1. Students will determine which combinations of ionic solutions form precipitates.

2. Students will identify the precipitate in each reaction.

3. Students will write net ionic equations for each reaction.

4. Students will write patterns or trends for certain cations and anions.

1. Construct a data table to record observations. Please list the chemicals in alphabetical order as above.

2. Add 1-2 drops of each pair of chemicals to a depression well on your spot plate.

3. If a precipitate forms, record that data including color.

4. If no reaction occurs, record N.R. on your data table.

5. Repeat procedures #1-4 for all chemical pairs listed on your data table.

The majority of ionic solids are soluble in water. Those that are not, form solid products called precipitates when two aqueous ionic solutions are mixed.

Recall that ionic compounds are made of positive and negative ions held together by the attractive, electrostatic forces that occur between oppositely charged particles. Soluble ionic compounds break apart completely into their respective ions when put in water. Example: NaCl _(s) when put into water yields Na⁺ _(aq) and Cl^- _(aq), and Ag(NO)_{3 (s)} also dissociates in water to form these respective ions - Ag⁺_(aq) and N0₃^-_(aq). It turns out that when these two solutions of sodium chloride and silver nitrate are mixed a solid falls out (precipitation occurs). The new mixture still contains Na⁺_(aq) and N0₃^- _(aq), however, the newly formed precipitate is AgCl _(s). The chemist describes this process first as a complete ionic equation.

Na⁺_(aq) + Cl^-_(aq) + Ag⁺_(aq) + NO₃^-_(aq) ---> AgCl _(s) + Na⁺_(aq) + NO₃^-_(aq)

Notice that the sodium and nitrate ions appear the same on both sides of the equation, that means they did not change; therefore, they are called spectator ions. Chemists like to write a more useful equation that describes only the changes that took place. They write a net ionic equation, which eliminates spectator ions.

NET: Ag⁺_(aq) + Cl^-_(aq) -----> AgCl _(s)

In this experiment, you will mix 13 different ionic solutions in all possible combinations to determine which combinations result in precipitate formation. Based on your results, you will write complete and net ionic equations for each reaction that has taken place. Use your solubility chart to determine what precipitates.

Dropper bottles with the following 0.20 M solutions.

1. barium chloride 7. potassium carbonate

2. cobalt chloride 8. potassium hydroxide

3. copper (II) sulfate 9. potassium iodide

4. iron (III) nitrate 10. silver nitrate

5. lead (II) nitrate 11. sodium carbonate

6. nickel (II) sulfate 12. sodium sulfide

1. Rinse your spot plates into the waste container provided.

2. Use Q-tips and soap to clean out each well, rinse with tap water.

3. Use distilled water as your final rinse.

4. Clean up your lab area and reagent table. Wash your hands thoroughly.

1. For each combination of solutions that GAVE A PRECIPITATE, write the correct net Ionic equation. Use a solubility chart to determine the identity of the precipitate.

2. Write these net ionic reactions in alphabetical order as they appear above. For example, the first reaction should be those containing barium chloride.

3. Carefully examine your net ionic reactions and data table, what patterns do you notice for certain precipitate reactions? What chemicals do not precipitate?