Class 9: Atoms and Molecules
Detailed Concepts:
- Atoms: Smallest unit of an element that retains its chemical properties, consisting of protons (positive, in nucleus), neutrons (neutral, in nucleus), and electrons (negative, in orbitals).
- Atomic Number (Z): Number of protons, defines the element (e.g., C: Z = 6).
- Mass Number (A): Protons + neutrons (e.g., C-12: 6 protons + 6 neutrons).
- Isotopes: Atoms with same Z but different A (e.g., C-12, C-14).
- Molecules: Two or more atoms chemically bonded, either homoatomic (e.g., O₂) or heteroatomic (e.g., H₂O). Molecules represent compounds or elemental forms.
- Laws of Chemical Combination:
- Law of Conservation of Mass: Total mass of reactants equals total mass of products in a closed system (e.g., 2H₂ + O₂ → 2H₂O, mass conserved).
- Law of Definite Proportions: A compound has a fixed composition by mass (e.g., H₂O: 11.11% H, 88.89% O).
- Law of Multiple Proportions: When two elements form multiple compounds, their mass ratios are simple whole numbers (e.g., CO vs. CO₂, C:O ratios 1:1 vs. 1:2).
- Atomic Mass: Mass of an atom relative to 1/12th of C-12, measured in atomic mass units (u). Example: H = 1 u, O = 16 u, C = 12 u.
- Molecular Mass: Sum of atomic masses of atoms in a molecule (e.g., CO₂: 12 + 2×16 = 44 u).
- Mole Concept:
- 1 mole = 6.022 × 10²³ particles (Avogadro’s number).
- Molar mass = molecular mass in grams (e.g., CO₂: 44 g/mol).
- Moles = Mass / Molar mass or Number of particles / 6.022 × 10²³.
- Percentage Composition: % of an element = (Number of atoms × Atomic mass / Molecular mass) × 100. Example: % C in CO₂ = (12/44) × 100 = 27.27%.
- Chemical Formulae:
- Empirical Formula: Simplest whole-number ratio of atoms (e.g., CH₂O for glucose).
- Molecular Formula: Actual number of atoms (e.g., C₆H₁₂O₆ for glucose, n = 6 × CH₂O).
- Steps: Calculate moles, divide by smallest, multiply for whole numbers.
- Stoichiometry: Quantitative relationships in reactions (e.g., 2H₂ + O₂ → 2H₂O; 2 moles H₂ react with 1 mole O₂).
Formulas:
- Moles:
- Moles = Mass / Molar mass.
- Moles = Number of particles / 6.022 × 10²³.
- Percentage Composition: % = (n × Atomic mass / Molecular mass) × 100.
- Empirical Formula: Moles of each element / Smallest mole value.
- Molecular Formula: n × Empirical formula, where n = Molecular mass / Empirical mass.
- Gas Volume at STP: 1 mole of gas = 22.4 L at standard temperature (0°C) and pressure (1 atm).
Applications:
- Competitive Exams:
- UPSC/PCS: Questions on mole calculations in environmental contexts (e.g., CO₂ emissions) or stoichiometry in industrial reactions (e.g., ammonia synthesis).
- SSC: Objective questions on percentage composition, empirical formulas, or isotope applications.
- Descriptive: Explain the law of multiple proportions with examples like CO and CO₂ or role of isotopes in carbon dating.
- Real-World:
- Industry: Mole calculations in fertilizer production (e.g., urea, (NH₂)₂CO).
- Environment: CO₂ mass calculations in greenhouse gas studies.
- Science: C-14 dating in archaeology, isotopic analysis in nuclear technology.
- Exam Tips:
- Master mole concept and stoichiometry for numerical questions.
- Link isotopes to nuclear chemistry or environmental science for mains.
Diagram (Textual Description):
- Atomic Structure: A central nucleus (protons + neutrons) surrounded by electron shells. For oxygen (O-16): 8 protons, 8 neutrons, 8 electrons (2 in 1s, 2 in 2s, 4 in 2p). Label shells and particles, showing isotopic variation (e.g., O-18: 10 neutrons).
Class 10: Metals and Non-Metals
Detailed Concepts:
- Metals vs. Non-Metals:
- Metals: Shiny, malleable, ductile, good conductors of heat/electricity (e.g., Fe, Cu, Al). Form positive ions by losing electrons (e.g., Na → Na⁺).
- Non-Metals: Dull, brittle, poor conductors (except graphite), form negative ions or covalent compounds (e.g., O, Cl). Examples: C (graphite, diamond), S, N.
- Physical Properties:
- Metals: High melting/boiling points (except alkali metals), sonorous, high density (except Na, K).
- Non-Metals: Low melting/boiling points (e.g., S melts at 113°C), non-sonorous, low density (except diamond).
- Chemical Properties of Metals:
- Reaction with Oxygen: Form basic or amphoteric oxides (e.g., 2Mg + O₂ → 2MgO, basic; Zn + O₂ → ZnO, amphoteric).
- Reaction with Water: Active metals (e.g., Na, K) react to form H₂ (e.g., 2Na + 2H₂O → 2NaOH + H₂); less reactive metals (e.g., Fe) react with steam.
- Reaction with Acids: Produce H₂ gas (e.g., Zn + 2HCl → ZnCl₂ + H₂). Less reactive metals (e.g., Cu) don’t react with dilute acids.
- Reaction with Non-Metals: Form ionic compounds (e.g., Na + Cl → NaCl).
- Chemical Properties of Non-Metals:
- Form acidic or neutral oxides (e.g., CO₂, acidic; CO, neutral).
- Form covalent compounds with other non-metals (e.g., CCl₄).
- Reactivity Series: Order of metal reactivity (K > Na > Ca > Mg > Al > Zn > Fe > Pb > H > Cu > Ag > Au). Determines displacement reactions (e.g., Zn + CuSO₄ → ZnSO₄ + Cu).
- Metallurgy:
- Extraction:
- Highly reactive metals (e.g., Na, Al): Electrolysis (e.g., Al from bauxite via Hall-Héroult process).
- Moderately reactive metals (e.g., Fe, Zn): Reduction (e.g., Fe₂O₃ + 3CO → 2Fe + 3CO₂).
- Low reactive metals (e.g., Au, Ag): Occur native or extracted by roasting.
- Refining: Electrolysis (e.g., Cu purification), distillation, or liquation.
- Extraction:
- Corrosion: Degradation of metals by environmental reactions (e.g., rusting: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O). Prevented by galvanization (Zn coating), painting, or sacrificial anodes.
- Alloys: Mixtures of metals or metals with non-metals (e.g., steel: Fe + C; brass: Cu + Zn). Enhance strength, corrosion resistance.
Formulas:
- Metal Oxides: 2M + O₂ → 2MO (e.g., 2Mg + O₂ → 2MgO).
- Displacement: M + NX → MX + N, where M is more reactive than N.
- Rusting: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O.
- Electrolysis of NaCl: 2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂.
Applications:
- Competitive Exams:
- UPSC/PCS: Questions on metallurgy in industrial contexts (e.g., steel production) or corrosion in infrastructure.
- SSC: Objective questions on reactivity series, alloy properties, or non-metal oxides.
- Descriptive: Discuss extraction of Al or Fe, or corrosion prevention methods.
- Real-World:
- Industry: Steel in construction, Al in aircraft, Cu in wiring.
- Environment: Acidic oxides (SO₂, CO₂) in acid rain, corrosion in pipelines.
- Technology: Alloys in machinery, non-metals (Si) in semiconductors.
- Exam Tips:
- Memorize reactivity series for displacement reactions.
- Link corrosion and acid rain to environmental science for mains.
Diagram (Textual Description):
- Reactivity Series: A vertical list: K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au. Arrows show decreasing reactivity, with displacement reactions (e.g., Zn + CuSO₄ → ZnSO₄ + Cu) and extraction methods (electrolysis for Na, Al; reduction for Fe, Zn).
Class 11: States of Matter
Detailed Concepts:
- States of Matter: Gas, liquid, solid, governed by intermolecular forces and kinetic energy.
- Gas Laws:
- Boyle’s Law: At constant temperature, P ∝ 1/V (PV = constant). Example: Doubling pressure halves volume.
- Charles’ Law: At constant pressure, V ∝ T (V/T = constant). Example: Heating a gas increases volume.
- Gay-Lussac’s Law: At constant volume, P ∝ T (P/T = constant).
- Avogadro’s Law: At constant T and P, V ∝ n (V/n = constant). 1 mole of gas = 22.4 L at STP (0°C, 1 atm).
- Combined Gas Law: PV/T = constant.
- Ideal Gas Equation: PV = nRT, where R = 0.0821 L atm mol⁻¹ K⁻¹ or 8.314 J mol⁻¹ K⁻¹.
- Real Gases: Deviate from ideal behavior at high pressure/low temperature due to intermolecular forces and molecular volume.
- Van der Waals Equation: [P + a(n/V)²](V – nb) = nRT, where a = intermolecular attraction, b = volume of molecules.
- Kinetic Theory of Gases:
- Gases consist of particles in random motion, with no intermolecular forces (ideal gas).
- Average kinetic energy ∝ temperature (KE = 3/2 kT per molecule, k = Boltzmann constant).
- Liquefaction of Gases:
- Achieved by high pressure and low temperature (e.g., NH₃, CO₂).
- Critical Temperature: Temperature above which a gas cannot be liquefied (e.g., CO₂: 31.1°C).
- Critical Pressure: Pressure needed to liquefy at critical temperature.
- Liquid State:
- Vapour Pressure: Pressure exerted by vapor in equilibrium with liquid, increases with temperature.
- Surface Tension: Force per unit length due to cohesive forces (e.g., water droplets form spheres).
- Viscosity: Resistance to flow, decreases with temperature (e.g., honey more viscous than water).
- Applications in Exams: Gas laws, real gas deviations, and liquid properties are key for numerical and conceptual questions.
Formulas:
- Boyle’s Law: P₁V₁ = P₂V₂ (constant T).
- Charles’ Law: V₁/T₁ = V₂/T₂ (constant P).
- Ideal Gas Equation: PV = nRT.
- Van der Waals Equation: [P + a(n/V)²](V – nb) = nRT.
- Kinetic Energy: KE = 3/2 nRT (per mole).
- Root Mean Square Speed: v_rms = √(3RT/M), where M = molar mass.
Applications:
- Competitive Exams:
- UPSC/PCS: Questions on gas laws in environmental contexts (e.g., atmospheric pressure) or industrial processes (e.g., gas storage).
- SSC: Objective questions on Boyle’s/Charles’ laws, critical temperature, or viscosity.
- Descriptive: Explain liquefaction of gases or surface tension in water conservation.
- Real-World:
- Industry: LPG/CNG compression, ammonia liquefaction in refrigeration.
- Environment: Gas behavior in climate models, vapor pressure in weather forecasting.
- Technology: Viscosity in lubricants, surface tension in inkjet printing.
- Exam Tips:
- Master gas law calculations and van der Waals corrections for numericals.
- Link gas behavior to environmental science for mains.
Diagram (Textual Description):
- Boyle’s Law: A graph of pressure (y-axis) vs. volume (x-axis), showing a hyperbolic curve (PV = constant). Mark points: high P, low V; low P, high V. Include a piston-cylinder to show compression.
Class 12: Coordination Compounds
Detailed Concepts:
- Coordination Compound: Complex with a central metal ion/atom surrounded by ligands (molecules/ions donating electron pairs).
- Key Terms:
- Ligand: Donates lone pair to metal (e.g., NH₃, Cl⁻, H₂O). Types:
- Monodentate: One binding site (e.g., CN⁻, NH₃).
- Polydentate: Multiple binding sites (e.g., ethylenediamine (en), bidentate; EDTA, hexadentate).
- Coordination Number (CN): Number of ligand bonds to metal (e.g., CN = 6 in [Co(NH₃)₆]³⁺).
- Oxidation State: Charge on metal after ligand removal (e.g., [Fe(CN)₆]⁴⁻, Fe = +2).
- Coordination Sphere: Central metal + ligands, enclosed in square brackets (e.g., [Co(NH₃)₅Cl]²⁺).
- Ligand: Donates lone pair to metal (e.g., NH₃, Cl⁻, H₂O). Types:
- Nomenclature (IUPAC):
- Name ligands first (anionic, neutral, cationic in order).
- Use prefixes (di-, tri-) for multiple ligands.
- Name metal with oxidation state in parentheses.
- Example: [Co(NH₃)₅Cl]Cl₂ → Pentaamminechlorocobalt(III) chloride.
- Isomerism:
- Structural:
- Ionization: Different ions produced (e.g., [Co(NH₃)₅SO₄]Br vs. [Co(NH₃)₅Br]SO₄).
- Linkage: Ligand binds via different atoms (e.g., NO₂⁻ as nitro (-NO₂) or nitrito (-ONO)).
- Coordination: Different ligand-metal arrangements (e.g., [Co(NH₃)₆][Cr(CN)₆] vs. [Cr(NH₃)₆][Co(CN)₆]).
- Hydrate: Water as ligand or crystal water (e.g., [Cr(H₂O)₆]Cl₃ vs. [Cr(H₂O)₅Cl]Cl₂·H₂O).
- Stereoisomerism:
- Geometrical: Cis (adjacent) vs. trans (opposite) in octahedral/square planar complexes (e.g., [Co(NH₃)₄Cl₂]⁺: cis vs. trans).
- Optical: Non-superimposable mirror images (e.g., [Co(en)₃]³⁺, chiral).
- Structural:
- Valence Bond Theory (VBT):
- Metal uses hybrid orbitals to bond with ligands.
- Example: [Ni(CN)₄]²⁻ (Ni²⁺, d⁸): dsp² hybridization, square planar, diamagnetic.
- Example: [NiCl₄]²⁻ (Ni²⁺, d⁸): sp³ hybridization, tetrahedral, paramagnetic.
- Crystal Field Theory (CFT):
- Ligands split d-orbitals of metal into t₂g and e_g (octahedral) or t₂ and e (tetrahedral).
- Crystal Field Splitting Energy (Δo): Energy difference, depends on ligand strength (spectrochemical series: I⁻ < Cl⁻ < H₂O < NH₃ < en < CN⁻).
- High-Spin vs. Low-Spin: Depends on Δo vs. pairing energy (e.g., [Fe(H₂O)₆]²⁺: high-spin, [Fe(CN)₆]⁴⁻: low-spin).
- Magnetic Properties: Paramagnetic (unpaired electrons) vs. diamagnetic (all paired). Example: [Fe(H₂O)₆]³⁺ has 5 unpaired electrons, paramagnetic.
Formulas:
- Oxidation State: Metal charge + Ligand charges = Complex charge.
- Crystal Field Splitting: Δo for octahedral, Δt = 4/9 Δo for tetrahedral.
- Magnetic Moment: μ = √(n(n+2)) BM, where n = number of unpaired electrons.
- Coordination Number and Geometry:
- CN = 4: Tetrahedral (sp³) or square planar (dsp²).
- CN = 6: Octahedral (d²sp³ or sp³d²).
Applications:
- Competitive Exams:
- UPSC/PCS: Questions on coordination compounds in catalysis (e.g., Wilkinson’s catalyst) or biology (e.g., hemoglobin).
- SSC: Objective questions on nomenclature, isomerism, or magnetic properties.
- Descriptive: Explain role of coordination compounds in chemotherapy (e.g., cisplatin) or metal extraction.
- Real-World:
- Medicine: Cisplatin ([Pt(NH₃)₂Cl₂]) for cancer treatment.
- Industry: Ni purification via [Ni(CO)₄], EDTA in water softening.
- Biology: Chlorophyll (Mg complex), hemoglobin (Fe complex).
- Exam Tips:
- Master nomenclature and isomerism for objective questions.
- Link CFT to magnetic properties and colors for mains.
Diagram (Textual Description):
- Octahedral Splitting in CFT: Show a d-orbital energy diagram. Five d-orbitals (d_xy, d_xz, d_yz, d_x²-y², d_z²) split into lower t₂g (3 orbitals) and higher e_g (2 orbitals) in octahedral field. Label Δo, show electron filling for [Fe(H₂O)₆]³⁺ (high-spin, 5 unpaired electrons).