1. Review of Solubility and Precipitation Reactions

2. Aqueous Solutions Compounds dissolved into water.
Can contain molecules or ions in a solution.
How do you distinguish between ion or molecule?

3. DISSOCIATION !!
The ability of a compound to breakdown in a solution into individual ions
Ionic Compounds
Break down into cations and anions
Electrical conductors—ions flow through solution
Molecular Compound
Compound remains intact as “molecules,” no breakdown
Generally NOT electrical conductors

4. Solubility
How much solute dissolves in a solution to produce a saturated solution
Temperature and Pressure dependent
Increase with increasing temperature
Increases with decreasing temperature (ex. Water in lake)
Pressure increases, solubility increases (ex. Soda can)
TEMPERATURE:
Solubility increase with increasing temperature—applies for most substances, good for liquids
Solubility increases with decreasing temperature—applies to gases (talk about how this applies to aquatic life—eutrophic vs. anoxic conditions with oxygen in water supporting organisms)
PRESSURE:
-not a big influence for solids/liquids
-huge influence on gases
-increased solubility of gas when partial pressure increases in area above solution (p. 476)
**Soft Drinks**--gas escapes from soda as pressure decreases and causes solubility to decrease

5. Which compounds are soluble in water
Which compounds are soluble in water? 1) BaCl2 2) Pb (NO3)2 3) Na2S 4) BaCO3 5) PbS
Soluble
Insoluble

6. Precipitation Prediction
Write the reactants in ionic from
breakdown into ionic form if compounds are soluble
leave as molecules if insoluble
Determine the solubility of the products.
Use solubility Rules

7. Precipitation Predictions (cont.)
3) Check to see if one product is insoluble in water.
Product will fall out of solution, identified as precipitate
4) Write the net ionic equation
Displays which ions are directly involved in the reaction, produce the precipitate
Ions existing on BOTH sides of the equation are “spectator ions” (do NOT participate in precipitate formation)
Spectator ions are eliminated

8. Example 1:
MgSO KOH
Write the net ionic equation. Will a precipitate form?

9. Solubility Product Constant (Ksp)

10. Chemical Equilibrium so far-----
Gases
Acids and Bases
Slightly soluble Salts
Many ionic compounds—only a small fraction dissolves

11. Example 1: BaSO4(s)  Ba+2(aq) + SO4-2(aq)
Indicates salt exists in “solubility equilibrium”—some dissolves, some does not

12. Solubility Product Constant (Ksp)
Equilibrium constant for slightly soluble salts
Indicates equilibrium between solid salt and the ions found in a solution when it dissociates
Expression represents the product of the concentrations of ions in equilibrium
Temperature dependent
Values found in table along with solubility equation (p. 678, Appendix C—p.A18)

13. Example 2:
Write a Ksp expression for an equilibrium in a saturated aqueous solution of iron (III) phosphate and for an equilibrium in a separate aqueous solution of chromium (III) hydroxide.

14. Ksp and Solution Molarity
Ksp is an equilibrium constant, NOT concentration
Molarity/concentration separate from Ksp
Ksp values
Considered estimates due to ion attractions to other ions in solution and Ksp values not exact
Used only for slightly soluble salts
Increase concentration of ions, increase Ksp value

15. Example 3:
A saturated aqueous solution of silver carbonate contains 32 mg of Ag2CO3 per liter at 20°C. Calculate the Ksp for Ag2CO3 at 20°C.

16. Example 4:
Based on a Ksp value of 1.4 x 10-5 at 25°C for silver sulfate, calculate this compound’s molar solubility at 25°C.
Ag2SO4(s)  2Ag+(aq) + SO4-2(aq)

17. Example 5:
Without any calculations, arrange the following in order of INCREASING molar solubility.
MgF2, CaF2, PbCl2, PbI2

18. Homework
Ksp Practice Problem Set #1