Solutions covers how to express concentration (molarity, molality, mole fraction, mass percent), how solubility of solids and gases responds to temperature and pressure, and Henry’s and Raoult’s laws for vapour pressure of ideal and non-ideal mixtures. Its heart is the four colligative properties, which depend only on the number of solute particles, plus the van’t Hoff factor that corrects them for dissociation and association. It is a high-yield NEET chapter because its formulas reliably appear as numerical problems and the colligative concepts link directly to molar-mass determination.
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Concentration Terms
Concentration says how much solute sits in a given amount of solvent or solution.
Molality & mole fraction are temperature-independent; molarity & normality change with T.
- Mole fraction: x_A + x_B = 1 (moles of one / total moles)
- Mass % = (mass solute / mass solution) × 100; ppm for very dilute
- Molarity ↔ molality conversion needs the density
Solubility & Henry’s Law
Solubility is the max solute that dissolves at a given temperature; gases behave opposite to solids.
K_H is the Henry’s-law constant (pressure units); higher K_H → lower gas solubility.
- Most solids: solubility rises with T (endothermic); pressure barely matters
- Gases: solubility rises with pressure, falls with T (warm soda goes flat)
- K_H increases with T; diluted-air scuba tanks avoid the ‘bends’
Raoult’s Law
Each volatile component’s partial vapour pressure equals its pure value times its mole fraction.
For a non-volatile solute only solvent vaporises: p_solution = p°_solvent · x_solvent.
- A non-volatile solute always lowers vapour pressure (x_solvent < 1)
- Ideal solution obeys Raoult at all x: ΔH_mix = 0, ΔV_mix = 0
- Examples: benzene + toluene, n-hexane + n-heptane
Non-ideal Solutions & Azeotropes
When A-B forces differ from A-A/B-B forces, solutions deviate and form constant-boiling azeotropes.
Azeotropes boil at constant composition and can’t be split by simple distillation.
- Positive: weaker A-B forces, VP higher than predicted (ethanol + water)
- Negative: stronger A-B forces, VP lower than predicted (acetone + chloroform)
- Positive ΔH > 0 & ΔV > 0; negative ΔH < 0 & ΔV < 0
Colligative Properties
These depend only on the NUMBER of solute particles, not their nature.
1 mol glucose and 1 mol urea shift the solution by the same amount.
- Relative lowering of vapour pressure = solute mole fraction
- Built on molality because it is temperature-proof
- They reveal dissociation or association of the solute
Boiling-point Elevation
A non-volatile solute raises the solvent’s boiling point.
K_b = molal elevation (ebullioscopic) constant; m = molality.
- ΔT_b = T_b(solution) − T_b(pure solvent)
- More solute particles → larger ΔT_b
- K_b is a property of the solvent, not the solute
Freezing-point Depression
A solute lowers the freezing point, which is why salt melts ice on roads.
K_f = molal depression (cryoscopic) constant; m = molality.
- ΔT_f = T_f(pure solvent) − T_f(solution)
- Depends only on particle count, not identity
- Used to find molar mass of an unknown solute
Osmotic Pressure
The pressure needed to stop solvent flowing across a semipermeable membrane into the solution.
R = 0.0821 L atm mol⁻1 K⁻1, T in kelvin; equal-π solutions are isotonic.
- Most accurate colligative property, best for macromolecules
- Molar mass: M = w R T / (π V)
- Isotonic = same π; hyper/hypotonic shrink or swell cells
Van’t Hoff Factor
Factor i corrects colligative formulas when a solute dissociates or associates.
Modified: ΔT_b = i K_b m, ΔT_f = i K_f m, π = i C R T.
- Dissociation i > 1: NaCl → 2 (i ≈ 2), BaCl2 → 3, K4[Fe(CN)6] → 5
- Association i < 1: benzoic acid dimerises in benzene (i ≈ 0.5)
- No change i = 1: glucose, urea, sucrose
Degree of Dissociation/Association
From i you can back out how completely a solute splits or pairs up.
n = number of ions (dissociation) or molecules combining (association).
- Higher i → larger colligative effect → lower observed molar mass
- Dissociation raises particle count; association lowers it
- Abnormal molar mass arises whenever i ≠ 1
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Frequently Asked Questions
A colligative property depends only on the number of solute particles in a solution and not on their chemical identity. The four colligative properties are relative lowering of vapour pressure, elevation of boiling point, depression of freezing point, and osmotic pressure.
Relative lowering of vapour pressure equals the solute mole fraction; elevation of boiling point is delta T_b = K_b times molality; depression of freezing point is delta T_f = K_f times molality; and osmotic pressure is pi = CRT. When a solute dissociates or associates, each formula is multiplied by the van’t Hoff factor i.
Molarity is moles of solute per litre of solution, while molality is moles of solute per kilogram of solvent. Molarity changes with temperature because the solution volume expands, but molality is temperature-independent since it uses mass, which is why colligative-property formulas use molality.
The van’t Hoff factor i is the ratio of the observed colligative effect to the value expected for an undissociated solute. It is greater than 1 for solutes that dissociate, such as NaCl giving about 2 and BaCl2 about 3, and less than 1 for solutes that associate, such as benzoic acid dimerising in benzene; for non-electrolytes like glucose and urea i equals 1.
Yes, Solutions is part of the Class 12 NEET chemistry syllabus and is a high-scoring chapter built on direct numerical formulas. Henry’s law relates the partial pressure of a gas above a solution to its mole fraction in the liquid (p = K_H times x), whereas Raoult’s law relates the vapour pressure of a volatile liquid component to its mole fraction (p_A = p_A pure times x_A); Henry’s law is essentially the special case for gas solubility.