Chemistry Notes – Set 4: Detailed Guide for UPSC, PCS, SSC Competitive Exams

Class 9: Structure of the Atom

Detailed Concepts:

• Atomic Models:
  • Thomson’s Model: Atom as a sphere of positive charge with embedded electrons (plum pudding model), disproved due to lack of nucleus.
  • Rutherford’s Model: Based on alpha particle scattering experiment:
    • Atom has a dense, positively charged nucleus.
    • Electrons orbit the nucleus, most of atom is empty space.
    • Limitations: Couldn’t explain electron stability or spectra.
  • Bohr’s Model: Electrons in fixed orbits with quantized energy levels, applicable to hydrogen-like atoms.
    • Electrons absorb/emit energy to jump between orbits.
    • Energy of electron: E_n = –2.18 × 10⁻¹⁸ J (Z²/n²), where Z = atomic number, n = principal quantum number.
• Subatomic Particles:
  • Proton: Positive charge (+1), mass ≈ 1 u, in nucleus.
  • Neutron: Neutral, mass ≈ 1 u, in nucleus.
  • Electron: Negative charge (–1), negligible mass (1/1836 u), in orbitals.
• Atomic Number and Mass Number:
  • Atomic Number (Z): Number of protons, defines element (e.g., C: Z = 6).
  • Mass Number (A): Protons + neutrons (e.g., C-12: A = 12).
  • Isotopes: Same Z, different A (e.g., C-12, C-14; 6 protons, 6 vs. 8 neutrons).
  • Isobars: Same A, different Z (e.g., C-14, N-14).
• Electron Arrangement:
  • Electrons occupy shells (K, L, M) with maximum capacities: 2n² (n = shell number).
  • Example: Na (Z = 11): 2 (K), 8 (L), 1 (M).
  • Valency: Number of electrons an atom gains, loses, or shares to achieve octet (e.g., Na: valency 1, loses 1 electron; Cl: valency 1, gains 1 electron).
• Atomic Spectra:
  • Emission Spectra: Light emitted when electrons drop to lower energy levels (e.g., hydrogen’s line spectrum: Balmer series in visible region).
  • Absorption Spectra: Dark lines where electrons absorb specific wavelengths.
• Applications in Exams: Atomic models, isotopes, and valency are key for objective and conceptual questions.

Formulas:

• Bohr’s Energy: E_n = –2.18 × 10⁻¹⁸ J (Z²/n²) for hydrogen-like atoms.
• Wavelength (Rydberg Formula): 1/λ = R_H (1/n₁² – 1/n₂²), R_H = 1.097 × 10⁷ m⁻¹.
• Energy of Photon: E = hν = hc/λ, h = 6.626 × 10⁻³⁴ J·s.
• Electron Capacity: 2n² electrons per shell (n = 1, 2, 3…).

Applications:

• Competitive Exams:
  • UPSC/PCS: Questions on isotopes in nuclear chemistry (e.g., C-14 dating) or Bohr’s model in spectroscopy.
  • SSC: Objective questions on electron arrangement, valency, or atomic models.
  • Descriptive: Explain Rutherford’s scattering experiment or applications of isotopes.
• Real-World:
  • Nuclear Science: Isotopes in radiocarbon dating, medical imaging (e.g., Tc-99m).
  • Technology: Spectroscopy in material analysis, lasers (electron transitions).
  • Environment: Isotopes in tracing pollution sources.
• Exam Tips:
  • Focus on Bohr’s model and isotope applications for objective questions.
  • Link spectra to environmental or analytical chemistry for mains.

Diagram (Textual Description):

• Rutherford’s Scattering Experiment: A thin gold foil bombarded by alpha particles. Most pass through, some deflect at large angles, few bounce back. Label nucleus (small, dense, positive) and electron cloud (sparse), showing alpha particle paths.

---

Class 10: Carbon and Its Compounds

Detailed Concepts:

• Carbon’s Uniqueness:
  • Tetravalent (4 valence electrons), forms covalent bonds.
  • Catenation: Ability to form long chains or rings (e.g., hydrocarbons).
  • Isomerism: Different structures for same molecular formula (e.g., butane, isobutane: C₄H₁₀).
• Allotropes:
  • Diamond: Tetrahedral covalent network, hard, insulator.
  • Graphite: Layered structure, conducts electricity due to delocalized electrons.
  • Fullerenes: Cage-like (e.g., C₆₀), used in nanotechnology.
• Hydrocarbons:
  • Saturated: Single bonds, alkanes (CₙH₂ₙ₊₂, e.g., CH₄).
  • Unsaturated: Double/triple bonds, alkenes (CₙH₂ₙ, e.g., C₂H₄), alkynes (CₙH₂ₙ₋₂, e.g., C₂H₂).
• Functional Groups:
  • Alcohol (-OH, e.g., CH₃OH).
  • Aldehyde (-CHO, e.g., CH₃CHO).
  • Ketone (>C=O, e.g., CH₃COCH₃).
  • Carboxylic acid (-COOH, e.g., CH₃COOH).
  • Halogen (-X, e.g., CH₃Cl).
• Nomenclature (IUPAC):
  • Longest carbon chain as base (e.g., methane, ethane, propane).
  • Functional group gets lowest number (e.g., CH₃CH₂OH: ethanol).
• Chemical Properties:
  • Combustion: Hydrocarbons burn to form CO₂ + H₂O (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O).
  • Oxidation: Alcohols to aldehydes/acids (e.g., CH₃CH₂OH → CH₃CHO with Cu, 573 K).
  • Addition: Unsaturated hydrocarbons add H₂, halogens (e.g., C₂H₄ + H₂ → C₂H₆, Ni catalyst).
  • Substitution: Alkanes react with halogens in sunlight (e.g., CH₄ + Cl₂ → CH₃Cl + HCl).
• Soaps and Detergents:
  • Soap: Sodium/potassium salts of fatty acids (e.g., sodium stearate, C₁₇H₃₅COONa).
  • Micelle Formation: Hydrophilic head (carboxylate) interacts with water, hydrophobic tail traps oil, enabling cleaning.
  • Detergents: Synthetic, work in hard water (e.g., sodium dodecyl sulfate).
• Applications in Exams: Nomenclature, reactions, and environmental impacts are key for objective and descriptive questions.

Formulas:

• Alkanes: CₙH₂ₙ₊₂ (e.g., CH₄, C₂H₆).
• Alkenes: CₙH₂ₙ (e.g., C₂H₄).
• Alkynes: CₙH₂ₙ₋₂ (e.g., C₂H₂).
• Combustion: CₙH₂ₙ₊₂ + (3n+1)/2 O₂ → nCO₂ + (n+1)H₂O.
• Ethanol Oxidation: CH₃CH₂OH + [O] → CH₃CHO + H₂O.

Applications:

• Competitive Exams:
  • UPSC/PCS: Questions on hydrocarbons in fuels (e.g., methane in biogas) or soaps in environmental contexts.
  • SSC: Objective questions on nomenclature, functional groups, or allotropes.
  • Descriptive: Explain micelle formation or environmental impact of hydrocarbons.
• Real-World:
  • Energy: Methane in natural gas, ethanol as biofuel.
  • Industry: Graphite in electrodes, detergents in cleaning.
  • Environment: CO₂ from combustion in greenhouse effect.
• Exam Tips:
  • Master nomenclature and reaction types for objective questions.
  • Link hydrocarbons to environmental science (e.g., global warming) for mains.

Diagram (Textual Description):

• Micelle Formation: A spherical micelle with hydrophilic heads (carboxylate groups) facing outward toward water and hydrophobic tails (hydrocarbon chains) inward, trapping oil. Label hydrophilic and hydrophobic parts, showing oil droplet inside.

---

Class 11: Thermodynamics

Detailed Concepts:

• Thermodynamics: Study of energy transformations in physical and chemical processes.
• System and Surroundings:
  • System: Part under study (open: exchanges matter/energy; closed: exchanges energy; isolated: exchanges neither).
  • Surroundings: Everything outside the system.
• Thermodynamic Properties:
  • State Functions: Depend only on current state (e.g., internal energy U, enthalpy H, entropy S, Gibbs free energy G).
  • Path Functions: Depend on process path (e.g., heat q, work w).
• Laws of Thermodynamics:
  • Zeroth Law: If two systems are in thermal equilibrium with a third, they are in equilibrium with each other (defines temperature).
  • First Law: Energy conservation, ΔU = q + w, where w = –PΔV (work done by volume change).
  • Second Law: Entropy of an isolated system increases for spontaneous processes (ΔS_total > 0).
  • Third Law: Entropy of a perfect crystal at 0 K is zero.
• Enthalpy (H): H = U + PV, heat content at constant pressure (ΔH = q_p).
  • Exothermic: ΔH < 0, heat released (e.g., combustion).
  • Endothermic: ΔH > 0, heat absorbed (e.g., evaporation).
• Gibbs Free Energy (G): G = H – TS, determines spontaneity:
  • ΔG < 0: Spontaneous.
  • ΔG = 0: Equilibrium.
  • ΔG > 0: Non-spontaneous.
  • ΔG = ΔH – TΔS (at constant T).
• Thermochemical Equations:
  • Show enthalpy changes (e.g., C + O₂ → CO₂, ΔH = –393.5 kJ/mol).
  • Types: Formation (ΔH_f), combustion (ΔH_c), neutralization.
• Hess’s Law: ΔH is independent of path, sum of step enthalpies equals total (e.g., ΔH for C → CO₂ via CO).
• Bond Enthalpy: Energy to break 1 mole of bonds (e.g., C–H: 413 kJ/mol). ΔH = Σ(Bond energies of reactants) – Σ(Bond energies of products).
• Applications in Exams: Enthalpy, Gibbs free energy, and Hess’s law are key for numerical and conceptual questions.

Formulas:

• First Law: ΔU = q + w.
• Work: w = –PΔV (constant pressure).
• Enthalpy: ΔH = ΔU + PΔV or ΔH = ΔU + Δn_gRT (for gases).
• Gibbs Free Energy: ΔG = ΔH – TΔS.
• Hess’s Law: ΔH_total = ΣΔH_steps.
• Entropy: ΔS = q_rev / T.
• Bond Enthalpy: ΔH = Σ(BE_reactants) – Σ(BE_products).

Applications:

• Competitive Exams:
  • UPSC/PCS: Questions on thermodynamics in energy production (e.g., combustion) or environmental contexts (e.g., heat engines).
  • SSC: Objective questions on ΔH calculations, spontaneity, or Hess’s law.
  • Descriptive: Explain spontaneity using ΔG or Hess’s law in industrial processes.
• Real-World:
  • Energy: Combustion in power plants, efficiency in heat engines.
  • Industry: Enthalpy in chemical manufacturing (e.g., ammonia synthesis).
  • Environment: Entropy in climate change models.
• Exam Tips:
  • Master ΔG and Hess’s law for numerical questions.
  • Link thermodynamics to environmental science (e.g., fuel efficiency) for mains.

Diagram (Textual Description):

• Exothermic Reaction Profile: Graph of energy (y-axis) vs. reaction coordinate (x-axis). Reactants at higher energy, transition state at peak, products at lower energy. Label ΔH (negative), activation energy (E_a), and transition state.

---

Class 12: Electrochemistry

Detailed Concepts:

• Electrochemistry: Study of chemical processes involving electron transfer, linking chemical and electrical energy.
• Electrochemical Cells:
  • Galvanic (Voltaic) Cell: Spontaneous redox reaction generates electricity (e.g., Zn | Zn²⁺ || Cu²⁺ | Cu, E°_cell = 1.10 V).
    • Anode: Oxidation (e.g., Zn → Zn²⁺ + 2e⁻).
    • Cathode: Reduction (e.g., Cu²⁺ + 2e⁻ → Cu).
    • Salt bridge (e.g., KCl) maintains charge neutrality.
  • Electrolytic Cell: Non-spontaneous reaction driven by external voltage (e.g., NaCl electrolysis).
• Electrode Potential:
  • Standard Electrode Potential (E°): Measured against standard hydrogen electrode (SHE, E° = 0 V) at 25°C, 1 M, 1 atm.
  • Example: E°_Cu²⁺/Cu = 0.34 V, E°_Zn²⁺/Zn = –0.76 V.
• Cell Potential: E_cell = E_cathode – E_anode. For spontaneity, E_cell > 0.
• Nernst Equation: Adjusts potential for non-standard conditions: E = E° – (0.059/n) log Q, where Q = reaction quotient, n = electrons transferred.
• Gibbs Free Energy: ΔG = –nFE_cell, negative for spontaneous reactions.
• Conductance:
  • Specific Conductance (κ): Conductivity of 1 m³ solution (S m⁻¹).
  • Molar Conductance (Λ_m): Λ_m = κ × 1000 / M (S cm² mol⁻¹).
  • Cell Constant: G* = l/A (m⁻¹), where l = electrode distance, A = area.
• Faraday’s Laws:
  • First Law: Mass deposited (m) = Z × Q, where Z = electrochemical equivalent, Q = charge.
  • Second Law: m ∝ equivalent weight (Z = M/nF), F = 96485 C/mol.
• Applications:
  • Batteries: Lead-acid (cars), Li-ion (electronics).
  • Electroplating: Coating metals (e.g., Cu, Ag) for corrosion resistance.
  • Corrosion: Electrochemical process (e.g., rusting), prevented by sacrificial anodes (e.g., Zn).

Formulas:

• Cell Potential: E_cell = E°_cathode – E°_anode.
• Nernst Equation: E = E° – (0.059/n) log Q (at 25°C).
• Gibbs Free Energy: ΔG = –nFE_cell.
• Molar Conductance: Λ_m = κ × 1000 / M.
• Faraday’s First Law: m = (Q × M) / (n × F).
• Specific Conductance: κ = G × (l/A).

Applications:

• Competitive Exams:
  • UPSC/PCS: Questions on batteries in renewable energy or corrosion in infrastructure.
  • SSC: Objective questions on Nernst equation, cell potential, or Faraday’s laws.
  • Descriptive: Explain electrochemical cells in fuel cells or electroplating in industry.
• Real-World:
  • Energy: Fuel cells (H₂-O₂) for clean energy, batteries in EVs.
  • Industry: Electroplating in jewelry, Cu purification.
  • Environment: Corrosion prevention in pipelines, ships.
• Exam Tips:
  • Master Nernst equation and Faraday’s laws for numerical questions.
  • Link electrochemistry to environmental science (e.g., battery recycling) for mains.

Diagram (Textual Description):

• Galvanic Cell: Zn anode in ZnSO₄ solution, Cu cathode in CuSO₄ solution, connected by a salt bridge (KCl). Electrons flow from Zn to Cu via wire, ions migrate through salt bridge. Label anode (oxidation), cathode (reduction), and E_cell.