# Solubility of Metal Salts

 Page 1/4 Date 11.11.2021 Size 30.25 Kb. #57668
1   2   3   4
Solubility of Metal Salts

## Solubility of Metal Salts

TOXICITY NOTE

Several of the metal cations used in this experiment can be toxic in high enough concentrations. While the concentrations of the solutions you will be using are small, nonetheless, when metals are in solution they are more easily absorbed into the body. Use caution when handling these materials, wear gloves, and assume every drop you see on the laboratory counters to contain these toxic metals. Be sure to place all waste materials in the appropriate waste containers.

Learning Outcomes

1. Use qualitative chemical reactions to establish solubility rules for 5 anions and 8 cations.

2. Gain an elementary understanding of metal complexes and learn to recognize signs of their presence in solution.

3. Practice recording thorough observations so conclusions can be made from the recorded data.

4. Practice balancing molecular equations for double displacement reactions of salts.
Introduction

In this experiment the student will be investigating some rules of solubility. By adding a series of anions to a selected group of metal cations, a classification system will be devised that can assist in the identification of metals in solution. You may want to refer to your course text for the rules of solubility.

One of the classifications will be due to the formation of a sulfide precipitate. While most of the anions will be provided to you in the form of the ammonium or sodium salts already dissolved in aqueous solution, you will produce the sulfide anion by a chemical reaction prior to its reaction with the metal cation. The source of sulfide will be hydrogen sulfide produced by the exposure of sodium sulfide to moist air.
Na2S(aq) + 2 H2O(l) H2S(g) + 2 NaOH (aq) Eq. 1
The hydrogen sulfide in turn may react with a metal cation already present in the test tube according to the equation below where “M” is any metal and in this example a charge of +3 was used (Eq.2) and we will assume that the salt M2S3 forms a precipitate. Since all metal nitrates are soluble in water, all our metals cations are in solution as metal nitrates for the lab. Naturally, for each metal cation one will have to balance the equation according to the stoichiometry of the reaction and the charge on the metal.
2 M(NO3)3(aq) + 3 H2S(g)  M2S3(s) + 6 HNO3(aq) Eq. 2
Other classifications will be due to the possible solubility of metal carbonates(CO32-) and sulfates(SO42-). An example of an equation for a salt containing carbonate is shown below(Eq. 3). Again “M” is used instead of a specific metal and this time we assume a charge on the metal ion of +2.
M(NO3)2(aq) + (Na)2CO3(aq) MCO3(s) + 2NaNO3(aq) Eq. 3
A particularly interesting precipitate is that of metal hydroxides (OH-). Some metal hydroxide precipitates can be forced back into solution by adding excess hydroxide. An example of this phenomenon can be seen with the lead II ion. The hydroxide forms an initial precipitate (Eq 4a) and with excess hydroxide re-dissolves (Eq 4b).
Pb(NO3)2 (aq) + 2 NaOH (aq)  Pb(OH)2 (s) + 2 NaNO3 (aq) Eq. 4a
Pb(OH)2 (s) + 2 NaOH (aq)  Na2[Pb(OH)4] (aq) Eq. 4b
The Pb(OH)42- species is called a metal complex and note that the compound Na2[Pb(OH)4] has (aq) after it indicating that it is soluble in water. Metal complexes are defined as metal cations covalently bonded to anions and/or neutral molecules. During the tests with anions other than hydroxide, you should observe the formation of one or two complexes. In this experiment, if a solution containing a metal cation changes color but a precipitate does not form when sulfide, carbonate, sulfate or thiocyanate is added, you can assume that a complex formed. If the precipitate dissolves in excess NaOH after forming a precipitate with a limited amount of NaOH, you may assume a complex formed in solution with or without a color change. See if you can recognize those situations in which complexes instead of precipitates form.
Remember when making your observations that all precipitates do not look alike. Precipitates can be colored or white. If the precipitate forms fairly large particles, then it will likely sink to the bottom of the test tube. If the precipitate forms very small particles of low density then it will likely remain suspended in the liquid forming a cloudy mixture. In this case you may not see it collect at the bottom on the test tube but you may see particles collect on the sides of the test tube.