Chemistry 30

Solubility of Various Substances - demonstration of various soluble and insoluble substances

Ionic Compounds produce Ions in Solution - Lesson 1.1

Students will need to understand that when ionic compounds dissolve they form ions in solution which will conduct a current. For example:

NaCl_(s)→ Na⁺_(aq) + Cl^-_(aq)

Conductivity can easily be demonstrated with a simple conductivity indicator which can be purchased from most scientific catalogues. Boreal used to carry an inexpensive version (search: "conductivity indicator"), or one could be made without too much effort.

Demonstrate and compare the conductivity of various solutions. Sugar and salt, for example, both dissolve well in water, but the sugar solution will not conduct a current since ions do not form. Acids are molecular compounds that will produce ions in solution, and thus will conduct a current.

See also the Prentice Hall video, Electrolytes and Non-electrolytes

 

 Solubility of Various Substances - Lesson 1.1

Demonstrate that some substances dissolve very well in water (sugar, salt) while others have a very low solubility (chalk dust - calcium carbonate, iodine crystals, oil and water). Iodine crystals dissolve poorly in water but do dissolve well in alcohol (producing a tincture of iodine).

 

Supersaturated Solutions - Lesson 1.1

Sodium thiosulfate, Na₂S₂O₃, readily forms a supersaturated solution. Fill the bottom few cm of a test tube or small beaker with the compound and add a small amount of water (5 mL of water would be sufficient for 50 g Na₂S₂O₃). Make sure not to add so much water that the entire sample dissolves. This dissolving process is endothermic, and the test tube will cool noticeably.

Gently heat the test tube in a warm water bath until all of the sodium thiosulfate has dissolved. Allow to cool to room temperature. An ice bath may be used but is not required. The excess sodium thiosulfate will remain in solution as a supersaturated solution.

Add one crystal of sodium thiosulfate. This will rapidly and dramatically cause the excess sodium thiosulfate to crystallize and come out of solution. The reaction is noticeably exothermic.

The recrystallization can also be shown to the class by placing the supersaturated solution in a petrie dish on an overhead , or allow each student or pair of students to prepare their own supersaturated solution.

The solid may be used again to form another supersaturated solution.

Sodium acetate trihydrate, NaC₂H₃O₂ · 3H₂O, works well also. Dissolve 125 g in 50 mL of distilled water, and heat to form a supersaturated solution.

With some practice you can carefully pour the supersaturated solution onto a glass plate containing a few seed crystals to create a tower of solid sodium thiosulfate out of the clear solution. View a demonstration

 

Temperature and Solubility - Lesson 1.3

Most solids demonstrate an increase in solubility as temperature increases.

This can be demonstrated by filling a large test tube about half full with sugar. Add water at room temperature until the test tube is nearly full. Stir to dissolve as much of the sugar as possible, but there should be solid sugar remaining - if not, add more sugar.

Place the test tube in a hot water bath or heat carefully over a bunsen burner until the solution is near the boiling point. The sugar should dissolve completely.

Some solids actually demonstrate a decrease in solubility as temperature increases. This will be noticed in solids in which the dissolving process is exothermic. Increasing the temperature will favor the endothermic process, and the substance will come out of the dissolved state.

Dissolve 37 g of calcium acetate in 100 mL of water and place in a Florence flask or other suitable container. Gently heat in a hot water bath or over a bunsen burner. The students should observe that the solution becomes cloudy as the calcium acetate comes out of solution.

Next place the container in an ice water bath. The calcium acetate should re-dissolve.

A Blue Precipitate - Lesson 3.4

A blue precipitate of Cobalt(II) hydroxide can be produced by adding a few drops of cobalt chloride solution to about 100 mL of a saturated limewater solution (calcium hydroxide).

Cobalt Chloride solution preparation. The concentration of this solution is not a critical factor. For demonstration purposes, dissolve a gram or two of CoCl₂ in about 25 mL of water. May be stored in a dropper bottle.

Limewater solution preparation. Prepare a saturated limewater solution by dissolving 1.5 g Ca(OH)₂ in 1 L water. Shake vigorously for a minute or two then let sit overnight. Filter off the clear solution - this is the limewater.

 

Common Ion Effect - Lesson 3.5

A simple demonstration of the common ion effect.

Half-fill a beaker with 1M NH_{3 (aq)}.

The NH₃ is in equilibrium with the ammonium ion, NH₄^+:

NH3 + H2O		NH4+ + OH–
colorless		magenta

Add a few drops of phenolphthalein. The solution will turn a magenta color, indicating a basic solution due to OH^- ions.

Next add a little solid NH₄Cl.  This will cause the concentration of NH₄⁺ to increase. The magenta color should disappear as equilibrium shifts to the left, demonstrating Le Châtelier's Principle.

Supersaturated Solutions

Several web sites show interesting ways to cause supersaturated solutions to recrystallize. Search terms: crystallization supersaturated solution

University of Wisconsin-Madison Chemistry Department Demonstration Lab

 

Dissolution of NaCl in Water

One of the excellent videos from the Prentice Hall "General Chemistry: Principles and Modern Applications" web site. Check this site for other useful videos.

Electrolytes and Nonelectrolytes

This Prentice Hall video illustrates the difference between electrolytic and nonelectrolytic solutions, due to the presence of free moving ions in the former.

Solution Formation from a Solid

This Prentice Hall video illustrates how a standard solution is prepared.

Gas Pressure and Solubility

This Prentice Hall videos help explain the relationship between pressure and the solubility of gases.