Activity 1: What Dissolves and What Does Not: Establishing Some Rules
This activity may used with basic, general, and honors students.
These solutions include some heavy metal ions that are toxic in the environment. Although they are safe to use when handled properly, you must dispose of these ions carefully. A small-scale approach is suggested to reduce the volume of solutions required.
|Set I||Set II||Set III||Set IV||Set V|
Prepare 10% solutions of the reagents by dissolving 10 g of each in 90 mL water. If all of the sample does not dissolve let it settle and use the clear solution.
Place the solutions in dropper bottles and arrange in sets. Prepare at least five complete set-ups. These solutions are stable and can be used many times. Assign only one set to each student pair. (Beral™ pipets in an audio cassette case are also a convenient arrangement.)
Acetate sheets, or small glass plates work well. Acetate sheets are inexpensive and eliminate the risk of glass cuts. You may wish to photocopy copies of the grid and give several to each student, or make transparencies of the grid. Small-scale cell well plates work nicely. Students can rinse them and dry with Q-tips Grid sheets on colored paper enhance visibility of white precipitates.
You might find it convenient to provide a large chart on the chalkboard and have students record their data as they complete each part of the laboratory\par activity. The following chart is suggested as a possible format.
Place an (I) for Insoluble in the space where a precipitate was formed.
Place an (S) in spaces where there was no precipitate.
Generally, we consider anything insoluble if it has a solubility of less than 0.1 g per 100 mL. You may find it necessary to declare some compounds as slightly soluble (ss). Silver sulfate fits this category.
When students compile their data and compare their observations, it will become apparent that certain combinations of solutions produce precipitates. Further, it can be determined that only certain pairs of ions react to produce precipitates. When class data are gathered, some generalizations regarding the solubility of ionic compounds can be drawn. It is important that the general rules for solubility be developed as a result of student observations, and not presented as something to be memorized.
This activity gives students experience in (1) developing laboratory skills, (2) making and recording observations, (3) drawing conclusions and making generalizations from data, and (4) working as part of a team.
Students should be introduced to dissociation, the separation of ions in aqueous solution. They should identify all ions present in a solution and write correct ionic symbols. They must understand that precipitates form only when an insoluble product results from the recombination of ions in solution. They should also be able to write net ionic equations for precipitation reactions, eliminating spectator ions.
You might select, as an example, the combination of silver nitrate and sodium chloride solutions. Do this as a demonstrationsimply add one to the other; a dense white precipitate of silver chloride will form. Ask students to list all ions present in the combined solutions and eliminate those that would not precipitate. Show students that the precipitate could not be sodium nitrate by dissolving a mall amount of this compound in water to show that it is very soluble. Stress that the process of elimination to determine that the precipitate must be silver chloride is based on all four potential salts in each reaction plus recording all six solution salts as soluble in the beginning of the activity.
Having students pool and compare their data provides an opportunity to discuss experimental error and the need to collect a large quantity of data before establishing conclusions. Students should be prepared to support their conclusions with careful observations and data collection. They should also realize that the rules, theories, and hypotheses that chemists use are the direct or indirect result of laboratory experimentation. This laboratory activity allows every student to contribute to forming the rules by which chemists try to describe and understand their surroundings.
Student results will be encouragingly consistent. If some question arises as to whether or not a precipitate formed, you can mix several milliliters of the solutions in question in front of the class and ask for their observations.
The patterns match those found in general rules for water solubilities of ionic substances.
a. No precipitate d. No precipitate b. Precipitate e. Precipitate c. No precipitate f. No precipitate
After discussing student data and creating a set of generalizations regarding solubilities of certain compounds, call attention to the solubility rules in the textbook, or distribute a copy. Compare these two sets of rules. Spot check student understanding by posing possible combinations of solutions. (What would happen if solutions of lead nitrate and potassium carbonate were mixed? What would the precipitate be?)
You may find that barium hydroxide and sodium hydroxide (Set 4) produce a precipitate due to dissolved carbon dioxide in the sodium hydroxide, forming barium carbonate. If this happens, explain to students that it is not due to the combination of ions from the original compounds.
It is important that students understand what their laboratory activity has produced regarding knowledge of solubility rather than having students simply memorize solubility rules. Allow students to refer to the class-generated solubility chart and predict the solubility of various ionic compounds, or the precipitate formed when various ionic solutions are mixed.
Ask students to design sets of solutions to check the solubility of compounds not included on their charts.
Have students write net ionic equations, eliminating all spectator ions.