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📑 In This Chapter Guide (Table of Contents)
1. Henry’s Law & Raoult’s Law
Henry’s Law: At constant temperature, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas present above the liquid surface: p = K_H · x (where K_H is Henry's law constant; higher K_H means lower gas solubility at a given pressure).
Raoult’s Law for Liquid-Liquid Volatile Solutions: For any volatile component in solution, its partial vapor pressure is directly proportional to its mole fraction: p₁ = p₁° · x₁ and p₂ = p₂° · x₂.
Ideal vs. Non-Ideal Solutions:
- Ideal Solutions: Obey Raoult’s law across all concentrations.
ΔH_mix = 0,ΔV_mix = 0. A-B intermolecular attractions equal A-A and B-B. Example: n-hexane + n-heptane, benzene + toluene. - Non-Ideal with Positive Deviation: A-B attractions are weaker than pure components.
ΔH_mix > 0,ΔV_mix > 0. Forms minimum boiling azeotropes. Example: Ethanol + Acetone. - Non-Ideal with Negative Deviation: A-B attractions are stronger (hydrogen bonding).
ΔH_mix < 0,ΔV_mix < 0. Forms maximum boiling azeotropes. Example: Chloroform + Acetone.
2. The Four Colligative Properties
Colligative properties depend solely on the total number of solute particles present in solution, independent of their chemical nature:
| Colligative Property | Mathematical Equation | Calculation of Solute Molar Mass (M₂) |
|---|---|---|
| 1. Relative Lowering of Vapor Pressure (RLVP) | (P₁° - P₁) / P₁° = i · x₂ |
M₂ = [ (W₂ · M₁) / W₁ ] × [ P₁° / (P₁° - P₁) ] |
| 2. Elevation of Boiling Point | ΔT_b = T_b - T_b° = i · K_b · m |
M₂ = (1000 · K_b · W₂) / (ΔT_b · W₁) |
| 3. Depression of Freezing Point | ΔT_f = T_f° - T_f = i · K_f · m |
M₂ = (1000 · K_f · W₂) / (ΔT_f · W₁) |
| 4. Osmotic Pressure (π) | π = i · C · R · T = i · (n₂ / V) · R · T |
M₂ = (W₂ · R · T) / (π · V) (Ideal method for determining molar mass of proteins and polymers) |
3. Van’t Hoff Factor (i) & Abnormal Molar Mass
The Van’t Hoff factor (i) accounts for association or dissociation of ionic solute particles in solution:
i = (Normal / Theoretical Molar Mass) / (Observed / Experimental Molar Mass)
i = Total number of moles of particles after association/dissociation / Initial moles
- For Dissociation (Electrolytes like NaCl, K₂SO₄):
i = 1 + (n - 1)α(where n = number of ions produced, α = degree of dissociation). For complete dissociation of NaCl (n=2),i = 2; for K₂SO₄ (n=3),i = 3. - For Association (Carboxylic acids dimerizing in benzene):
i = 1 - (1 - 1/n)α. For dimerization of acetic acid in benzene (n=2),i < 1(~0.5).
💡 Frequently Asked Questions (FAQ)
❓ Why is osmotic pressure measurement preferred over other colligative properties for polymers and biomolecules?
Osmotic pressure is measured at room temperature (preventing thermal denaturation of sensitive proteins) and uses molarity (which is convenient to measure). Furthermore, biomacromolecules have enormous molar masses, resulting in negligibly small boiling point elevation or freezing depression, whereas osmotic pressure generates measurable magnitude even in dilute solutions.
❓ What are isotonic solutions? What happens when RBC is placed in hypertonic saline?
Two solutions having identical osmotic pressure at the same temperature are called isotonic solutions; no net osmosis occurs between them. When a Red Blood Cell (RBC) is placed in hypertonic saline (>0.9% NaCl), water flows out of the cell via exosmosis, causing the RBC to shrink and crenate.
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