1. Properties of Solutions
Chapter 13 BLB 11th

2. Expectations: g ↔ mol (using molar mass) g ↔ mL (using density)
Other conversions: temp., pressure, etc.
Solve for any variable in a formula.
Distinguish between molecular and ionic compounds.
Convert between different concentration units.
Describe the properties of solutions.

3. 13.1 The Solution Process Solution – homogeneous mixture
Solute – present in smaller quantity
Solvent – present in larger quantity
Intermolecular forces are rearranged when a solute and solvent are mixed.

4. Making a Solution Solute molecules separate (endothermic)
Solvent molecules separate (endothermic)
Formation of solute-solvent interactions (exothermic)
ΔHsoln = total energy
ΔHsoln – enthalpy change for the formation of a solution; exothermic – usually favorable; endothermic – usually unfavorable

6. Will a solution form?
Solute-solvent interaction must be stronger or comparable to the separation of solute and solvent particles.
Intermolecular forces play a key role.
Entropy (disorder) is also a factor.
Disorder is favorable. (2nd law of thermodynamics)
Solution formation increases entropy.
Dissolve vs. react (p )

7. Entropy in Solution Formation
Ionic compound
very ordered
As the ionic compound dissolves, it
becomes more disordered.

8. 13.2 Saturated Solutions and Solubility
Saturated solution – solution is in equilibrium with undissolved solute.
Solute + solvent ⇌ solution
Unsaturated – less solute than saturated
Supersaturated – more solute than saturated
dissolution
crystallization

9. A Saturated Solution
A dynamic equilibrium – ions continually exchange between the solid and solution form.

10. 13.3 Factors Affecting Solubility
Like dissolves like, i.e. same polarity.
Polar solutes are soluble in polar solvents.
Nonpolar solutes are soluble in nonpolar solvents.
If two liquids: miscible or immiscible
Examples:
✔ water + alcohol, NaCl + water, hexane + pentane
✘ water + hexane, NaCl + benzene, oil + water

11. Fat- and Water-Soluble Vitamins

12. 13.3 Factors Affecting Solubility
Pressure Effects (for gases in any liquid solvent)
Solubility increases as the partial pressure above the solution increases.
Henry’s Law: Sg = kPg
Sg – solubility of gas
k – Henry’s Law constant; conc./pressure units
Pg – partial pressure of gas above solution

13. The partial pressure of O2 in your lungs varies from 25 to 40 torr
The partial pressure of O2 in your lungs varies from 25 to 40 torr. What molarity of O2 can dissolve in water at each pressure? The Henry’s Law constant for O2 is 6.02 x 10-5 M/torr.

14. 13.3 Factors Affecting Solubility
Temperature Effects
For solids: Solubility ↑ as temperature ↑ - usually.
If ΔHsoln > 0 (endothermic)
If ΔHsoln < 0 (exothermic)
For gases: Solubility ↓ as temperature ↑ - always.
Kinetic energy plays a primary role.
Entropy is also a factor.

15. Ionic compounds

16. Gases
(In liquids)

17. 13.4 Ways of Expressing Concentration
Mass %, volume %, and ppm

18. 13.4 Ways of Expressing Concentration, cont.
Mole fraction, molarity, and molality

19. 44. A solution contains 80.5 g ascorbic acid (C6H8O6) in 210 g water and has a density of 1.22 g/mL at 55°C. Calculate mass %, X, m, and M.

20. 51. Commercial aqueous nitric acid has a density of 1
51. Commercial aqueous nitric acid has a density of 1.42 g/mL and is 16 M. Calculate mass % of HNO3.

21. 43. A sulfuric acid solution containing 571
43. A sulfuric acid solution containing g of H2SO4 per liter of solution has a density of g/cm3. Calculate the mass %, mole fraction, molality, and molarity.

22. Concentration Problems
Practice!
See Figure 13.19, p. 545, for conversion map.
Several examples on pp

23. 13.5 Colligative Properties
The addition of a solute to a pure solvent:
Lowers the vapor pressure
Lowers the freezing point
Raises the boiling point
Causes movement through a semipermeable membrane (osmosis)
Depends on the number of solute particles (moles), not the identity; more particles the greater the effect
Ionic compounds cause an even greater effect.

25. 1. Lowering the vapor pressure
Addition of solute blocks the solvent from evaporation.
More solute, less vapor, lower vapor pressure
Raoult’s Law (for a nonvolatile solute):
PA = XAPA° PA – solvent v. p. over solution
(PA < PA°) PA° – pure solvent v. p.
XA – mole fraction of solvent

26. 1. Lowering the vapor pressure, cont.
When a volatile solute is added, both the solvent and solute contribute to the vapor pressure.
“Expanded” Raoult’s Law:
Ptotal = PA + PB = XAPA° + XBPB°
If a solution obeys Raoult’s Law, it is an ideal solution.
Nonideal solutions have strong intermolecular interactions which lower the vapor pressure of the solution even further.

27. Vapor Pressure Lowering

28. 62a. Calculate the vapor pressure above a solution of 32
62a. Calculate the vapor pressure above a solution of 32.5 g C3H8O3 (glycerin-nonvolatile) in 125 g water at 343 K. The vapor pressure of water at 343 K is torr.

29. 63. A solution is made from equal masses of water and ethanol (C2H5OH)
63. A solution is made from equal masses of water and ethanol (C2H5OH). Calculate the vapor pressure above the solution at 63.5°C. The vapor pressures of water and ethanol are 175 and 400. torr, respectively, at 63.5°C.

30. 2. Boiling point elevation 3. Freezing point depression
Since a solution has a lower vapor pressure:
A higher temperature is needed to boil solution
A lower temperature is needed to freeze solution.
To calculate effect:
b.p. ↑ ΔTb = Kb·m solution − solvent
f.p. ↓ ΔTf = Kf·m solvent − solution
ΔT – difference between boiling or freezing points of the pure solvent and solution
K – boiling or freezing pt. dep. constant (specific to solvent)
m – molality

31. 69a. Calculate the freezing and boiling points of a solution that is 0
69a. Calculate the freezing and boiling points of a solution that is 0.40 m glucose in ethanol. For ethanol: f.p °C, b.p. 78.4°C, Kf = 1.99 °C/m, Kb = 1.22 °C/m

32. 72. Calculate the molar mass of lauryl alcohol when 5
72. Calculate the molar mass of lauryl alcohol when 5.00 g of lauryl alcohol is dissolved in kg benzene (C6H6). The freezing point of the solution is 4.1°C. For benzene: f.p. 5.5°C, Kf = 5.12 °C/m

33. 4. Osmosis
Osmosis – movement of solvent molecules through a semipermeable membrane from a region of lower solute concentration to a region of higher solute concentration
Driving force – to dilute the higher concentration
Continues until:
Equilibrium is reached between two solutions, or
External pressure prevents further movement.

35. Osmosis in red blood cells
Hypertonic solution Hypotonic solution

36. 4. Osmosis, cont. Osmotic pressure
P = M R T P – osmotic pressure (atm)
M – molarity
R – L∙atom/mol∙K
T – temperature (K)
Good technique for measuring molar mass of large molecules like proteins

37. 4. Osmosis, cont. Applications:
Kidney dialysis
Intercellular transport
Reverse osmosis – apply external pressure to reverse the flow of solvent molecules
Water purification – alternative to salt ion exchange
Desalination – purification of salt water

38. 78. A dilute aqueous solution of an organic compound is formed by dissolving 2.35 g in water to form L of solution. The resulting solution has an osmotic pressure of atm at 25°C. Calculate the molar mass of the compound.

39. 13.6 Colloids Colloid or colloidal dispersion
Intermediate between a solution and a suspension
Dispersing medium – analogous to solvent
Dispersing phase – analogous to solute; typically large molecules with high molar masses
Does not settle
Tyndall effect – particles scatter light

40. Tyndall Effect

41. Types of Colloids

42. Surfactants
Change the surface properties so that two things that would not normally mix do
Emulsifying agent
Soap
Detergent

43. Hydrophobic – water-fearing (nonpolar) Hydrophilic – water-loving (polar)

44. Action of soap on oil