1. Unit 12: Properties of Solutions
CHM 1046: General Chemistry and Qualitative Analysis
Unit 12: Properties of Solutions
Dr. Jorge L. Alonso
Miami-Dade College – Kendall Campus
Miami, FL
Textbook Reference:
Chapter # 14
Module # 2

2. Solutions
Solutions (soln) are homogeneous mixtures of two or more pure substances.
The solvent (solv) is present in greatest abundance.
All other substances are solutes (solu).
Volumetric flask
{PrepASolu}

3. Solutions
Solutions are homogeneous mixtures of two or more pure substances.
In a solution, the solute (< 50%) is dispersed uniformly throughout the solvent (> 50%).
Homogeneous Heterogeneous

4. Student, Beware!: solution vs. reaction
H2O
H2O
Cu(NO3)2 (s) Cu(NO3)2 (aq)
Mg (s) + 2 HCl (aq) 
H2 (g) + MgCl2 (aq)
Just because a substance disappears when it comes in contact with a solvent, it doesn’t mean the substance dissolved.
Dissolution is a physical change—you can get back the original solute by evaporating the solvent.
If you can’t, the substance didn’t dissolve, it reacted.

5. How Does a Solution of Salts in Water Form?
The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.
Example:
{NaCl + H2O*}
NaCl (s) Na+ (aq) Cl- (aq)
+ H20
dissociation

6. Solubility in water
Covalent Compounds: many are insoluble gases, polar covalent are mostly soluble.
Elements: mostly insoluble solids, liquids & gases.
Ionic Compounds: many are soluble.
SOLUBILITY RULES: for Ionic Compounds (Salts)
All salts of alkali metals (IA) are soluble.
All NH4+ salts are soluble.
All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
All Cl-, Br-, and I- are soluble except for Ag+, Pb2+, and Hg22+ salts.
All SO42- are soluble except for Pb2+, Sr2+, and Ba2+.
All O2- are insoluble except for IA metals Ca2+, Sr2+, and Ba2+ salts.
{Soluble metal oxides form hydroxides: CaO Ca OH-}
7. All OH- are insoluble except for IA metals, NH4+ & slightly soluble Ca 2+ Ba2+ & Sr2+
6. All salts containing the anions: CO32-, PO43-, AsO43-, S2- and SO32- are insoluble except fro IA metals and NH4+ salts.
7. For salts containing the anions not mentioned above (e.g., CrO42-, Cr2O72-, P3-, C2O42- etc.) assume that they are insoluble except for IA metals and NH4+ salts, unless, otherwise informed.
H2O

7. Energetics of Solutions Heats (Enthalpy) of Solution (Hsoln)
Exothermic solution process
Hot & Cold Packs
H2O
CaCl2 (s) Ca 2+(aq) + 2Cl- (aq)
+ Heat
Hsoln = kJ/mol
CaCl2
+ H2O
NH4NO3
+H2O
Endothermic solution process
H2O
Heat +
NH4NO3 (s) NH4+ (aq) + NO3-(aq)
Hsoln = kJ/mol

8. How Does a Solution Form? (3 events)
(3) Formation of Ion-dipole interactions
(2) Separation of solvent molecules
(1) Separation of solute molecules
H2O
NaCl
Crystal Lattice
H Lattice Energy
+788 kJ/mol Endothermic
Hvap
+41 kJ/mol Endothermic
Exothermic
H- bonding
Energy of Solution Hsoln = + or -
If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

9. Energy Changes in Solution
The enthalpy change of the overall process depends on H for each of these steps.
{EnerSoln} - +

10. Why do Endothermic Reactions occur Spontaneously?
Exothermic solution process
Endothermic solution process
H2O
CaCl2 (s) Ca 2+ (aq) + 2Cl- (aq)
Hsoln = kJ/mol
H2O
NH4NO3 (s) NH4+ (aq) NO3-(aq)
Hsoln = kJ/mol

11. Concentration of Solutions (General)
Unsaturated
Less than the maximum amount of solute for that temperature is dissolved in the solvent.
{UnsatSoln}
Saturated
Solvent holds as much solute as is possible at that temperature.
Dissolved solute is in dynamic equilibrium with solid solute particles.
Solubility of NaCl in H2O = g/100 mL (25 °C)
NaC2H4O2 = 76 g/100 ml (0°C)

12. Concentration of Solutions (General)
{*SuperSat1}
{*SuperSat2}
Supersaturated
Solvent holds more solute than is normally possible at that temperature.
These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

13. Factors Affecting Solubility A. Solids and Liquids:
Polarity
Temperature
B. Gases :
Molar Mass
Pressure – Henry’s Law
Temperature

14. Factors Affecting Solubility of Solids and Liquids: (1) Polarity
Chemists use the axiom “like dissolves like”:
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.
hydrophobic water-hating, nonpolar,
hydrophilic water-loving, polar,
(hydrocarbon chain)
[Know molecular shapes and polar molecules]
{Water + Oil}

15. Factors Affecting Solubility of Solids and Liquids: (1) Polarity
Chemists use the axiom “like dissolves like”:
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.
Polar
Non-Polar
Polar heads
Non-Polar tails
[Know molecular shapes and polar molecules]

16. Factors Affecting Solubility of Solids and Liquids: (1) Polarity
Glucose (which has hydrogen bonding) is very soluble in water, while…..
…..cyclohexane (which only has dispersion forces) is not.

17. Factors Affecting Solubility of Solids and Liquids: (1) Polarity
Vitamin A is soluble in nonpolar compounds (like fats).
Vitamin C is soluble in water.
{*Iodine + Water then + CCl4 }

18. Factors Affecting Solubility of Solids and Liquids: (2) Temperature
Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

19. Factors Affecting Solubility of Gases in Solution: (1) Molar Mass
In general, the solubility of gases in water increases with increasing molar mass.
Which gas dissolves best in water?
g-MM
28 g/n
32 g/n
40 g/n
84 g/n
Why?
Larger molecules have stronger London dispersion forces.

20. Factors Affecting Solubility of Gases in Solution: (2) Pressure
The solubility of liquids and solids does not change appreciably with pressure.
The solubility of a gas (Sg) in a liquid is directly proportional to its pressure (Pg).
{Henrys Law}
Sg Pg

21. Factors Affecting Solubility of Gases in Solution: (2) Pressure
Henry’s Law:
{movie}
Sg Pg
Sg = kPg
where
Sg is the solubility of the gas;
Pg is the partial pressure of the gas above the liquid.
k is the Henry’s law constant for that gas in that solvent; Pg Sg
@ 250C, PN2= 0.75 atm, SN2= M. What is SN2 when PN2 = 1.50 atm?
Ans.= M

22. Factors Affecting Solubility of Gases in Solution: (3) Temperature
The opposite is true of gases:
Gases in carbonated soft drinks are less soluble when warmed.
Warm lakes have less O2 dissolved in them than cool lakes.

23. Colligative Properties
How does the presence of a solute change the properties of a solvent?
Vapor pressure?
Boiling point?
Freezing (melting) point?
Osmotic pressure?
Solution: Solvent + Solute
Pure Solvent

24. Colligative Properties
Properties that depend only on the CONCENTRATION (number of solute particles present), not on the identity of the solute particles.
Vapor pressure lowering
2. Boiling point elevation
3. Freezing point
depression
4. Osmotic pressure
PA = XA  PA
Tb = Kb  m
Tf = Kf  m
 = MRT
Pure solvent
Solvent + Solute

25. Ways of Expressing Concentrations of Solutions
(1) Mass Percentage (%)
(2) Parts per Million (ppm) & Parts per Billion (ppb)
(3) Mole Fraction (X)
(4) Molarity (M)
(5) Molality (m)

26. Mass Percentage (%) (parts per Hundred)
mass of solu A in solution
total mass of solution
Mass % of A =
 100 or 102
hundred
50. g of NaCl in 100.mL of H2O
Parts per Million (ppm)
mass of solu A in solution
total mass of solution
ppm =
 106
million
g of NaCl in
100.mL of H2O
Parts per Billion (ppb)
mass of solu A in solution
total mass of solution
ppb =
 109
billion
0.000,000,5 g of NaCl in 100.mL of H2O

27. total moles in solution (A+B+∙∙∙)
Mole Fraction (X)
moles of substance A
total moles in solution (A+B+∙∙∙)
XA =
You can calculate the mole fraction of either the solvent or of the solute — make sure you find the one you need for your calculations!
What are the mole fractions of methanol and water in a solution containing 128g CH3OH (MW=32.0) and 202 mL of H2O (MW=18.0)? Which is the solute and the solvent?
Mole fraction of Solvent:
Mole fraction of Solute:

28. Molarity (M) mol of solute M = L of solution
Because volume is temperature dependent, molarity can change with temperature.
{The volume of most substances expand with increasing temperature}
What is the molarity of a solution containing 128g CH3OH (MW=32.0,  C) and 202 mL of H2O (MW=18.0,

29. Molality (m) moles of solute m = kg of solvent
Because both moles and mass do not change with temperature, molality (unlike molarity) is not temperature dependent.
What is the molality of a solution containing 128g CH3OH (MW=32.0,  C) and 202 mL of H2O (MW=18.0,

30. Changing Molarity to Molality
If we know the density of the solution, we can calculate the molality from the molarity, and vice versa.
Problem: What is the molality (m) of a 2.0M NaCl solution at 250C if its density is 1.2 g/mL?
mol of solute
kg of solvent
m =
mol of solute
L of solution
M =
= 1.8 m NaCl

31. Colligative Properties
Properties that depend only on the number of solute particles present, not on the identity of the solute particles.
Vapor pressure lowering
2. Boiling point elevation
3. Freezing point
depression
4. Osmotic pressure
PA = XA  PA
Tb = Kb  m
Tf = Kf  m
 = MRT
Pure solvent
Solvent + Solute

32. 1. Vapor Pressure (Pvap ) Lowering
Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.
{PureSolv}
Therefore, the vapor pressure of a solution (PA ) is lower than that of the pure solvent (PA).
PA = XA  PA
{*Solution}

33. Raoult’s Law: PA = XA P A where XA is the mole fraction of solvent A
{VapPress.WaterVsEthGlycol}
PA = XA P A
where
XA is the mole fraction of solvent A
P A is the normal vapor pressure of the pure solvent A at that temperature
P A is the new vapor pressure of solution at that temperature
Consider: XA= XA= XA= 0.5

34. 2. Freezing Point Depression & 3. Boiling Point Elevation
f.pt.
b.pt.
f.pt. Solution
f.pt. Pure solvent
Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

35. Freezing Point Depression
{f.pt. Equil}
{f.pt. Lower}
The change in freezing point is proportional to the molality of the solution :
Tf = Kf  m
Here Kf is the molal freezing point depression constant of the solvent.
Tf is subtracted from the normal freezing point of the solvent.

36. Boiling Point Elevation
The change in boiling point is proportional to the molality of the solution:
Tb = Kb  m
where Kb is the molal boiling point elevation constant, a property of the solvent.
Tb is added to the normal boiling point of the solvent.

37. Freezing Point Depression & Boiling Point Elevation
Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved.
Tf = Kf  m
Tb = Kb  m

38. Diffusion
The movement of particles from an area of high concentration to lower concentration until a homogeneous solution has been formed.
Osmosis
The movement of water from an area of high concentration to lower concentration (diffusion) across a semipermeable membrane (SPM), until a homogeneous solution has been formed (equilibrium). The SPM allows smaller particles (water) to pass through, but blocks other larger particles.
Diffusion of water across a SPM.

39. Osmosis
In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the area of lower solvent concentration (higher solute concentration).
Solvent
Solution
{*OsmoPress}

40. Solution side (Low Conc. H2O)
{*OsmoSemiPerMemb}
Solvent side
(High Conc. H2O)
Solution side (Low Conc. H2O)
SPM

41. Osmosis in Cells Low Water High Water
If the water concentration outside the cell is more than that inside the cell, the solution is hypotonic (low in solute).
Water will flow into the cell, and hemolysis results.
High Water
H2O
H2O
(low solute)
H2O
H2O
H2O
H2O
H2O
H2O
Low Water
H2O
H2O
(high solute)
H2O
H2O
H2O
H2O
H2O
H2O
H2O
{EggOsmoPres}

42. Osmosis in Blood Cells High Water
If the water concentration outside the cell is lower than that inside the cell, the solution is hypertonic (high in solute).
Water will flow out of the cell, and crenation results.
Isotonic: equal concentrations.
H2O
(high solute)
(salt water)
H2O
H2O
H2O
H2O
H2O
High Water
H2O
H2O
(salt
water)
(low solute)
H2O
H2O
H2O
H2O
H2O
H2O
(salt water)
(salt water)

43.  V = n RT  = ( )RT 3. Osmotic Pressure  = MRT n V
The pressure required to stop osmosis, known as osmotic pressure, , is
 V = n RT n V
 = ( )RT
 = MRT
where M is the molarity of the solution

44. Reverse Osmosis Reverse Osmosis
What exactly is Reverse Osmosis?
Reverse Osmosis is the method of purifying liquid by pushing it through a semi-permeable membrane.Osmosis is the process by which water and nutrients are supplied to living cells. The natural flow of water is from a dilute to a concentrated solution across a cell wall. Cell walls are natural, selective semi-permeable membranes, which allow certain materials to pass through but reject others.In a reverse osmosis system, pressure is utilised to reverse the natural flow through a synthetic membrane so that pure water molecules pass through and impurities are flushed away.
                                          
What exactly is Reverse Osmosis?
Reverse Osmosis is the method of purifying liquid by pushing it through a semi-permeable membrane.Osmosis is the process by which water and nutrients are supplied to living cells. The natural flow of water is from a dilute to a concentrated solution across a cell wall. Cell walls are natural, selective semi-permeable membranes, which allow certain materials to pass through but reject others.In a reverse osmosis system, pressure is utilised to reverse the natural flow through a synthetic membrane so that pure water molecules pass through and impurities are flushed away.
                                          
Reverse Osmosis
Reverse Osmosis
{RevOsmo}
{RevOsmo}
The method of purifying liquid by pushing it through a semi-permeable membrane. Pressure is utilized to reverse the natural osmotic flow through a synthetic membrane so that pure water molecules pass through and impurities are flushed away.
Contaminated water
Contaminated water
Membrane
Membrane
Pure
H2O
H2O
H2O
H2O
H2O
H2O
H2O

45. Determining Molar Mass from Colligative Properties
We can use the effects of a colligative property such as freezing point depression to determine the molar mass (MM) of a compound.
A solution of an unknown nonvolatile electrolyte was prepared by dissolving 0.250g of the substance in 40.0 g of CCl4. The b. pt. of the resultant solution was C higher than that of the pure solvent. What is MM of solute? The Kb for CCl4 is C/m.
Tb = Kb  m

46. Determining Molar Mass from Colligative Properties
We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound.
The osmotic pressure of a protein solution was 250C to be 1.54 torr. The solution contained 3.50 mg of protein dissolved in water to make 5.00 mL of solution. Determine the molar mass (MM) of the protein.
 = MRT

47. Colligative Properties & Molar Mass
Vapor pressure lowering:
Boiling point elevation:
Freezing point depression:
Osmotic pressure:
PA = XA  PA
=  PA
Tb = Kb  m
= Kb 
Tf = Kf  m
= Kf 
 = MRT
= RT

48. Colligative Properties of Electrolytes
Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.
H2O
Non-electrolyte:
C6H12O6 (s) C6H12O6 (aq) 1η
1 η molecules
H2O
CaCl2 (s) Ca 2+ (aq) + 2Cl- (aq)
Electrolyte: 1η
3 η ions
You would expect a 1 M solution of CaCl2 to show three times the change in freezing point that a 1 M solution of C6H12O6, however …………….

49. van’t Hoff Factor
One mole of NaCl in water does not really give rise to two moles of ions.
Some Na+ and Cl− reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.
Reassociation is more likely at higher concentration.
Therefore, the number of particles present is concentration dependent.

50.  = i MRT The van’t Hoff Factor PA = i XAPA Tb = i Kb  m
We modify the previous equations for Colligative Properties by multiplying them by the van’t Hoff factor, i
Equations
Vapor pressure lowering:
Boiling point elevation:
Freezing point depression:
Osmotic pressure:
PA = i XAPA
Tb = i Kb  m
Tf = i Kf  m
 = i MRT

51. Colloidal Dispersion
homogeneous mixtures of particles larger than individual ions or molecules, but too small to be settled out by gravity.

52. Tyndall Effect
The scattering of light by colloids.

53. Colloids in Biological Systems
Some molecules have a polar, hydrophilic (water-loving) end and a nonpolar, hydrophobic (water-hating) end.
H2O
H2O
H2O
Polar Head
H2O
H2O
H2O
Sodium stearate is one example of such a molecule.
H2O
H2O
H2O
H2O
H2O
Non- Polar Tail
H2O
H2O
H2O
H2O
H2O
Soap Micelles in water solution

54. Colloids in Biological Systems
These molecules can aid in the emulsification of fats and oils in aqueous solutions.
{Micelle}
{Cotton Fabric Softener}

55. Suspension:
heterogenous fluid containing solid particles that are sufficiently large for sedimentation. Usually they must be larger than 1 micrometre.[1] The internal phase (solid) is dispersed throughout the external phase (fluid) through mechanical agitation, with the use of certain excipients or suspending agents.

61. 2006 (A)

65. 2003 A

66. Energetics of Solutions Heats of Solution (Hsoln)
Recyclable Hot Pack: Super Saturated Solution of Sodium Acetate
Na+(aq) + C2H3O2_(aq) NaC2H3O2(s)
crystallization
heating
∆Hsoln = - 67 kJ/n
{HotPack Na Acet}