Revision Notes: Class 12 Chemistry Chapter 1 Solutions

A condensed, formula-first recap of every key definition and result in Chapter 1: Solutions — ideal for last-minute board exam revision. These Class 12 Chemistry Chapter 1 notes are ideal for quick revision just before exams.

Last Updated: October 24, 2026

Common Mistakes Students Make in Solutions

Revision Notes: Class 12 Chemistry Chapter 1 – 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

Frequently Asked Questions

What is the difference between molarity and molality as concentration terms?
Molarity is the number of moles of solute per litre of solution and changes with temperature since volume expands or contracts, while molality is moles of solute per kilogram of solvent and remains unaffected by temperature changes.

What does Raoult Law state about the vapour pressure of a solution?
Raoult Law states that the partial vapour pressure of each volatile component in a solution is directly proportional to its mole fraction in the solution, forming the basis for understanding ideal and non-ideal solutions.

What to Revise First (and Last) in This Chapter

Written by Satish

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