosion, metabolism.

Examples: Zn + Cu²⁺ → Zn²⁺ + Cu (Zn oxidized, Cu reduced).

Practice Tips: Assign ON and balance equations like Cr₂O₇²⁻ + Fe²⁺ in acid.

Unit 8: Organic Chemistry – Some Basic Principles and Techniques 

Introduction to organic compounds.

Key Subtopics and Detailed Explanation:

• General Introduction: Organic = carbon compounds (except oxides, carbonates). Tetravalency of C, catenation.
• Purification Methods: Crystallization, sublimation, distillation, chromatography.
• Qualitative Analysis: Detection of C, H, N, S, halogens (Lassaigne's test).
• Quantitative Analysis: C/H by combustion, N by Dumas/Kjeldahl.
• Classification: Aliphatic/aromatic, saturated/unsaturated, functional groups (e.g., -OH alcohol).
• IUPAC Nomenclature: Root + suffix + prefix. E.g., CH₃CH₂OH = ethanol.
• Electronic Effects: Inductive (±I), resonance (±R), hyperconjugation.
• Bond Fission: Homolytic (free radicals •), heterolytic (carbocation⁺, carbanion⁻).
• Reaction Intermediates: Electrophiles (e⁺ seekers), nucleophiles (e⁻ donors).
• Types of Reactions: Addition, substitution, elimination.

Practice Tips: Name compounds like 2-methylpropane; identify effects in phenol.

Unit 9: Hydrocarbons 

Compounds of C and H.

Key Subtopics and Detailed Explanation:

• Alkanes: General formula C_nH_{2n+2}. Nomenclature (straight/branched). Isomerism (chain, position). Conformation: Ethane staggered/ eclipsed (staggered stable).
  • Physical: Non-polar, boiling increases with mass.
  • Chemical: Halogenation (free radical: initiation, propagation, termination), combustion, pyrolysis.
• Alkenes: C_nH_{2n}. Double bond. Geometrical isomerism (cis-trans in >C=C<).
  • Preparation: Dehydrohalogenation, dehydration.
  • Reactions: Addition (H₂, HX - Markovnikov: H to less H carbon; peroxide effect anti-Markovnikov for HBr), ozonolysis (cleave double bond), oxidation.
• Alkynes: C_nH_{2n-2}. Triple bond. Acidic H in terminal.
  • Preparation: From vicinal dihalides.
  • Reactions: Addition (2 molecules), acidic nature (react with Na).
• Aromatic Hydrocarbons: Benzene C₆H₆. Resonance (Kekule structures), aromaticity (Huckel's 4n+2 e⁻).
  • Reactions: Electrophilic substitution (nitration: HNO₃+H₂SO₄ → NO₂⁺), sulphonation, halogenation, Friedel-Crafts (alkylation/acylation).
  • Directive Effects: Ortho-para (activators like -OH), meta (deactivators like -NO₂).
  • Toxicity: Benzene carcinogenic.

Practice Tips: Mechanisms for addition/substitution; draw resonance for benzene.