1. Solutions

2. Various Types of Solutions
Example
State of Solution
State of Solute
State of Solvent
Air, natural gas
Gas
Vodka, antifreeze
Liquid
Brass
Solid
Carbonated water (soda)
Seawater, sugar solution
Hydrogen in platinum

3. Solution Composition

4. Molarity

5. Exercise
You have 1.00 mol of sugar in mL of solution. Calculate the concentration in units of molarity.

6. Exercise
You have a 10.0 M sugar solution. What volume of this solution do you need to have 2.00 mol of sugar?

7. Exercise
Consider separate solutions of NaOH and KCl made by dissolving g of each solute in mL of solution. Calculate the concentration of each solution in units of molarity. 7

8. Mass Percent

9. Exercise
What is the percent-by-mass concentration of glucose in a solution made my dissolving 5.5 g of glucose in 78.2 g of water?

10. Mole Fraction

11. Exercise
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the mole fraction of H3PO4. (Assume water has a density of 1.00 g/mL.)

12. Molality

13. Exercise
A solution of phosphoric acid was made by dissolving 8.00 g of H3PO4 in mL of water. Calculate the molality of the solution. (Assume water has a density of 1.00 g/mL.)

14. Formation of a Liquid Solution
Separating the solute into its individual components (expanding the solute).
Overcoming intermolecular forces in the solvent to make room for the solute (expanding the solvent).
Allowing the solute and solvent to interact to form the solution.

15. Steps in the Dissolving Process

16. Steps in the Dissolving Process
Steps 1 and 2 require energy, since forces must be overcome to expand the solute and solvent.
Step 3 usually releases energy.
Steps 1 and 2 are endothermic, and step 3 is often exothermic.

17. Enthalpy (Heat) of Solution
Enthalpy change associated with the formation of the solution is the sum of the ΔH values for the steps:
ΔHsoln = ΔH1 + ΔH2 + ΔH3
ΔHsoln may have a positive sign (energy absorbed) or a negative sign (energy released).

18. Enthalpy (Heat) of Solution

19. Concept Check
Explain why water and oil (a long chain hydrocarbon) do not mix. In your explanation, be sure to address how ΔH plays a role.
Oil is a mixture of nonpolar molecules that interact through London dispersion forces, which depend on molecule size. ΔH1 will be relatively large for the large oil molecules. The term ΔH3 will be small, since interactions between the nonpolar solute molecules and the polar water molecules will be negligible. However, ΔH2 will be large and positive because it takes considerable energy to overcome the hydrogen bonding forces among the water molecules to expand the solvent. Thus ΔHsoln will be large and positive because of the ΔH1 and ΔH2 terms. Since a large amount of energy would have to be expended to form an oil-water solution, this process does not occur to any appreciable extent. 19

20. The Energy Terms for Various Types of Solutes and Solvents
Hsoln
Outcome
Polar solute, polar solvent
Large
Large, negative
Small
Solution forms
Nonpolar solute, polar solvent
Large, positive
No solution forms
Nonpolar solute, nonpolar solvent
Polar solute, nonpolar solvent

21. In General
One factor that favors a process is an increase in probability of the state when the solute and solvent are mixed.
Processes that require large amounts of energy tend not to occur.
Overall, remember that “like dissolves like”.

22. Structural Effects:
Polarity
Pressure Effects:
Henry’s law
Temperature Effects:
Affecting aqueous solutions

23. Pressure Effects Henry’s law: C = kP
C = concentration of dissolved gas
k = constant
P = partial pressure of gas solute above the solution
Amount of gas dissolved in a solution is directly proportional to the pressure of the gas above the solution.

24. A Gaseous Solute

25. Temperature Effects (for Aqueous Solutions)
Although the solubility of most solids in water increases with temperature, the solubilities of some substances decrease with increasing temperature.
Predicting temperature dependence of solubility is very difficult.
Solubility of a gas in solvent typically decreases with increasing temperature.

26. The Solubilities of Several Solids as a Function of Temperature

27. The Solubilities of Several Gases in Water

28. An Aqueous Solution and Pure Water in a Closed Environment

29. Vapor Pressures of Solutions
Nonvolatile solute lowers the vapor pressure of a solvent.
Raoult’s Law:
Psoln = observed vapor pressure of solution
solv = mole fraction of solvent
= vapor pressure of pure solvent

30. A Solution Obeying Raoult’s Law

31. Nonideal Solutions
Liquid-liquid solutions where both components are volatile.
Modified Raoult’s Law:
Nonideal solutions behave ideally as the mole fractions approach 0 and 1.

32. Vapor Pressure for a Solution of Two Volatile Liquids

33. Summary of the Behavior of Various Types of Solutions
Interactive Forces Between Solute (A) and Solvent (B) Particles
Hsoln
T for Solution Formation
Deviation from Raoult’s Law
Example
A  A, B  B  A  B
Zero
None (ideal solution)
Benzene-toluene
A  A, B  B < A  B
Negative (exothermic)
Positive
Negative
Acetone-water
A  A, B  B > A  B
Positive (endothermic)
Ethanol-hexane 33

34. Concept Check
For each of the following solutions, would you expect it to be relatively ideal (with respect to Raoult’s Law), show a positive deviation, or show a negative deviation?
Hexane (C6H14) and chloroform (CHCl3)
Ethyl alcohol (C2H5OH) and water
Hexane (C6H14) and octane (C8H18)
a) Positive deviation; Hexane is non-polar, chloroform is polar.
b) Negative deviation; Both are polar, and the ethyl alcohol molecules can form stronger hydrogen bonding with the water molecules than it can with other alcohol molecules.
c) Ideal; Both are non-polar with similar molar masses. 34

35. Colligative Properties
Depend only on the number, not on the identity, of the solute particles in an ideal solution:
Boiling-point elevation
Freezing-point depression
Osmotic pressure

36. Boiling-Point Elevation
Nonvolatile solute elevates the boiling point of the solvent.
ΔT = Kbmsolute
ΔT = boiling-point elevation
Kb = molal boiling-point elevation constant
msolute= molality of solute

37. Freezing-Point Depression
When a solute is dissolved in a solvent, the freezing point of the solution is lower than that of the pure solvent.
ΔT = Kfmsolute
ΔT = freezing-point depression
Kf = molal freezing-point depression constant
msolute= molality of solute

38. Changes in Boiling Point and Freezing Point of Water

39. Exercise
A solution was prepared by dissolving g glucose in g water. The molar mass of glucose is g/mol. What is the boiling point of the resulting solution (in °C)? Glucose is a molecular solid that is present as individual molecules in solution.

40. Exercise
You take 20.0 g of a sucrose (C12H22O11) and NaCl mixture and dissolve it in 1.0 L of water. The freezing point of this solution is found to be °C. Assuming ideal behavior, calculate the mass percent composition of the original mixture, and the mole fraction of sucrose in the original mixture.

41. Exercise
A plant cell has a natural concentration of m. You immerse it in an aqueous solution with a freezing point of –0.246°C. Will the cell explode, shrivel, or do nothing?

42. Osmosis – flow of solvent into the solution through a semipermeable membrane.
= MRT
= osmotic pressure (atm)
M = molarity of the solution
R = gas law constant
T = temperature (Kelvin)

45. Exercise
When 33.4 mg of a compound is dissolved in 10.0 mL of water at 25°C, the solution has an osmotic pressure of 558 torr. Calculate the molar mass of this compound.

46. van’t Hoff Factor, i
The relationship between the moles of solute dissolved and the moles of particles in solution is usually expressed as:

47. Ion Pairing
At a given instant a small percentage of the sodium and chloride ions are paired and thus count as a single particle.

48. Examples
The expected value for i can be determined for a salt by noting the number of ions per formula unit (assuming complete dissociation and that ion pairing does not occur).
NaCl i = 2
KNO3 i = 2
Na3PO4 i = 4

49. Ion Pairing Ion pairing is most important in concentrated solutions.
As the solution becomes more dilute, the ions are farther apart and less ion pairing occurs.
Ion pairing occurs to some extent in all electrolyte solutions.
Ion pairing is most important for highly charged ions.

50. Modified Equations

51. A suspension of tiny particles in some medium.
Tyndall effect – scattering of light by particles.
Suspended particles are single large molecules or aggregates of molecules or ions ranging in size from 1 to 1000 nm.

52. Types of Colloids

53. Coagulation Destruction of a colloid.
Usually accomplished either by heating or by adding an electrolyte.