1. Ions in Solution
Chapter 14

2. I. Ionic Compounds in Aqueous Solution (Aqueous - water is solvent)
A. Theory of Ionization
1. Faraday - current causes ions to form
a. Electrolytes
b. Nonelctrolytes
2. Arrhenius - ionization of molecules in water produces ions

3. B. Dissolving Ionic Compounds
1. The solution process for ionic compounds
a. Hydration - solution process with water as solvent
b. Factors affect # of water molecules needed for hydration:
1) size of ion
2) charge of ion

4. 2. Heat of solution for ionic compounds
heat of hydration - energy released when ions become surrounded by water
a) Exothermic - releases heat ; negative heat of solution
b) Endothermic - absorbs heat ; positive heat of solution

5. 3. Dissociation - separation of ions when an ionic compound dissolves
NaCl ---> Na+(aq) +Cl-(aq)
1 mol mol mol
CaCl2 ---> Ca +2(aq) + 2Cl -(aq)
1 mol mol mol

6. C. Ionic Equations and Precipitation Reactions
1. Reactions in Solution
a. Precipitate (ppt) - insoluble substance formed through a chemical reaction in a solution
b. Some double replacement reactions produce ppt; others form a gas or water.

7. c. Solubility Table
1. i - insoluble - forms a ppt
2. ss - slightly soluble - formation of a slight ppt
3. s - soluble - no ppt forms

8. 2. Writing Ionic Equations
a. Write formula for compound.
sodium chloride = NaCl
b. Write the compound as ions: NaCl becomes Na+ + Cl-
c. Check solubility table to determine if a ppt forms
d. If all combinations give ‘s’ - reaction is NR
e. If one combination gives either ‘i’ or ‘ss’ - then a reaction takes place

9. e. Overall ionic equation includes all ions those that form a ppt and those that are referred to as ‘spectator ions’ because they do not form a ppt
f. Net ionic equation includes only those ions that form a ppt; cancel out the spectator ions on both sides of the equation.

10. Examples:
Write the overall ionic equation and the net ionic equation that occurs when aqueous solutions of zinc nitrate and ammonium sulfide are combined.
A solution of sodium sulfide is combined with a solution of iron(II) nitrate. Write the net ionic equation for any reaction that occurs.

11. II. Molecular Electrolytes (Polar covalent molecules can form electrolytes)
A. The solution process for molecular electrolytes
1. Polar molecules in water - opposite dipoles attract - if strong enough bond breaks and the molecule is separated into simpler charged parts

12. 2. Ionization - formation of ions from solute molecules by the action of the solvent
[Dissociation: ionic compounds Ionization: polar compounds]

13. B. The Hydronium Ion
1. H+ is only a proton, smaller than any other ion - it is attracted to others so strongly it does not have any independent existence
H H2O ---> H3O +
hydrogen ion water hydronium ion

14. C. Strong and Weak Electrolytes
1. Strong % ions
2. Weak - low concentration of ions

15. III. Properties of Electrolyte Solutions
A. Conductivity of Solutions
1. Strong-weak: degree of ionization
2. Concentrated-dilute; amount of solute-solvent
3. Ionization of H2O
2H2O ---> H3O+ + OH-

16. B. Colligative Properties of Electrolyte Solutions
1. Electrolytes affect colligative properties more than nonelectrolytes
Example: Compute the bp and fp for a solution made by adding21.6 g of NiSO4 to100 g of water.

17. 2. Theory vs Reality
a)Theory - electrolytes reduce fp by 2,3 times - depending on # of ions
b) Reality - reduces more than nonelectrolytes , but not as much as predicted
c) Reason - because ions are attracted to each other in water - more concentrated solutions have higher attraction for each other because they are closer together

18. 3. “Ideal Solution” - dilute enough that the ions have the expected activity

19. 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

20. 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
2. Freezing Point Depression - solution has a lower fp than solvent ∆ tf = Kfm
∆ tf- freezing point change Kf- molal freezing point constant m - molality of the solution

21. Example: What is the fp of water in a solution of 17
Example: What is the fp of water in a solution of g C12H22O11 and 200 g of water?

22. 3. Boiling Point Elevation - solution has a higher bp than solvent
∆ tb = Kbm
∆ 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?

23. 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

24. Example: When 1.56 g of an unknown , nonelectrolyte solute is dissolved in 200 g H2O, the ∆ tf = Co. Determine the molar mass.