17: Electrolytes & Colligative Properties of Solutions

1. Electrolyte

A substance that ionizes when dissolved in a polar solvent such as water

2. Solvation

The stabilization of ionic species by a shell of surrounding solvent molecules, through interactions between solvent and solute molecules

3. Solvated ion

An ion in solution that has become stable because it is surrounded by a shell of solvent molecules

4. Hydration sphere

The shell of water molecules around an ion when the solvent is water

5. Classifying electrolytes by ionization (with examples)

Strong electrolytes: ionize 100% (or almost), main species are solvated ions and solvent, almost all are ionic compounds, single arrow (e.g., KCl, (NH4)2SO4, NaCl, HCl, NaOH). Weak electrolytes: partially ionize (1–10%), main species are un-ionized solute, reversible double arrow with an equilibrium between forward and reverse reactions; typically carboxylic acids and amines (e.g., CH3COOH, CH3NH2). Non-electrolytes: don't ionize at all, usually molecular compounds (e.g., sugar C6H12O6, ethanol C2H5OH)

6. Classifying electrolytes by conductivity

Ions in solution let electricity pass, and current strength depends on electrolyte strength: strong electrolytes conduct well (H2O + NaCl conducts because of Na+ and Cl− charges), weak electrolytes conduct weakly (acetic acid), and non-electrolytes don't conduct (pure water, ethanol, sugar)

7. Colligative property

A property that depends on the number of solute particles in solution, not on their nature (1 mol glucose = 1 mol Na+ = 1 mol SO4 2− in effect); concentration expressed as mole fraction or molality

8. Vapor pressure of a liquid or solid at temperature T

The equilibrium pressure the vapor reaches when a liquid (or solid) is placed in an empty, closed vessel at temperature T; increases with temperature because more molecules have enough energy to escape

9. Vapor pressure as an equilibrium of two processes

In a closed vessel, evaporation (surface molecules go into the vapor) and condensation (vapor molecules return to the liquid) happen at the same time; when their rates become equal, the pressure of the vapor is the vapor pressure. The same applies to a solid

10. Raoult's Law

The vapor pressure of an ideal solution: P = X solvent × P°, where P° is the vapor pressure of the pure solvent and X solvent = n solvent / (n solute + n solvent). Adding solute lowers X solvent, and since X < 1, P < P° always

11. Vapor pressure lowering equation (ΔP)

ΔP = P − P° = −X solute × P° (using X solvent = 1 − X solute); it is negative, which makes sense because adding solute always lowers the vapor pressure below that of the pure solvent

12. Van't Hoff factor (i)

Corrects for the fact that the moles of solute put into solution don't always equal the moles of particles that determine a colligative effect: final moles of particles = i × initial moles of solute (e.g., NaCl i = 2, Na3PO4 i = 4)

13. How the actual Van't Hoff factor is calculated

i = actual number of moles of particles / expected number of moles of particles without dissociation (the nominal solute concentration); in practice it is determined experimentally (e.g., from osmotic pressure or freezing point data)

14. Why i can differ from the theoretical value

Association of particles in solution: ion pairs form from Coulombic attraction between oppositely charged ions, and carboxylic acids associate through intermolecular H-bonds. Association increases with concentration, so measured i decreases as concentration increases

15. Interpreting i = 1, i > 1, i < 1

i = 1: no association or dissociation (typical for molecular compounds). i > 1: dissociation occurs; if lower than the ideal value, some association is also happening (ion pairs or H-bonds). i < 1: only association, or both with more association than dissociation

16. Boiling point

The temperature at which a substance's vapor pressure equals the external pressure; molecules have enough energy to overcome the IMFs holding them in the liquid

17. Boiling point vs. external pressure

The higher the external pressure, the higher the boiling point

18. Normal boiling point

The boiling point when the external pressure is 1 atm

19. Why a solution boils higher than the pure solvent

The solute lowers the solvent's vapor pressure, so more energy (a higher temperature) is needed to raise the vapor pressure up to the external pressure

20. Boiling point elevation equation

ΔTb = Tb,solution − Tb,solvent = i Kb m, where i = Van't Hoff factor, Kb = ebullioscopic constant (positive, depends on the solvent; 0.512 °C/m for water), and m = molality of the solute

21. Freezing point depression equation

ΔTf = Tf,solution − Tf,solvent = −i Kf m (the solution freezes lower), where i = Van't Hoff factor, Kf = cryoscopic constant (a property of the solvent; 1.86 °C/m for water), and m = molality. Example: 0.500 m glucose (i = 1) → ΔTf = 0.930 °C, so it freezes at −0.930 °C and boils at 100.256 °C

22. Effect of a solute on the solvent's phase diagram

The liquid-vapor line shifts down (lower vapor pressure), so the boiling point rises and the freezing point falls; the liquid region becomes wider on both ends

23. Semipermeable membrane

A membrane that lets certain molecules or ions pass but not others (e.g., water passes, larger solute molecules and hydrated ions don't)

24. Osmotic pressure

Solvent (water) flows through a semipermeable membrane from pure solvent into the solution side, raising the solution's level until equilibrium; the osmotic pressure (Π) is the pressure that would have to be applied to the solution to stop this flow

25. Osmotic pressure equation

Π = i n solute R T / V solution = i R T C solute, where Π = osmotic pressure, i = Van't Hoff factor, n = moles of solute, V = volume of solution, C = molarity, R = gas constant, T = temperature in K. Example: 2.460 atm from 1.000 g in 100.0 mL at 300 K → n = 0.00999 mol → molar mass = 100. g/mol

26. Hypertonic, hypotonic, isotonic solutions

Hypertonic: higher solute concentration than the cell, so water leaves and the cell shrinks. Hypotonic: lower solute concentration (e.g., pure water), so water enters and the cell swells or bursts. Isotonic: same concentration, no net water flow. This is why freshwater fish can't live in the ocean and vice versa