1. Chapter 13 Solutions

2. Types of Mixtures Solutions Suspensions Colloids

3. A. Solutions 1. Definitions a. Soluble - capable of being dissolved
b. Solution - homogeneous mixtures
2. Nature of Solutions
a. Solvent - dissolving medium
b. Solute - substance dissolved
c. Electrolytes vs. nonelectrolytes

4. 3. Properties of Solutions
a. Particle type - ions, atoms, small molecules
b. Particle size nm
c. Effect of light - no scattering
d. Effect of gravity - stable
e. Filtration - cannot be separated
f. Uniformity- Homogeneous

5. 4. Types of Solutions
a. Gaseous solutions
b. Liquid solutions
c. Solid solutions (alloys)
solute solvent example
oxygen nitrogen air mixture
water drops air humid day
gold silver k gold

6. B. Suspensions 1. Definition - heterogeneous mixture 2. Properties
a. Particle type: large particles or aggregates
b. Particle size: 100 nm and larger
c. Effect of light: gives Tyndall effect

7. d. Effect of gravity: unstable
e. Filtration: can be separated
f. Uniformity: heterogeneous

8. C. Colloids 1. Definition: mixture between a solution and a suspension
2. Properties
a. Particle type: large molecule or particle
b. Particle size: nm

9. c. Effect of light: gives Tyndall effect
d. Effect of gravity: stable
e. Filtration: cannot be separated
f. Uniformity: borderline

10. Tyndall Effect

12. Is solute ionic compd or moelcular compound?
Dissolve solute in water and check if solution will conduct electricity
Ionic compound in aqueous form will conduct electricity, due to ions ability to carry charge – electrolyte
Ex. NaCl(aq) will conduct electricity

13. Molecular compound dissolved in water will not conduct electricity, because molecules are neutral – nonelectrolytes.
Ex. Sugar, C12H22O11 dissolved in water will not conduct electricity.

14. The Solution Process A. Factors Affect Rates of Dissolving
1. Surface area - smaller is faster
2. Agitation - stirring/shaking is faster
3. Heat
a. Liquids and solids - increase in heat mostly increases rate
b. Gases - increase in heat decreases rate

15. 4. Pressure
a. Liquids - no affect
b. Gases - increase pressure increase rate

16. B. Solubility
1. Saturated - maximum amount of dissolved solute
2. Unsaturated - less solute than saturated
3. Supersaturated - more dissolved solute than saturated

17. 4. Definitions:
a. Solution equilibrium - physical state in which the opposing processes of dissolving and crystallizing occur at an equal rate
b. Solubility - amount of substance required to form a saturated solution with a specific amount of solvent at a specified temperature

18. Solubility Table

19. C. Factors Affecting Solubility
1. Types of solvents and solutes
a. “like dissolves like” –
polar molecules will dissolve polar molecules and ionic compds due to attractions of opposite poles.
Nonpolar molecules will dissolve nonpolar molecules
b. miscible and immiscible

20. 2. Pressure
a. Liquids and solids - little to no effect
b. Gases - solubility increases with increase pressure
Henry’s law - solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid
Example: CO2 stays dissolved in soda with a pressure of over 5 atm.

21. 3. Temperature
a. Liquids and solids - increase temperature increases solubility (mostly)
b. Gases - increase temperature decreases solubility

22. D. Heats of Solution 1. Definitions
a. Heat of solution - amount of heat energy absorbed or released when a solute dissolves in a specific amount of solvent
b. Endothermic - positive heat of solution

23. c. Exothermic - negative heat of solution
d. Solvated - a solute molecule is surrounded by solvent molecules

24. 2. Solution formation
a. Solute-solute broken up: requires energy
b. Solvent-solvent broken up: requires energy
c. Solute-solvent attraction: releases energy
3. Positive heat of solution (endothermic)
when a + b > c
4. Negative heat of solution (exothermic)
when a + b < c

25. III. Concentration of Solutions
Concentration of solutions – amt of solute in given amt of solvent or solution
Universal solvent is water if not mentioned in the problem.

26. A. % by Mass
number of grams of solute dissolved in 100 g of solution #g of solute X g solution
Example: What is the % by mass of a solution prepared by dissolving 4.0 g HC2H3O2 in 35 g H20?

27. B. Molarity Number of moles of solute in one liter of solution
M = moles of solute L of solution
Example: What is the molarity of a solution composed of 5.85 g KI dissolved in enough water to make a 125 mL solution?

28. Ex. How would prepare 1. 00 L of 0. 500 M CuSO4 · 5H2O solution
Ex. How would prepare 1.00 L of M CuSO4 · 5H2O solution? Give clear directions.

29. C. Molality
Number of moles solute per kilogram of solvent m = moles solute kg of solvent
Example: How many grams of AgNO3 are needed to prepare a m solution in 250 g of water?

30. IV. Colligative Properties of Solutions
A. Definition - a property that depends on the number of solute particles but is independent of their nature
1. Nonelectrolytes - 1 solute particle
2. Electrolytes - # of solute particles dependent on # ions
NaCl: AgNO3: 2
MgCl2: 3 K3PO4: 4

31. B. Types of Colligative Properties
1. Vapor Pressure Lowering - the tendency for molecules to escape from a liquid to a gas is less in a solution than a pure solvent

32. 2. Freezing Point –
solution has a lower fp than solvent
Explanation: when solute is added to solvent, the solute particles interfere with the solvent particles from coming together as a solid so the temp is lowered to get the solution to freeze. Illustrate.

33. Δ tf=freezing point change= orig. f.p.– new f.p.
The amount of decrease in freezing temperature when a solute is dissolved is called freezing point depression,
Δtf = Kfm
Δ tf=freezing point change= orig. f.p.– new
f.p.
Kf=molal freezing point constant m = molality of the solution
Example: What is the fp of water in a solution of g C12H22O11 and 200 g of water?

34. 3. Boiling Point –
solution has a higher bp than solvent
Explanation: When solute is added to solvent, the solute particles interfere with the solvent particles from boiling up so the temp is raised to get the solution to boil. Illustrate.

35. The amount of increase in boiling temperature when a solute is dissloved is called boiling point elevation, Tb
Δ tb = Kbm
Δ tb = change bp Kb = molal boiling pt constant
m = molality
Example: What is the bp of a solution that is made by adding 20 g C12H22O11 in 500g H20?

36. C. Determination of Molar Mass of a Solute
1. Determine Δtf(Δ tb)
2. Determine m
Δ t = Km
3. If ionic divide by number of particles
4. Calculate moles of solute
m X kg of solvent
5. Molar mass = mass of solute moles of solute

37. Example: When 1.56 g of an unknown , nonvolatile solute is dissolved in 200 g H2O, the Δ tf = Co. Determine the molar mass.