Solubility trends of metal halides

In this experiment, students use drops of a variety of silver and metal halide solutions to explore similarities and differences in their solubility. The experiments show that while silver fluoride is soluble, other silver halides are insoluble, and that fluorides behave differently from the rest of the halides.

This is a class experiment in which students can very quickly and simply investigate trends in the solubilities of halides of some metals on a microscale, using only drops of solution. The experiment should take about ten minutes.

Equipment

Apparatus

• Eye protection
• Worksheet containing a grid on which drops of solution are combined (available for download below)
• Clear plastic sheet such as OHP film

Chemicals

Access to the following solutions contained in plastic pipettes (see notes 10 and 11 below):

• Calcium nitrate, 0.2 M
• Lithium bromide, 0.2 M
• Sodium fluoride, 0.2 M
• Sodium chloride, 0.2 M
• Potassium bromide, 0.2 M
• Potassium iodide, 0.2 M
• Silver nitrate, 0.1 M

Health, safety and technical notes

• Read our standard health and safety guidance.
• Wear eye protection throughout.
• Calcium nitrate solution, Ca(NO₃)₂(aq) – see CLEAPSS Hazcard HC019B.
• Lithium bromide solution (similar to lithium chloride), LiBr(aq) – see CLEAPSS Hazcard HC047b. 
• Sodium fluoride solution, NaF(aq) – see CLEAPSS Hazcard HC095C. 
• Sodium chloride solution, NaCl(aq) – see CLEAPSS Hazcard HC047b and CLEAPSS Recipe Book RB082.
• Potassium bromide solution, KBr(aq) – see CLEAPSS Hazcard HC047b.  
• Potassium iodide, KI(aq) – see CLEAPSS Hazcard HC047b and CLEAPSS Recipe Book RB072. 
• Silver nitrate solution, AgNO₃(aq) – see CLEAPSS Hazcard HC087 and CLEAPSS Recipe Book RB077.  Note that even this dilute solution can stain fingers and clothing.
• It may be useful to colour code the pipettes containing the solutions, to prevent confusion.
• All the solutions must be made up using deionised or distilled water. This is to prevent any unwanted cloudiness appearing on mixing with silver nitrate because of the presence of chloride ions in tap water.

Procedure

1. Cover the worksheet containing the table with the plastic sheet.
2. Put one drop of each of the halide ion solutions in the appropriate boxes.
3. Add one drop of silver nitrate solution to the drop of halide solution in each of the boxes in the first column.
4. Repeat the process for the second column using lithium bromide solution.
5. Repeat the process for the third column using calcium nitrate solution.
6. Record all your observations when the solutions are mixed.

Teaching notes

The first column compares the solubilities of silver halides. Silver fluoride is soluble, so no precipitate forms. The other silver halides are all insoluble and form as precipitates. The reactions can be represented by the ionic equation:

Ag⁺(aq) + X^–(aq) → AgX(s), where X = Cl, Br and I.

It may be possible to distinguish the silver halides by colour: silver chloride is white (but darkens rapidly in strong light), silver bromide is off-white or cream, and silver iodide is pale yellow.

For the lithium halides in the second column, only lithium fluoride is insoluble, but a precipitate does not always form at these concentrations.

For the calcium halides, only the fluoride is insoluble, forming a white precipitate.

These experiments show that fluorides behave differently from the rest of the halides in each series. In Groups 1 and 2, the behaviour of the fluorides is not typical of the rest of the halides. This is because all the other halides of the metals in these groups are soluble.

At a suitable level, these trends can be discussed with students in terms of ion size, lattice energy and hydration energies of the aqueous ions. The small size of the fluoride ion is the main reason for it behaving differently from the other halide ions here.