1. 1 Colligative Properties of Solutions

2. 2 Colligative Properties Colligative properties are physical properties of solutions that change when adding a solute to a solvent. These properties are changed by the number of particles dissolved, NOT the type of particle.  Vapor Pressure  Boiling Point  Freezing Point  Osmotic Pressure

3. 3 Electrolytes and Nonelectrolytes Electrolyte – a solution that contains ions. It will conduct electricity.  Created when water (solvent) dissociates an ionic compound or ionizes a polar molecule.  NaCl dissociates into Na + and Cl -  1 mole of NaCl forms 2 moles of ions.  AlCl 3 dissociates into Al 3+ and 3Cl -  1 mole of AlCl 3 forms 4 moles of ions.

4. 4  The more ions (particles) in solution, the stronger the electrolyte, the greater change in the colligative properties. Nonelectrolyte – a solution that does not contain ions. Does not conduct electricity.  If I add 1 mole of sugar to a solution, there will be 1 mole of particles.  Nonelectrolytes do affect colligative properties, but not as much as electrolytes.

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7. 7 Vapor Pressure Vapor pressure – the pressure exerted when liquid particles have escaped the liquid’s surface and enter the gaseous state. In a closed container, the solvent particles reach a dynamic equilibrium where the number of liquid particles evaporating is equal to the number of gaseous particles condensing. When a solute is added to a solvent, the intermolecular forces prevent the solvent particles to enter the gaseous state. So the vapor pressure decreases.

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9. 9 Number of solute particles ↑ Solution Vapor Pressure ↓ at constant temperature Only if the solute is nonvolatile – does not evaporate easily. (ionic compounds, polar molecules with high melting points).

10. 10 Boiling Point Elevation When adding a nonvolatile solute to a solvent the boiling point increases. The vapor pressure is decreased, so it takes a higher temperature to give them enough energy to escape into the gaseous state. The amount of boiling point elevation depends on the amount of solute dissolved, not the type. 1 mole of a nonvolatile solute will raise the boiling point of 1 kg of water by 0.512°C.

11. 11 The amount of change depends on the solvent and the concentration. Molal Boiling point constant, K bp  table 2 on page 448 of the textbook  Water K bp = 0.512 ° C/m (m = molality)  Benzene K bp = 2.53 ° C/m Boiling point elevation = ∆ T bp  ∆ T bp = K bp m m = molality  T soln = T bp + K bp m When you add salt to water, it boils at a higher temperature!

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13. 13 What is the boiling point when 11.4 g of ammonia (NH 3 ) is dissolved in 200.0 mL of water?

14. 14 What is the boiling point of a solution containing 85.42 g of copper(II) sulfate in 100.0  mL of water?

15. 15 Freezing Point Depression When adding a nonvolatile solute to a solvent the freezing point decreases. The vapor pressure is decreased, so the solvent freezes at a lower temperature. The amount of freezing point depression depends on the amount of solute dissolved, not the type. 1 mole of a nonvolatile solute will lower the freezing point of 1 kg of water by 1.86°C.

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17. 17 The amount of change depends on the solvent and the concentration of the solute. Molal Freezing point constant, K fp  table 2 on page 448 of textbook  Water K fp = -1.86 °C/m (m = molality)  Benzene K fp = -5.12 °C/m Freezing point depression = ∆ T fp  ∆ T fp = K fp m m = molality  T soln = T fp + K fp m When you add salt to ice, the freezing point decreases, and the ice melts at a temperature less than 0°C.

18. 18 Molality = moles solute/kg solvent You can measure the molar mass of a solute based on the change in the boiling or freezing points. You measure the change in the freezing or boiling point, and calculate the molality.  ∆ T fp = K fp m T soln = T fp + K fp m  ∆ T bp = K bp m T soln = T bp + K bp m  m = moles solute/kg of solvent.  You measured the mass of the solute and the mass of the solvent.

19. 19  Molar Mass (g/mole) =

20. 20 What is the freezing point of a mixture when 85.0 g of sodium chloride is dissolved in 392 mL of water?

21. 21 What is the molality of a sugar solution that freezes at -5.0°C?

22. 22 When 6.00 g of an unknown is dissolved in 125 mL of water, the solution’s freezing point is lowered by -1.38°C. What is the molar mass of the unknown?

23. 23 How many grams of antifreeze C 2 H 4 (OH) 2 would be required per 500.0mL of water to prevent the water from freezing down to -20.0°C?

24. 24 Osmotic Pressure Osmosis – movement of a solvent (water) from an area of higher solvent concentration (low solute) to an area of lower solvent (high solute) concentration through a semi- permeable membrane. A special form of diffusion. Semi-permeable membrane – A barrier that allows water to pass through, but blocks many solute particles such as sugar and salt.

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28. 28 As the water flows from the pure solvent side to dilute the solution side, the pressure builds up as a result of the added volume. Osmotic pressure – external pressure that must be applied to stop osmosis. This pressure exists on the solution side. Osmotic pressure depends on the concentration of the solute particles, not the type. Used to preserve fruit (add sugar) and meat (add salt) because bacteria become dehydrated and die.

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30. 30 Reason you cannot drink salt water. The water in your body rushes to dilute the salt water and you dehydrate and die. Osmotic pressure =  = MRT  M = concentration in moles/liter, molarity  R = ideal gas constant  T = temperature, Kelvin.