Solutions
Solutions is the domain that tests how concentration, vapour pressure and colligative properties connect. Kerala Plus Two questions usually combine a numerical part with a conceptual part.
Concentration and solubility
Molarity M = moles of solute per litre of solution and changes with temperature; molality m = moles of solute per kilogram of solvent and does not. Mole fraction is dimensionless, ppm is parts per million by mass, and normality is equivalents per litre. Henry's law p = K_H x explains why gas solubility falls as temperature rises.
Raoult's law and deviations
For a non-volatile solute, relative lowering of vapour pressure equals the mole fraction of solute. Ideal solutions obey Raoult's law; positive deviation occurs when A–B forces are weaker than A–A and B–B forces, while negative deviation occurs when A–B forces are stronger, as in chloroform–acetone. Azeotropes cannot be separated by ordinary fractional distillation.
Colligative properties
ΔTb = i Kb m, ΔTf = i Kf m and π = iCRT. The van't Hoff factor i is the observed colligative effect divided by the calculated effect: i > 1 for dissociation, i < 1 for association such as benzoic acid dimerising in benzene.
Worked example
18.0 g glucose (M = 180 g mol−1) in 500 g water gives n = 0.100 mol and m = 0.100/0.500 = 0.200 mol kg−1. ΔTf = 1.86 × 0.200 = 0.372 K, so the solution freezes at −0.372 °C. For 5.85 g NaCl (M = 58.5 g mol−1) in the same water, n = 0.100 mol and m = 0.200 mol kg−1; with i ≈ 2, ΔTf ≈ 0.744 K. Convert mass to moles, then to molality, and apply i only to electrolytes.
Mark-losing traps
Using solution mass instead of solvent mass, forgetting i, treating molarity as molality at all concentrations, and mixing units in π = iCRT are common errors. In π, C must be mol L−1 and R = 0.08314 L bar K−1 mol−1 for pressure in bar.