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Solutions – Class 12 Chemistry Chapter 1 Notes & NCERT PDF

Class 12ChemistryChapter 1NCERT book: Chemistry Part I - Class 12

Notes and a simple summary of Chapter 1 of the NCERT Class 12 Chemistry book, with the official chapter PDF, flashcards, an MCQ quiz and an AI tutor for your doubts.

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Chapter 1: Solutions

Introduction

Most substances we meet in daily life are not pure — they are mixtures. Air, sea water, brass, medicines and even the blood in our veins are solutions. This unit studies homogeneous mixtures of two or more components, how we express their composition, how gases and solids dissolve, how solutions behave when they boil or freeze, and how these behaviours let us find the molar mass of an unknown solute.

A solution is a homogeneous mixture of two or more components whose composition and properties are uniform throughout. The component present in the largest quantity is the solvent; the others are solutes.

Key Concepts

1.1 Types of Solutions

A binary solution has two components. Depending on the physical state of solvent and solute there are nine types, grouped as:

  • Gaseous solutions — e.g. mixture of oxygen and nitrogen gases; chloroform mixed with nitrogen gas; camphor in nitrogen gas.
  • Liquid solutions — e.g. oxygen dissolved in water; ethanol in water; glucose in water.
  • Solid solutions — e.g. hydrogen in palladium; amalgam of mercury with sodium; copper dissolved in gold.

    • 1.2 Expressing Concentration of Solutions

  • Mass percentage (w/w) = (mass of component ÷ total mass of solution) × 100. A 10% glucose solution has 10 g glucose in 90 g water.
  • Volume percentage (V/V) — used for liquid–liquid solutions; e.g. 35% ethylene glycol (an antifreeze) in water lowers the freezing point of water to 255 K.
  • Mass by volume percentage (w/V) — grams of solute per 100 mL of solution; common in medicine and pharmacy.
  • Parts per million (ppm) = (parts of component ÷ total parts) × 10⁶; used for trace amounts, e.g. dissolved oxygen or pollutants.
  • Mole fraction (x) = moles of component ÷ total moles. For a binary solution, x₁ + x₂ = 1.
  • Molarity (M) = moles of solute ÷ volume of solution in litres.
  • Molality (m) = moles of solute ÷ mass of solvent in kg.
    • Mass %, ppm, mole fraction and molality are independent of temperature; molarity depends on temperature because volume changes with temperature.

    1.3 Solubility

    Solubility is the maximum amount of a substance that can be dissolved in a specified amount of solvent at a specified temperature.

  • Solids in liquids — "like dissolves like": polar solutes dissolve in polar solvents, non-polar in non-polar. A solution in which no more solute can dissolve at a given temperature and pressure is saturated. If dissolution is endothermic (ΔH_sol > 0), solubility increases with temperature (Le Chatelier). Pressure has no significant effect because solids and liquids are highly incompressible.
  • Gases in liquids — solubility increases with pressure and decreases with rise in temperature (dissolution of gases is exothermic). This is why aquatic species are more comfortable in cold water.

    • Henry's law: the partial pressure of a gas in the vapour phase is proportional to its mole fraction in the solution:

    p = K_H · x

    A higher K_H means lower solubility at the same partial pressure. Applications given in the book:

  • Soft drinks and soda water are sealed under high pressure to increase the solubility of CO₂.
  • Scuba divers breathe air at high pressure; on rapid ascent dissolved gases form bubbles in blood — the painful condition called bends. Tanks are filled with air diluted with helium (11.7% helium, 56.2% nitrogen and 32.1% oxygen).
  • At high altitudes the low partial pressure of oxygen causes low blood-oxygen — the condition called anoxia.

    • 1.4 Vapour Pressure of Liquid Solutions — Raoult's Law

    For a solution of two volatile liquids, the partial vapour pressure of each component is directly proportional to its mole fraction:

    p₁ = p₁⁰ x₁ and p₂ = p₂⁰ x₂; p_total = p₁⁰ + (p₂⁰ − p₁⁰) x₂

    The mole fraction of a component in the vapour phase is yᵢ = pᵢ / p_total; the vapour is richer in the more volatile component. Raoult's law is a special case of Henry's law with K_H = p⁰. For a non-volatile solute, only the solvent contributes: p = p₁⁰ x₁.

    1.5 Ideal and Non-ideal Solutions

  • Ideal solutions obey Raoult's law over the entire range of concentration; ΔH_mix = 0 and ΔV_mix = 0. A–B interactions are similar to A–A and B–B. Examples: n-hexane + n-heptane, bromoethane + chloroethane, benzene + toluene.
  • Positive deviation — A–B interactions weaker than A–A/B–B; vapour pressure higher than predicted (e.g. ethanol + acetone; carbon disulphide + acetone).
  • Negative deviation — A–B interactions stronger (hydrogen bonding); vapour pressure lower (e.g. phenol + aniline; chloroform + acetone).
  • Azeotropes boil at constant temperature with the same composition in liquid and vapour. Large positive deviation → minimum boiling azeotrope (ethanol–water, about 95% ethanol by volume). Large negative deviation → maximum boiling azeotrope (nitric acid–water, about 68% nitric acid by mass, boiling at 393.5 K).

    • 1.6 Colligative Properties

    Properties that depend only on the number of solute particles (not their nature):

  • Relative lowering of vapour pressure: (p₁⁰ − p₁)/p₁⁰ = x₂ ≈ (w₂ M₁)/(M₂ w₁)
  • Elevation of boiling point: ΔT_b = K_b · m (K_b = ebullioscopic constant; 0.52 K kg mol⁻¹ for water)
  • Depression of freezing point: ΔT_f = K_f · m (K_f = cryoscopic constant; 1.86 K kg mol⁻¹ for water)
  • Osmotic pressure: Π = CRT = (n₂/V)RT. Osmosis is the flow of solvent through a semipermeable membrane from pure solvent (or dilute solution) to concentrated solution. Solutions with equal osmotic pressure are isotonic; 0.9% (mass/volume) NaCl is normal saline, isotonic with blood cells. In a hypertonic solution cells shrink; in a hypotonic solution they swell.
  • Reverse osmosis — applying pressure larger than the osmotic pressure pushes pure solvent out through the membrane; used in desalination of sea water.

    • Osmotic pressure is preferred for finding molar masses of macromolecules like proteins and polymers, because it is measured at room temperature and gives a sizeable value even for very dilute solutions.

    1.7 Abnormal Molar Masses and van't Hoff Factor

    Solutes that dissociate (e.g. KCl → K⁺ + Cl⁻) give more particles and show lower-than-expected molar mass; solutes that associate (e.g. ethanoic acid or benzoic acid dimerising in benzene) give fewer particles and higher molar mass. The van't Hoff factor:

    i = normal molar mass / abnormal molar mass = observed colligative property / calculated colligative property

    Modified equations: ΔT_b = i K_b m, ΔT_f = i K_f m, Π = i CRT. For dissociation i > 1; for association i < 1.

    Worked Examples

  • Molarity (Ex. 1.2): 5 g NaOH in 450 mL. Moles = 5/40 = 0.125 mol → M = 0.125 × 1000/450 = 0.278 M.
  • Henry's law (Ex. 1.4): N₂ at 0.987 bar, K_H = 76.48 kbar. x = 0.987/76480 = 1.29 × 10⁻⁵; 1 L water has 55.5 mol, so n ≈ 1.29 × 10⁻⁵ × 55.5 = 7.16 × 10⁻⁴ mol = 0.716 mmol.
  • Raoult's law (Ex. 1.5): x(CH₂Cl₂) = 0.688, p⁰ values 415 and 200 mm Hg. p_total = 200 + (415 − 200) × 0.688 = 347.9 mm Hg; y(CH₂Cl₂) = 285.5/347.9 = 0.82 — vapour is richer in the more volatile CH₂Cl₂.
  • Relative lowering (Ex. 1.6): 0.5 g solute in 39 g benzene; p falls from 0.850 to 0.845 bar → M₂ = 170 g mol⁻¹.
  • Boiling-point elevation (Ex. 1.8): 1.80 g solute in 90 g benzene raises b.p. by 0.88 K; K_b = 2.53 → M₂ = (2.53 × 1.8 × 1000)/(0.88 × 90) ≈ 58 g mol⁻¹.
  • Freezing-point depression (Ex. 1.10): 1.00 g in 50 g benzene, ΔT_f = 0.40 K, K_f = 5.12 → M₂ = 256 g mol⁻¹.
  • Osmotic pressure (Ex. 1.11): 1.26 g protein in 200 cm³, Π = 2.57 × 10⁻³ bar at 300 K → M₂ = wRT/(ΠV) ≈ 6.1 × 10⁴ g mol⁻¹ (the book prints 61,022 g mol⁻¹; direct calculation gives about 61,039).
  • van't Hoff factor (Ex. 1.13): acetic acid, calculated ΔT_f = 1.86 × 0.0106 = 0.0197 K, observed 0.0205 K → i = 0.0205/0.0197 ≈ 1.041, showing slight dissociation.

    • Formula Sheet

  • x₁ + x₂ = 1; M = n/V(L); m = n/kg solvent
  • Henry: p = K_H x
  • Raoult: pᵢ = pᵢ⁰ xᵢ; p_total = p₁⁰ + (p₂⁰ − p₁⁰) x₂; yᵢ = pᵢ/p_total
  • Relative lowering: Δp/p₁⁰ = x₂
  • ΔT_b = i K_b m; ΔT_f = i K_f m; Π = i CRT
  • M₂ = K_b w₂ × 1000/(ΔT_b w₁); M₂ = K_f w₂ × 1000/(ΔT_f w₁); M₂ = w₂RT/(ΠV)

    • Board Exam Focus

  • Numericals on molarity/molality/mole fraction, Henry's law, Raoult's law and all four colligative properties — show units and substitute step by step.
  • Define with example: ideal solution, azeotrope, osmotic pressure, isotonic solution, van't Hoff factor.
  • Reasoning: Why is osmotic pressure preferred for proteins? (room temperature measurement, sizeable value for dilute solutions, biomolecules unstable at high temperature.) Why do we add salt to icy roads / ethylene glycol to car radiators? (freezing-point depression.)
  • Answer outline — positive vs negative deviation: state interaction strength, effect on vapour pressure, sign of ΔH_mix and ΔV_mix, one example each, and the type of azeotrope formed.

    • Key Terms

  • Solute (विलेय), Solvent (विलायक)
  • Molarity (मोलरता), Molality (मोललता)
  • Mole fraction (मोल अंश)
  • Henry's law (हेनरी का नियम)
  • Raoult's law (राउल्ट का नियम)
  • Azeotrope (स्थिरक्वाथी मिश्रण)
  • Colligative property (अणुसंख्य गुणधर्म)
  • Osmosis (परासरण), Osmotic pressure (परासरण दाब)
  • Reverse osmosis (प्रतिलोम परासरण)
  • Isotonic (समपरासरी)
  • Ebullioscopic constant (क्वथनांकमापी स्थिरांक)
  • Cryoscopic constant (हिमांकमापी स्थिरांक)
  • van't Hoff factor (वान्ट हॉफ गुणक)

    • Common Mistakes

  • Using mass of solution instead of mass of solvent in molality.
  • Forgetting that a higher K_H means lower gas solubility.
  • Writing ΔT_f = K_f × molarity — it must be molality.
  • Forgetting the van't Hoff factor for electrolytes (NaCl gives i ≈ 2).
  • Mixing units of R: use 0.083 L bar mol⁻¹ K⁻¹ with bar and litres, 8.314 J mol⁻¹ K⁻¹ with SI units.

    • 💡 Key Learning Points

    • ✓Express concentration as mass %, ppm, mole fraction, molarity and molality, and know which depend on temperature.
    • ✓Apply Henry's law and Raoult's law, including vapour-phase composition.
    • ✓Distinguish ideal and non-ideal solutions and the two kinds of azeotropes.
    • ✓Use the four colligative properties to find molar masses of solutes.
    • ✓Explain abnormal molar masses using the van't Hoff factor.

    👨‍🏫 Teaching Tips

    • →Demonstrate a soda bottle opening to discuss Henry's law and CO₂ solubility.
    • →Place a raw mango slice in salt solution and a wilted flower in water to show osmosis.
    • →Plot p_total versus x₂ on the board for ideal, positive and negative deviation cases.
    • →Give a colligative-property table: ask students to decide which uses molality vs molarity.
    • →Run a timed numericals drill on Examples 1.8, 1.10 and 1.11 with unit checking.
    • →Discuss why salt is spread on icy roads and ethylene glycol added to car radiators.
    • →Use a concept map linking van't Hoff factor to dissociation and association.

    📋 Assessment Questions

    1. Calculate molality, molarity and mole fraction from given masses and densities.
    2. Find the number of millimoles of a gas dissolved using Henry's law.
    3. Explain positive and negative deviation with examples and azeotrope types.
    4. Determine the molar mass of a solute from ΔT_b, ΔT_f or Π.
    5. Calculate and interpret the van't Hoff factor for a weak electrolyte.

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    Keep learning

    All chapters in Chemistry Part I - Class 12

    1. 1. Solutions
    2. 2. Electrochemistry
    3. 3. Chemical Kinetics
    4. 4. The d- and f-Block Elements
    5. 5. Coordination Compounds