1. Chapters 4 & 11: Properties of Solutions
AP CHEMISTRY

2. 4.1 Water, the Common Solvent
Many common chemical reactions occur in water, or aqueous solution. To understand how chemical species interact in solution, we must first understand water, the universal solvent.
Water is an excellent solvent due to:
Its shape; water is a bent molecule.
Electrons aren’t shared evenly (oxygen is more electronegative)
Electrons spend more time close to O than to H
This uneven distribution of charge makes water polar.

3. Hydration Water is held together by covalent bonds
When water surrounds an ionic crystal, the H end attracts the anion and the O end attracts the cation. This process is called hydration.
Hydration causes salts to dissolve. H2O also dissolves polar covalent substances such as C2H5OH.

4. Water and Nonpolar Molecules
H2O doesn’t dissolve nonpolar covalent substances.
The difference in a substances’ ability to dissolve is due to its interaction with itself, and the solvent solution.

5. 4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes
Review:
Solute-the substance that dissolves
Solvent-the substance that does the dissolving
Electrical conductivity-the ability of a solution to conduct an electrical current
Strong electrolyte-a highly ionized solution that easily conducts electrical current
Weak electrolyte-a solution with few ions that does not conduct a current very well
Nonelectrolyte- a solution made of a soluble compound that does not ionize and thus does not conduct an electrical current.

7. Svante Arrhenius determined that the extent to which a solution can conduct an electrical current depends directly on the number of ions present.
A solute that ionizes completely conducts very well.
He said that the best conductors are:
Soluble salts (like sodium chloride or magnesium nitrate)
Strong acids like HCl(aq), HNO3(aq), H2SO4(aq)
strong bases that contain OH- (have a bitter taste, slippery feel and include NaOH and KOH)

8. Ex. Write an equation showing what happens when the salt sodium acetate dissolves in water.
NaC2H3O2  Na+ + C2H3O2-

9. Arrhenius went on to describe weak electrolytes
Arrhenius went on to describe weak electrolytes. He said they only ionize slightly and include weak acids and bases.
HC2H3O2  H+ + C2H3O2-
99% %
Ammonia (NH3) is a weak base NH3 + H2O  NH4+ + OH-

10. Solubility- g/given volume solvent or moles/given volume solution

11. 4.3 The Composition of Solutions
In order to perform stoichiometric calculations with solutions, we must know the nature of the solution and the amounts of chemical present.
M = moles of solute
liters of solution

12. Ex. Calculate the molarity of a solution made by dissolving 23
Ex. Calculate the molarity of a solution made by dissolving 23.4g of sodium sulfate in enough water to form 125 mL of solution.
23.4 g Na2SO4 1 mol Na2SO4 = mol Na2SO g Na2SO mol = 1.32 M L

13. Ex. How many grams of Na2SO4 are required to make 350 mL of 0
Ex. How many grams of Na2SO4 are required to make 350 mL of 0.50 M Na2SO4?
0.350L mol Na2SO g Na2SO4 = 25 g
1 L mol Na2SO4

14. Dilution problem (M1V1 = M2V2) (1.000M)(V1) = (0.250M)(500.0mL)
Ex. What volume of M KNO3 must be diluted with water to prepare mL of M KNO3?
Dilution problem (M1V1 = M2V2)
(1.000M)(V1) = (0.250M)(500.0mL)
V1 = 125 mL

15. A standard solution is one where concentration is accurately known
A standard solution is one where concentration is accurately known. Read procedure for using volumetric flasks and types of pipets. We will be using both in several labs this year.

16. 11.1 Solution Composition
Solutions can generally be described as dilute (very little solute per volume of solvent) of concentrated (a lot of solute per volume of solvent). Another way to describe a solution is by looking at the mass of the solute in terms of the mass of the entire solution, or mass percent.
We can also examine the percent solute to solvent using the mole fraction of the solution, or χ.
Finally, we can look at the amount of solute in moles per amount of solvent in kg to find molality.

17. Formulas for Solutions
Mass % = Mass solute x 100 Mass of solution Mole Fraction χA = nA . nA + nB Molality = moles solute Kg solvent

18. Ex. A solution is made using 10. 5g ethanol (C2H5OH) in 1,200
Ex. A solution is made using 10.5g ethanol (C2H5OH) in 1,200.0 g of water to make a final volume of ml. Find the molarity, molality and mole fraction of this solution.
10.5g C2H5OH 1 mol C2H5OH = 0.228mol C2H5OH
46.08g C2H5OH
1,210.5 g soln = L soln
1,200.0g H2O/18.02g = mol H2O
M = mol / L = 0.188M C2H5OH
m = mol/ kg = 0.19 m C2H5OH
χC2H5OH = (0.228) =
( )

19. 11.2 The Energies of Solution Formation
When a solution forms, it takes place in three distinct steps:
Expanding the solute- The solute particles must move away from one another. This may form ions or molecules. (ΔH1)
Expanding the solvent- The solvent must break any internal bonds so that it can surround and hydrate the solute. (ΔH2)
The solute and solvent begin to interact to form the solution. (ΔH3)

20. Solution Formation-Exothermic

21. Solution Formation-Endothermic

22. Heat of Solution (ΔHsoln)
The heat of solution (ΔHsoln) is equal to the sum of heat produced or consumed in the three steps.
If the heat of solution is very positive, then energy must come into the system to make the solute dissolve, so the substance is insoluble.
If the heat of solution is very negative or only slightly positive, then energy is leaving the system as the solute dissolves and the solute is soluble.
The table below summarizes these interactions.

23. Heat of Solution ΔH1 ΔH2 ΔH3 ΔHsoln Outcome Polar solute and solvent
ΔH1
ΔH2
ΔH3
ΔHsoln
Outcome
Polar solute and solvent
Large positive value
Large negative value
Small positive value
Solution forms
Nonpolar solute, polar solvent
No solution forms
Nonpolar solute and solvent
Polar solute, nonpolar solvent

24. 11.3 Factors Affecting Solubility
Although we will discuss many of these in other chapters, the following summarizes the reasons for the values of H in the chart above.
Structural effects- the shape of solute molecules affects their polarity, and thus their ability to dissolve in solvents.
Ex. Vitamin A vs. Vitamin C

25. 11.3 Factors Affecting Solubility
Pressure effects- as pressure increases on the surface of a solvent, it may increase the rate of dissolution of a solid or liquid solute, but it does not affect the solubility of these solutes. If the solute is a gas, however, increasing pressure on the system will increase the solubility of the gas in the solvent.
Temperature effects- again, as temperature increases in a solvent, the rate of dissolution will increase. Many solid and liquid solutes will become more soluble, but a few solid and liquid solutes, and almost all gaseous solutes, will become less soluble as temperature increases.

26. 11.4 The Vapor Pressure of Solutions
When a solution is created, the properties of the solvent are often changed. For example, adding a solution of water and antifreeze will allow your car to keep running, whether the temperature outside is very high or very low.

27. Volatile VS Nonvolatile
A volatile compound is one that vaporizes easily due to weak internal bonding (ex. Acetone)
A nonvolatile compound has strong internal bonding and thus does not vaporize easily (ex. water)
When a volatile solute is added to a solvent like water, the volatile compound will vaporize to some extent, adding to the amount of particles above the surface of the solution. In a closed container, this increases the vapor pressure.
Nonvolatile solutes, however, will not change the amount of vaporized particles. These types of solutions are called ideal solutions.

28. This behavior was described by Francois Raoult, in his law
This behavior was described by Francois Raoult, in his law. Psoln = χsolvent .Psolvent
You can see that the more solute is added, the smaller the value for the mole fraction of solvent, and thus the smaller the pressure.
If a solute is added that actually attracts water molecules, it may lower vapor pressure in a closed container by keeping water in a liquid state. Solutes like salt actually bind to water and keep it from vaporizing, therefore reducing the number of water vapor particles above the surface of the solution.

29. 11.8 Colloids Matter can be separated as shown in the diagram below.
One of the ways we can classify solutions is to look at the particles they are made of.
Homogenous solutions are made up of relatively small particles that are permanently suspended in solvent.
Heterogeneous solutions are made up of relatively large particles that will settle out of solution over time.

30. Colloids
Colloids fall into a special category because they are made up of intermediately sized particles that stay suspended due to interactions with the surrounding solvent.
The only way to distinguish between a colloid and a solution is to shine a light through them. Colloid particles are large enough that they bounce light, creating a “beam” you can see (called the Tyndall effect). Solutions will not show a “beam” effect.