Revision Notes: Class 12 Chemistry Chapter 2 Solutions

A condensed, formula-first recap of every key definition and result in Chapter 2: Solutions — ideal for last-minute board exam revision. These Class 12 Chemistry Chapter 2 notes are ideal for quick revision just before exams.

Revision Notes: Class 12 Chemistry Chapter 2 – Solutions

• Solution = homogeneous mixture of two or more non-reacting substances; solute (lesser amount) + solvent (greater amount).
• 9 types of solutions based on physical state of solute/solvent: gas–gas, gas–liquid, gas–solid, liquid–gas, liquid–liquid, liquid–solid, solid–gas, solid–liquid, solid–solid.
• Mass percentage (w/w %) = (mass of component ÷ total mass of solution) × 100.
• Volume percentage (v/v %) = (volume of component ÷ total volume of solution) × 100.
• Mass by volume percentage (w/v %) = (mass of solute in g ÷ volume of solution in mL) × 100.
• Parts per million (ppm) = (mass/moles/volume of component ÷ total mass/moles/volume) × 10^6.
• Mole fraction, x_A = n_A ÷ (n_A+n_B); sum of all mole fractions in a solution = 1.
• Molarity (M) = moles of solute ÷ volume of solution (L); unit mol L−1; temperature-dependent (volume changes with T).
• Molality (m) = moles of solute ÷ mass of solvent (kg); unit mol kg−1; temperature-independent.
• Solubility of a solid in a liquid depends on temperature and, per Le Chatelier’s principle, increases with T if dissolution is endothermic and decreases if exothermic.
• Solubility of a gas in a liquid decreases with increasing temperature (dissolution of gases is exothermic) and increases with increasing pressure.
• Henry’s law: p = K_H · x, where p = partial pressure of gas, x = mole fraction of gas in solution, K_H = Henry’s law constant (depends on gas, solvent, and temperature); higher K_H = lower solubility at a given pressure.
• Applications of Henry’s law: carbonated drinks (CO2 solubility under pressure), scuba diving (N2 narcosis / decompression sickness), high-altitude anoxia (low O2 partial pressure).
• Raoult’s law (volatile solute): p1 = x1 p°1, p2 = x2 p°2; total pressure p_total = x1 p°1 + x2 p°2.
• Raoult’s law (non-volatile solute): relative lowering of vapour pressure (p°−p)/p° = x2 (mole fraction of solute); for dilute solutions, x2 ≈ n2/n1.
• Ideal solutions: obey Raoult’s law at all concentrations; ΔmixH = 0, ΔmixV = 0 (e.g., benzene + toluene, n-hexane + n-heptane).
• Non-ideal solutions show deviation from Raoult’s law: positive deviation (A–B forces weaker than A–A/B–B; ΔmixH > 0; p_total > ideal; e.g. ethanol+acetone, water+ethanol) and negative deviation (A–B forces stronger; ΔmixH < 0; p_total < ideal; e.g. chloroform+acetone, HNO3+water).
• Azeotropes: constant-boiling mixtures with the same composition in liquid and vapour phases; cannot be separated by simple distillation. Minimum-boiling azeotropes form from solutions with positive deviation; maximum-boiling azeotropes from negative deviation.
• Colligative properties depend only on the number of solute particles, not their identity: relative lowering of vapour pressure, elevation of boiling point, depression of freezing point, osmotic pressure.
• Elevation of boiling point: ΔTb = i · Kb · m, where Kb = molal elevation (ebullioscopic) constant, m = molality.
• Depression of freezing point: ΔTf = i · Kf · m, where Kf = molal depression (cryoscopic) constant.
• Kb and Kf are solvent-specific constants with units K kg mol−1; for water, Kf = 1.86 K kg mol−1 (standard/commonly used value).
• Osmosis: spontaneous flow of solvent from a region of lower solute concentration to higher solute concentration through a semi-permeable membrane.
• Osmotic pressure: π = CRT (van’t Hoff equation for dilute solutions), where C = molar concentration, R = gas constant, T = temperature (K); equivalently π = (n/V)RT; for an ionising/associating solute, π = iCRT.
• Isotonic solutions have equal osmotic pressure at a given temperature; a solution is hypertonic/hypotonic relative to another if its osmotic pressure is higher/lower.
• Reverse osmosis: applying pressure greater than osmotic pressure forces solvent to flow from concentrated to dilute solution through a semi-permeable membrane (used in seawater desalination).
• Van’t Hoff factor, i = normal molar mass ÷ observed (abnormal) molar mass = observed colligative property ÷ calculated (normal) colligative property = (total number of moles of particles after dissociation/association) ÷ (number of moles of formula units dissolved).
• Dissociation (solute breaks into n ions): i = 1 + (n−1)α, where α = degree of dissociation; i > 1 (e.g. NaCl, KCl, electrolytes).
• Association (n solute units combine into one aggregate): i = 1 − α + α/n = 1 − α(1−1/n); i < 1 (e.g. carboxylic acids dimerising in benzene).
• Modified colligative-property equations with van’t Hoff factor: ΔTb = i Kb m; ΔTf = i Kf m; π = i CRT; (p°−p)/p° = i x2.

More on This Chapter

NCERT Solutions | Extra HOTS Questions | Class 12 Chemistry Book

Written by Satish

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