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

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

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

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

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.

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?

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

Example # 1: Dilutions What is the molarity of a solution made by diluting .0500 L of a 4.74 M solution of HCl to 250.0 mL?

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?

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

Precipitation Reactions

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

Precipitation Reactions

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 

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