Colligative Properties of Solutions

1. Colligative Properties of Solutions …are properties or behaviors of solutions that depend only on the number of particles dissolved in solution. The chemical identity of the substance is not important since different substances can produce the same effect as long as the number of solute particles is the same.

2. Examples of Colligative Properties • Freezing point depression • Boiling point elevation • Osmosis

3. Definitions: Freezing point depression – this means that the temperature at which a solution freezes is lower than the temperature at which the pure solvent freezes. The amount of lowering depends on the solvent and concentration of particles.

4. Examples relating to FP Lowering: • Salt + H2O freezes below 0° C. • If the outside temperature is -10 °C a patch of ice would need to warm up to 0 ° C before it could melt. However, a mixture of water and salt may not be able to freeze until the temperature

5. is -15° C. So, the ice melts at -10° C when salt is thrown on it! There is enough heat in the environment even at -10° C to provide the heat of fusion (melting). Example (b): Ethylene glycol, also called anti-freeze, is added to the water in the radiator of an automobile to lower the freezing point to below -25° F.

6. Boiling Point Elevation – This means that the boiling point of a solution is higher than the boiling point of the pure liquid. Example: “anti-freeze” is added to the water in a car radiator to both lower the temperature at which the mixture freezes AND

7. raise the temperature at which the cooling water would boil. Another example: in making candy or jelly, a mixture of sugar and water is boiled until the temperature is well above 100° C. This corresponds to having boiled off enough water to cause the mixture to harden when cool.

8. In driving off water by boiling, the concentration of the dissolved sugar increases. In turn, this raises the boiling point. The correct concentration to make jelly or candy can be determined by measuring the boiling point of the mixture.

9. Osmosis is the most clinically and biologically important colligative property we will examine. Osmosis is the process in which solvent flows from a dilute solution into a more concentrated solution across a semi-permeable membrane.

10. Small solvent molecules like H2O can pass through a semi-permeable membrane, but larger solute particles can not pass through the barrier. The membrane is riddled with very small holes through which only solvent molecules can fit.

11. 1.0 M glucose H2O An experiment with osmosis uses a container separated into two compartments by a semi-permeable membrane (red line). Initially the two compartments are filled to the same level: Time passes

12. glucose H2O Later we find the liquid levels are unequal: H2O flows across the membrane to dilute the glucose solution: osmosis.

13. The pressure that would need to be applied to the surface of the glucose solution to keep solvent from flowing across the membrane is called the osmotic pressure of the solution. The osmotic pressure increases as the concentration of dissolved particles increases.

14. NaCl glucose Another experiment: 1.0 M NaCl 1.0 M glucose Time passes H2O flows into NaCl solution. Why?

15. Notice that 1.0 M NaCl and 1.0 M glucose (C6H12O6) do not contain the same number of dissolved particles: 1.0 M NaCl (ionic compound) = 1.0 M Na+ and 1.0 M Cl- (2 moles particles per Liter solution). C6H12O6 (molecule), no ionization.

16. A new concentration unit is useful in understanding osmosis. The osmolarity of a solution = (molarity of the dissolved compound) x (the number of particles produced per mole of compound). The abbreviation for osmolarity is osmol.

17. For 1.0 M NaCl we have: (1.0 mole NaCl/L) x (2 mole particles/mole NaCl) = 2 osmol. For 1.0 M C6H12O6 , (1.0 mole glucose/L) x (1 mole particles/mole glucose) = 1 osmol.

18. Reminder: ionic compounds dissociate in water to give the number of units indicated by the formula. Most molecules (covalent) do not dissociate when they dissolve in water. Glucose, starch, alcohol, urea, etc. do not dissociate.

19. To summarize: *Solvent flows across a semi- permeable membrane from a more dilute solution into a more concentrated solution. *Solvent flows across a semi-perm membrane from a solution of lower osmolarity to a solution of higher osmolarity.

20. *The solution of higher osmolarity has a higher osmotic pressure, so *Solvent flows from a solution of lower osmotic pressure across a membrane into a solution of higher osmotic pressure.

21. 0.75 M CaCl2 1.0 M KBr Sample problem: what happens over time in this case? • No water flows • b) Water flows from left to right () • c) Water flows from right to left ()

22. Calculate osmolarity values: 0.75 M CaCl2 x 3 moles particles = 2.25 osmol. 1.0 M KBr x 2 moles particles = 2.0 osmol. Water flows from KBr solution into CaCl2 solution ().

23. Clinical relevance Cell membranes are semi permeable. The fluid inside a typical red blood cell has about the same osmotic pressure as 0.89 (w/v) % NaCl, called physiological saline. 0.89 g NaCl/100 mL solution = 0.015 mole NaCl/100 mL = 0.15 M = 0.30 osmol.

24. 0.60 osmol 0.30 osmol What happens if the blood cell is placed in a 0.6 osmolar solution of NaCl? External solution has higher osmotic pressure: hypertonic

25. 0.60 osmol 0.30 osmol So: water flows out of cell into external solution…cell shrinks (crenation). H2O After Before

26. 0.10 osmol C6H12O6 0.30 osmol NaCl What happens if blood cells are placed in 0.10 osmol glucose? External solution has lower osmotic pressure than cell: hypotonic

27. 0.10 osmol C6H12O6 0.30 osmol NaCl So: water flows into cell from external solution…cell swells and may burst (lysis). H2O After before

28. When the external solution has the same osmotic pressure as inside the cell, the flow of water in and out the cell is exactly balanced: no change in cell size occurs. When the internal and external solutions have the same osmotic pressure they are isotonic.