Chapter 1: Solutions Long Questions | chemistry Class 12 CBSE

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Chapter 1: Solutions Long Questions | chemistry Class 12 CBSE | ByteNotes.in

Long Answer

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Question 1

Long Form

Compare the various methods of expressing the concentration of a solution and explain which methods are temperature-dependent and which are not.

Mass percentage, volume percentage, mass by volume percentage, and parts per million express concentration in terms of mass or volume ratios; mole fraction, molality, and molarity express it in molar terms.
Mass percentage (w/w), ppm, and mole fraction are defined using mass, which does not change with temperature, making them temperature-independent.
Molality is defined as moles of solute per kilogram of solvent and is also temperature-independent since it uses mass of solvent.
Molarity is defined as moles of solute per litre of solution; since the volume of liquid changes with temperature, molarity is temperature-dependent.
For reactions requiring precise stoichiometry at variable temperatures, molality or mole fraction is preferred; molarity is convenient for volumetric analysis at a fixed temperature.

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Question 2

Long Form

Explain the effect of temperature and pressure on the solubility of a solid and a gas in a liquid, applying Le Chatelier's Principle.

For solids in liquids, pressure has negligible effect because both phases are highly incompressible and remain unaffected by pressure changes.
Temperature significantly affects solid solubility: if dissolution is endothermic (ΔsolH > 0), solubility increases with temperature; if exothermic (ΔsolH < 0), solubility decreases — consistent with Le Chatelier's Principle.
For gases in liquids, increasing pressure increases the number of gaseous particles per unit volume above the solution, forcing more gas into solution (Henry's law); solubility increases with pressure.
Dissolution of gas in liquid is an exothermic process (similar to condensation); by Le Chatelier's Principle, increasing temperature shifts equilibrium backward, decreasing gas solubility.
This explains why aquatic organisms are more comfortable in cold water (higher dissolved oxygen) and why CO2 effervesces out of a warm soft drink more readily.

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Question 3

Long Form

Derive the expression for molar mass of a solute using elevation of boiling point, and explain why a solution always boils at a higher temperature than the pure solvent.

A solution boils when its vapour pressure equals atmospheric pressure. Since a non-volatile solute lowers the solvent's vapour pressure, a higher temperature is required to raise it to atmospheric pressure, resulting in a higher boiling point.
The elevation ΔTb = Tb − Tb° is directly proportional to molality for dilute solutions: ΔTb = Kb × m, where Kb is the ebullioscopic constant.
If w2 g of solute (molar mass M2) is dissolved in w1 g of solvent, molality m = (w2 × 1000) / (M2 × w1).
Substituting: ΔTb = Kb × (w2 × 1000) / (M2 × w1), rearranging to M2 = (1000 × w2 × Kb) / (ΔTb × w1).
By measuring ΔTb experimentally with known values of w1, w2, and Kb of the solvent, the molar mass M2 of the unknown solute can be determined.

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Question 4

Long Form

Describe the colligative property of osmotic pressure, its mathematical treatment, and explain why this method is particularly advantageous for determining molar masses of large molecules.

Osmotic pressure (Π) is the excess pressure that must be applied to a solution to prevent the passage of solvent through a semipermeable membrane into the solution.
For dilute solutions: Π = CRT = (n2/V)RT, where n2 is moles of solute, V is volume of solution, R is the gas constant, and T is temperature.
Molar mass can be determined: M2 = w2RT / (ΠV), where w2 is the mass of solute.
Osmotic pressure measurements are conducted near room temperature, making them suitable for thermally unstable biomolecules like proteins and enzymes.
Osmotic pressure is very large even for very dilute solutions (e.g., a 1 mg/mL polymer solution generates a measurable pressure), and polymer solutions use molarity instead of molality, avoiding weighing errors; hence this method gives accurate M2 values for macromolecules.

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Question 5

Long Form

Explain ideal and non-ideal solutions, the molecular basis for positive and negative deviations from Raoult's law, and give examples from the chapter for each.

Ideal solutions obey Raoult's law at all concentrations, with ΔmixH = 0 and ΔmixV = 0, because A–A, B–B, and A–B interactions are nearly equal (e.g., benzene-toluene, n-hexane-n-heptane).
Positive deviations occur when A–B interactions are weaker than A–A and B–B; molecules escape more easily, giving vapour pressure higher than predicted; ΔmixH > 0. Example: ethanol-acetone (acetone breaks H-bonds between ethanol molecules); also acetone-CS2.
Negative deviations occur when A–B interactions are stronger than A–A and B–B; molecules escape less easily, vapour pressure is lower than predicted; ΔmixH < 0. Example: chloroform-acetone (chloroform forms H-bond with carbonyl of acetone); also phenol-aniline.
Very large positive deviations form minimum boiling azeotropes (e.g., ethanol-water, ~95% ethanol); very large negative deviations form maximum boiling azeotropes (e.g., HNO3-water, 68% HNO3, bp 393.5 K).
Azeotropic mixtures boil at constant temperature with the same liquid and vapour composition and cannot be separated by fractional distillation.

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Question 6

Long Form

Analyse the concept of abnormal molar mass and the van't Hoff factor, giving examples of both association and dissociation of solutes.

Abnormal molar mass occurs when the experimentally determined molar mass from colligative properties differs from the theoretical value due to association or dissociation of solute particles.
For electrolytes that dissociate (e.g., KCl → K⁺ + Cl⁻), the number of particles doubles, so the observed molar mass is half the expected value; modified colligative properties are larger than calculated.
For associating solutes (e.g., acetic acid in benzene forms dimers via H-bonds: 2CH3COOH ⇌ (CH3COOH)2), particle count halves; observed molar mass is about twice the expected; colligative properties are smaller.
Van't Hoff factor i = Normal molar mass / Abnormal molar mass = Observed colligative property / Calculated colligative property. For KCl (dilute), i ≈ 2; for acetic acid in benzene, i ≈ 0.5; for K2SO4, i → 3.
Modified equations: ΔTb = i Kb m; ΔTf = i Kf m; Π = i n2RT/V; allowing correct calculation of colligative properties even for electrolytes.

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Question 7

Long Form

Explain the process of osmosis, the role of semipermeable membranes, and discuss the biological and practical significance of osmosis using specific examples from the chapter.

Osmosis is the spontaneous flow of solvent molecules through a semipermeable membrane from a region of lower solute concentration to higher solute concentration; the membrane allows small solvent molecules through but restricts larger solute molecules.
Biological membranes (e.g., pig's bladder, parchment) and synthetic membranes (cellophane) act as semipermeable membranes; solvent flows from dilute to concentrated side until equilibrium or until counterbalancing pressure is applied.
Practical significance: raw mangoes shrivel in brine because water leaves by osmosis; wilted flowers revive in fresh water as water enters cells; carrots placed in water regain firmness.
Medical significance: 0.9% (mass/volume) NaCl is isotonic with blood plasma; intravenous injections must match blood plasma concentration to prevent cell shrinkage (hypertonic) or swelling (hypotonic).
Osmosis enables water movement from soil into plant roots and upward in plants; preservation of meat by salting and fruits by adding sugar kills bacteria through osmosis, as bacteria lose water and shrivel.

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Long Answer

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Question 8

Long Form

Describe how colligative properties — specifically depression of freezing point — can be used to determine the degree of dissociation of a weak electrolyte in solution.

For a weak electrolyte like acetic acid in water: CH3COOH ⇌ H⁺ + CH3COO⁻; if x is degree of dissociation, n(1−x) moles remain undissociated and 2nx moles of ions are formed, giving total n(1+x) moles.
The van't Hoff factor i = n(1+x)/n = 1 + x, so degree of dissociation x = i − 1.
Experimentally, ΔTf (observed) is measured; ΔTf (calculated) is determined assuming no dissociation; i = ΔTf(observed) / ΔTf(calculated).
From i, x is obtained; knowing concentrations, the dissociation constant Ka = [H⁺][CH3COO⁻]/[CH3COOH] = (nx)²/[n(1−x)] can be calculated.
This method connects macroscopic measurements (freezing point depression) to microscopic equilibrium constants, demonstrating the power of colligative properties in physical chemistry.

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Question 9

Long Form

Explain Henry's law in detail, discuss the significance of Henry's law constant KH, and describe three important applications of this law.

Henry's law states that at constant temperature, the solubility of a gas in a liquid is directly proportional to its partial pressure above the liquid: p = KH × x, where x is the mole fraction of the gas in solution.
KH is a characteristic constant for a specific gas-solvent combination at a given temperature; higher KH means lower solubility (e.g., N2 has KH = 76.48 kbar while CO2 has KH = 1.67 kbar in water at ~298 K, meaning CO2 is far more soluble).
Application 1: Soft drinks and soda water — CO2 is dissolved under high pressure; when the bottle is opened, pressure drops and CO2 effervesces out.
Application 2: Scuba diving — at high underwater pressure, atmospheric gases dissolve more in blood; on rapid ascent, pressure drops and nitrogen forms bubbles causing painful 'bends'; divers use helium-diluted air to minimize this risk.
Application 3: High altitude — low partial pressure of O2 at high altitudes reduces O2 concentration in blood, causing anoxia (weakness, impaired thinking) in climbers.

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Long Answer

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Hard

Question 10

Long Form

Describe the vapour pressure of liquid-liquid solutions using Raoult's law and explain how the composition of vapour phase differs from the liquid phase.

For a binary mixture of two volatile liquids (components 1 and 2), Raoult's law gives partial pressures: p1 = x1 p1° and p2 = x2 p2°, where x1, x2 are mole fractions in liquid and p1°, p2° are vapour pressures of pure components.
Total pressure ptotal = p1 + p2 = p1° + (p2° − p1°)x2; a linear plot of ptotal vs. mole fraction lies between p1° and p2°.
The vapour phase composition is given by Dalton's law: y1 = p1/ptotal and y2 = p2/ptotal, where y1, y2 are mole fractions in vapour.
The more volatile component (higher vapour pressure) has a greater mole fraction in the vapour phase than in the liquid phase; at equilibrium, vapour is always richer in the more volatile component.
For example, in CHCl3 (p° = 200 mm Hg) and CH2Cl2 (p° = 415 mm Hg) mixture, the vapour phase is richer in CH2Cl2 (more volatile), as demonstrated by y(CH2Cl2) = 0.82 while x(CH2Cl2) = 0.688.

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Chapter 1: Solutions | Long Questions

Compare the various methods of expressing the concentration of a solution and explain which methods are temperature-dependent and which are not.

Explain the effect of temperature and pressure on the solubility of a solid and a gas in a liquid, applying Le Chatelier's Principle.

Derive the expression for molar mass of a solute using elevation of boiling point, and explain why a solution always boils at a higher temperature than the pure solvent.

Describe the colligative property of osmotic pressure, its mathematical treatment, and explain why this method is particularly advantageous for determining molar masses of large molecules.

Explain ideal and non-ideal solutions, the molecular basis for positive and negative deviations from Raoult's law, and give examples from the chapter for each.

Analyse the concept of abnormal molar mass and the van't Hoff factor, giving examples of both association and dissociation of solutes.

Explain the process of osmosis, the role of semipermeable membranes, and discuss the biological and practical significance of osmosis using specific examples from the chapter.

Describe how colligative properties — specifically depression of freezing point — can be used to determine the degree of dissociation of a weak electrolyte in solution.

Explain Henry's law in detail, discuss the significance of Henry's law constant KH, and describe three important applications of this law.

Describe the vapour pressure of liquid-liquid solutions using Raoult's law and explain how the composition of vapour phase differs from the liquid phase.

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