1. Unit 11 – Solutions and Ions in Aqueous Solutions
11.c – Molarity, Dilutions and Solubility Rules

2. Concentration
Measure of the amount of solute in a given amount of solute

3. Molarity is most often used to specify the concentration of a solution
number of moles of solute in one liter of solution
units: moles/liter = M

4. Example 1
21.0 g of NaOH is dissolved in enough water to make 500. mL of solution.

5. Example 2
3.7 moles of HCl is added to water to make 500. mL of solution. The Density of the solution is 1.10 g/mL.

6. Example # 3
You have a 6.0 M solution of sodium chromate. If you wanted to have a sample of solution that only contained .149 moles, how much of the solution would you need?

7. Dilutions
when water is added to a standard solution to decrease the molarity to a desired level

8. Example # 1: Dilutions
What is the molarity of a solution made by diluting L of a 4.74 M solution of HCl to mL?

9. Example #2: Dilutions
What volume of water would you add to 15.0 mL of a 6.77 M solution of nitric acid in order to get a 1.50 M solution?

10. Precipitation Reactions
Remember that all ionic compounds are considered strong electrolytes
dissociate completely in water
When two solutions are mixed, free ions are floating around
They collide together randomly
If an ionic compound forms that is insoluble, it is called a precipitation reaction

11. Precipitation Reactions

12. Precipitation Reactions
If no insoluble compound forms, no reaction has actually happened
It is just a solution with lots of free ions floating around
Spectator Ions
free ions not involved in chemical reaction

13. Precipitation Reactions

14. Solubility Rules
Rules that help us predict which compounds will be soluble/insoluble
See handout – you will have a copy of this on tests & quizzes 

15. Practice – are the following soluble (S) or insoluble (I)
Aluminum nitrate
Magnesium chloride
Rubidium sulfate
Silver chloride
Ammonium hydroxide
Lead (II) sulfide
Sodium phosphate
Iron (III) phosphate