CBSE Class 12 Chemistry Flashcards

24 sample cards from the Chemistry deck

Sampled from the end of the deck, so these are different cards from the ones shown on the syllabus page.

1. How does Raoult's law apply to a solution of a non-volatile solute in a volatile solvent?

   The relative lowering of vapour pressure equals the mole fraction of the solute: (p°_A − p_A)/p°_A = x_B (mole fraction of solute). The solute lowers the solvent's vapour pressure.

2. What is an ideal solution and what are its conditions?

   A solution that obeys Raoult's law over the entire concentration range, with ΔH_mixing = 0 and ΔV_mixing = 0, and where A–B interactions equal A–A and B–B interactions (e.g. benzene + toluene, n-hexane + n-heptane).

3. Distinguish positive and negative deviations from Raoult's law.

   Positive deviation: vapour pressure higher than expected, ΔH_mix > 0, ΔV_mix > 0, weaker A–B interactions (e.g. ethanol + acetone). Negative deviation: vapour pressure lower than expected, ΔH_mix < 0, ΔV_mix < 0, stronger A–B interactions (e.g. chloroform + acetone).

4. What are colligative properties? Name the four.

   Properties that depend only on the number of solute particles, not their nature: (1) relative lowering of vapour pressure, (2) elevation of boiling point, (3) depression of freezing point, and (4) osmotic pressure.

5. Write the formulas for elevation of boiling point and depression of freezing point.

   ΔT_b = K_b × m and ΔT_f = K_f × m, where K_b = molal elevation (ebullioscopic) constant, K_f = molal depression (cryoscopic) constant, and m = molality.

6. Define osmotic pressure and write its formula.

   Osmotic pressure (π) is the excess pressure that must be applied to a solution to just stop osmosis (flow of solvent through a semipermeable membrane). π = CRT = (n/V)RT, where C is molar concentration.

7. What are isotonic solutions?

   Two solutions having the same osmotic pressure at a given temperature; there is no net movement of solvent between them across a semipermeable membrane.

8. Define the van't Hoff factor (i) and give its formula.

   i = (normal molar mass)/(abnormal/observed molar mass) = (observed colligative property)/(calculated colligative property) = (total moles of particles after association/dissociation)/(moles before). It accounts for association or dissociation of solute.

9. What value of van't Hoff factor (i) indicates dissociation versus association?

   i > 1 indicates dissociation of solute (e.g. NaCl → i ≈ 2); i < 1 indicates association of solute (e.g. dimerization of acetic acid in benzene); i = 1 means no association/dissociation.

10. Distinguish between a galvanic (voltaic) cell and an electrolytic cell.

    A galvanic cell converts chemical energy into electrical energy via a spontaneous redox reaction (ΔG negative). An electrolytic cell uses electrical energy to drive a non-spontaneous reaction (electrolysis).

11. In a galvanic cell, which electrode is the anode and which is the cathode, and what charge does each carry?

    Oxidation occurs at the anode (negative terminal); reduction occurs at the cathode (positive terminal). Electrons flow from anode to cathode in the external circuit.

12. What is the role of a salt bridge in a galvanic cell?

    It completes the circuit and maintains electrical neutrality in the two half-cells by allowing ion flow, preventing accumulation of charge that would stop the cell reaction.

13. Define standard electrode potential (E°).

    The electrode potential measured relative to the standard hydrogen electrode (SHE) under standard conditions (1 M ion concentration, 1 bar gas pressure, 298 K). SHE is assigned E° = 0.00 V.

14. How is the standard cell potential (EMF) of a galvanic cell calculated?

    E°_cell = E°_cathode − E°_anode (both as reduction potentials) = E°_right − E°_left. A positive E°_cell indicates a spontaneous reaction.

15. Write the Nernst equation for a single electrode (M^n+ + ne⁻ → M).

    E = E° − (RT/nF) ln([reduced]/[oxidized]); at 298 K, E = E° − (0.0591/n) log(1/[M^n+]) for the electrode M^n+/M.

16. Write the Nernst equation for a cell at 298 K and relate E°_cell to equilibrium constant.

    E_cell = E°_cell − (0.0591/n) log Q. At equilibrium E_cell = 0 and Q = K, giving E°_cell = (0.0591/n) log K.

17. State the relation between standard cell potential and Gibbs free energy change.

    ΔG° = −nFE°_cell, where n = number of moles of electrons transferred and F = Faraday constant (96500 C/mol). Negative ΔG° (positive E°) means a spontaneous reaction.

18. Define conductivity (κ) and molar conductivity (Λ_m) of an electrolytic solution.

    Conductivity (κ) is the conductance of a solution in a cell of 1 m length and 1 m² area (κ = 1/ρ). Molar conductivity (Λ_m = κ/c, in S cm² mol⁻¹) is the conducting power of all the ions produced by one mole of electrolyte.

19. How do conductivity and molar conductivity change with dilution?

    Conductivity (κ) decreases with dilution (fewer ions per unit volume); molar conductivity (Λ_m) increases with dilution (more complete dissociation / greater ion mobility).

20. State Kohlrausch's law of independent migration of ions.

    The limiting molar conductivity of an electrolyte is the sum of the individual limiting molar conductivities of its cations and anions: Λ°_m = ν₊ λ°₊ + ν₋ λ°₋. It allows calculation of Λ°_m for weak electrolytes.

21. State Faraday's first law of electrolysis.

    The amount of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte: m = Z × Q = Z × I × t (Z = electrochemical equivalent).

22. State Faraday's second law of electrolysis.

    When the same quantity of electricity is passed through different electrolytes, the masses of substances liberated are proportional to their chemical equivalent weights (equivalent mass = molar mass / valence).

23. Differentiate primary and secondary batteries with examples.

    Primary batteries are non-rechargeable; the reaction occurs only once (e.g. dry cell, mercury cell). Secondary batteries are rechargeable and can be reused (e.g. lead storage battery, nickel-cadmium cell).

24. Describe the electrode reactions of the lead storage battery during discharge.

    Anode: Pb + SO₄²⁻ → PbSO₄ + 2e⁻; Cathode: PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O. The electrolyte is ~38% H₂SO₄; it is recharged by reversing the reactions. A fuel cell (e.g. H₂–O₂) instead converts the energy of fuel combustion directly into electricity.

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