Investigating the solubilities of lead halides
The lesson may begin with a discussion of general solubility patterns among common ionic compounds. From these patterns, students predict the pattern of solubilities among the lead halides, and what might happen when solutions containing lead ions and different halide ions are mixed. They can test their predictions experimentally using dilute solutions containing the relevant ions. They then follow this up by testing the solubilities of the lead halides when heated, and confirm the pattern by allowing the solutions to cool again. Throughout they use ionic equations when making predictions and explaining observations.
This is a straightforward class practical, based on predicting behaviour from known patterns of chemical properties. The known patterns of solubilities and of properties of Group 17 compounds will need to be familiar to students – if not these will need to be discussed first.
The practical itself will take about 25 minutes to the point at which the hot solutions are ready to cool down, plus however long it then takes for the precipitates to re-form on cooling.
Each working group will need
Boiling tubes (large test-tubes), 3
Beaker (250 cm
Heat resistant mat
The following solutions should be available in dropper bottles
Lead nitrate, 0.2 M (TOXIC, DANGEROUS FOR THE ENVIRONMENT), 20 cm
Potassium chloride, 0.2 M, 10 cm
3 (Note 2)
Potassium bromide, 0.2 M, 10 cm
3 (Note 2)
Potassium iodide, 0.2 M, 10 cm
3 (Note 2)
Refer to Health & Safety and Technical notes section below for additional information.
Health & Safety and Technical notes
Read our standard health & safety guidance
Wear eye protection. Wash hands after using lead nitrate solution.
Lead nitrate (TOXIC, DANGEROUS FOR THE ENVIRONMENT both as solid and as 0.2 M solution) - see CLEAPSS Hazcard and CLEAPSS Recipe Book.
Potassium chloride - see CLEAPSS Hazcard and CLEAPSS Recipe Book.
Potassium bromide - see CLEAPSS Hazcard and CLEAPSS Recipe Book.
Potassium iodide - see CLEAPSS Hazcard and CLEAPSS Recipe Book.
1 If dropper bottles are not available for dispensing solutions, students will also require dropping pipettes to dispense solutions.
2 Sodium halides can be used instead of potassium halides if preferred.
a Place 3 boiling tubes in a row in a test-tube rack.
b To each tube add about a 3 cm depth of 0.2 M lead nitrate solution.
c To the first tube add 5 drops of 0.2 M potassium chloride solution, and note what happens. Keep the mixture formed for g below.
d To the second tube add 5 drops of 0.2 M potassium bromide solution, and note what happens. Keep the mixture formed for g below.
e To the third tube add 5 drops of 0.2 M potassium iodide solution, and note what happens. Keep the mixture formed for g below.
f Heat each of the test tubes from b, c, d in turn in a low Bunsen flame until the mixtures are boiling. Allow them to boil very gently for a minute.
g Allow the three boiling tubes to cool down. If time is short, they may be cooled by standing in a beaker of cold water.
h Observe and record what happens in each of the three tubes as the mixtures in them cool.
This practical can be carried out either as an investigative experiment, as described above, or as a simple exploration without use of prior knowledge. In the latter case, students will not need to be familiar with solubility patterns of ions in solution and their reactions, or with ionic equations.
However, to make it an investigation, the points below need to be considered:
1 Students will need some generalizations about the solubilities of salts before they can be asked to predict whether or not a precipitate will form on mixing two solutions and what the precipitate will be. If these have not been taught previously, the lesson may need to start with a minimum of these generalizations:
all nitrates are soluble
all sodium and potassium salts are soluble
all chlorides are soluble except for the chlorides of lead and silver
all lead salts are insoluble in cold water except for the nitrate (and ethanoate).
2 Students also need to be able to think in terms of cations and anions ions behaving independently in solution. Diagrams such as those below can be helpful.
Students need to be able to translate these diagrams into the formalities of chemical equations; eg the ionic equation: 3
2+(aq) + 2I -(aq) →PbI 2(s)
and the full equation:
3) 2(aq) + 2KI(aq) →PbI 2(s) + 2KNO 3(aq)
The solubilities of the lead halides increase markedly with temperature, so that the three halides under investigation are all effectively soluble in boiling water. This means that on cooling these solutions, the lead halides will crystallise out again. For lead chloride and lead bromide, the effect is rapid and the crystals small, so their appearance returns to that of a precipitate.
However for lead iodide, especially on slower cooling, the effect of recrystallisation can be spectacular, with thin golden flaky crystals of lead iodide shimmering in suspension, and falling as golden rain to the bottom of the tube. If the students do not observe this phenomenon, it is well worth the teacher repeating this stage as a demonstration.
Silver and lead halides from this series.
Health & Safety checked, 2016
This Practical Chemistry resource was developed by the Nuffield Foundation and the Royal Society of Chemistry.
© Nuffield Foundation and the Royal Society of Chemistry
The variation of crystal size for lead iodide with rate of cooling can be used as a laboratory model to demonstrate how rate of cooling of magmas affects crystal size in the resulting igneous rocks:
Joint Earth Science Education Initiative
Page last updated October 2015