08 – Selective precipitation

1. Chemistry 30 – Unit 2 – Solubility – Ch. 16 in Text 08 – Selective precipitation

2. What is “Selective Precipitation”? • Selective precipitation is a method we can use to separate mixtures of ions one at a time from a mixture containing many soluble substances. • This can be very practical if we want to isolate a specific ion: • More technical applications would include drug testing etc. where we want to determine if a specific substance is present in a solution

3. In depth example: • For example, imagine we mixed aqueous AgNO3, Cu(NO3)2, and Mg(NO3)2 together in a beaker. • *Knowledge from out solubility table tells us the following information: • compounds containing nitrate are ALWAYS soluble (regardless of the cation) • NO3- would be considered a spectator ion when we mix these three substances together (thus, we no longer need to consider them) • Ignoring the NO3 we basically have a beaker full of Ag+, Cu2+, and Mg2+.

4. In depth example: • To use the skill of selective precipitation: • What we want to do now, is add aqueous solutions that will form a precipitate with each of the cations above one at a time. • When a precipitate is formed, we can filter it out of the solution because it is solid. The original beaker now has one less ion in it. • This process is continued until no ions are left.

5. In depth example: • To bring the cations out of solution, We are given the following solutions to use: • Na2S, NaCl, and NaOH. • *Knowledge from out solubility table tells us the following information: • sodium ions, Na+, also always form soluble compounds • Na+ will remain as a spectator ion and we can ignore it • Thus, we are really given solutions of the following anions: • S2-, Cl-, and OH-

6. In depth example: • Order of adding solutions: • The order that we add the sodium compounds is crucial because the anions may be insoluble with more than one of the cations in the beaker • To help us analyze our situation, we'll prepare a chart, with the cations we need to separate along one axis and anions along the other axis:

7. Chart:

8. Chart: • Use a table of solubility to determine which combinations produce soluble compounds (sol) and which ones form compounds of low solubility. These are the ones that will precipitate out of solution (ppt - short-hand for precipitate).

9. Chart:

10. Order of adding solutions: • We need to determine the correct order for adding the negative ions. Be very careful with this next step - and be sure to read your chart in the right direction.

11. Order of adding solutions: • If we add S2- first, we would get two cations to precipitate - Ag+ and Cu2+. This would not help us, as they still be mixed together, this time as solids. • Adding OH- first would be even worse - all three cations would now be mixed together as solid precipitates.

12. Order of adding solutions: • Only by adding Cl- first can we separate out one of the ions, Ag+. • The others would remain in solution. • The formula of our first precipitate would be AgCl(s). • It is important to realize that while both S2- and OH- also form a precipitate with Ag+, once it has been removed using Cl- we no longer need to worry about it - it's all gone.

13. Order of adding solutions: • Once we've added Cl-, we're removed Ag+ from solution. So if we now add S2- we can again form a single precipitate - CuS(s), thus removing Cu2+from the mixture. • This will leave only Mg2+ still in solution. Sometimes this last cation is allowed to remain in solution. But if we chose we can add OH- and produce our final precipitate, Mg(OH)2 (s)

14. To summarize: • 1) First add NaCl to remove Ag+ , forming the precipitate AgCl(s) • 2) Second, add Na2S to remove Cu2+, forming CuS(s) • 3) Last add NaOH to remove Mg2+, forming Mg(OH)2 (s)