1. The Solubility Rules

2. Many Ionic Compounds are Soluble in Water Water is a highly polar molecule, with strong dipoles Ionic bonds are extremely strong, but many cannot withstand constant collisions with water molecules. Ionic bonds are extremely strong, but many cannot withstand constant collisions with water molecules. The solubility of ionic compounds varies with temperature Discounting temperature, however, there is a wide range of solubility for ionic compounds.

3. Question 1 When calcium chloride is dissolved in water, to which end of the adjacent water molecules will a calcium ion be attracted? the oxygen end, which is the negative pole the oxygen end, which is the positive pole the hydrogen end, which is the negative pole the hydrogen end, which is the positive pole

4. Solubility Rules There are a set of rules that chemists use to predict precipitates in double replacement reactions. These rules are found on Table F; they have to be memorized in college classes

5. Note that Table F has much of the same information that Table E does: The more common polyatomic ions are listed on Table F along with their formulas in a slightly more accessible form.

6. Let’s Look at the Solubility Rules Some ions are always soluble, no matter what the counter ion is! What is the important group 1 ion not mentioned here? Why is seawater salty?

7. Some Ions are more complex… There are exceptions for halides and sulfates, which are generally soluble Heavy metals are usually insoluble with everything but nitrate and acetate

8. Based on Reference Table F, which solution will contain the highest concentration of iodide ions? lead iodide silver iodide mercury (II) iodide iron (II) iodide

9. Ions that are usually insoluble: Note that there are no truly insoluble compounds: tiny amounts of even the most insoluble compound do dissolve

10. What if there were no insoluble ions?? No fossils: no hard parts! No shelled animals (CaCO 3 ) No vertebrates (you and me – Ca 3 PO 4 ) No heavy metal toxicity: sulfur binds these metals at extremely low concentrations…

11. Exceptions to Insolubility Most, but not all exceptions are found on the other side of the table (Group 1, ammonium ion)

12. Fracking

14. Fracking Fluid Millions of gallons of water and other chemicals are used to liberate methane from shale deposits. Fracking fluid is basic: it contains lots of OH - ions. Much of it has to be removed from the ground for reprocessing: it is 1000x more radioactive when it comes back up!

17. WHY?!? Exposing rock to strong base liberates radium – a very radioactive decay product of uranium. Ra(OH) 2 is much more soluble than Mg(OH) 2 (milk of magnesia)

18. Using Table F, determine which of the following are soluble/insoluble. CompoundSoluble Insoluble Na 2 SO 4 Mg 3 ( PO 4 ) 2 BaSO 4 SrSO 4 KClO 3 K2SK2S NH 4 NO 3 Ba(OH) 2 LiOH (NH 4 ) 2 CrO 4

19. CompoundSoluble/Insoluble Pb(ClO 4 ) 2 ZnS KC 2 H 3 O 2 (NH 4 ) 2 SO 4 K 2 CrO 4 Co(OH) 2 HgCO 3 Hg 2 Cl 2 NiCl 2 NaC 2 H 3 O 2

20. According to Reference Table F, which of the following compounds will form a saturated solution that is most dilute? ammonium chloride calcium carbonate potassium iodide sodium nitrate

21. Solubility Rules and Double Replacement Reactions Assign oxidation numbers to each element, and find the charges of each polyatomic ion using table E. Write the ion combinations and their charges on the product side. Criss cross each charge to get the formula. Balance each equation. Determine the preciptate using Table F. For those that should have precipitates, indicate which product is the precipitate by using a down arrow↓ or (s).

22. +2 -1 +1 -2 +2 -2 +1 -1 ZnCl 2 + Na 2 CO 3 → ZnCO 3 + NaCl CaCl 2 + Pb(NO 3 ) 2 → FeBr 2 + K 2 CrO 4 →

23. AlCl 3 + Ba(OH) 2 → ZnCl 2 + AgNO 3 → KI + Pb(NO 3 ) 2 → KI + Ba(OH) 2 →

24. Lab 13: Double Replacement Reactions and Solubility Rules Objective To study various double replacement reactions, identifying precipitates and spectator ions. Introduction Combining solutions of soluble ionic compounds often results in a product that precipitates. The insoluble product can be determined by using the rules of solubility shown on Table F of your reference tables.

25. Materials Well plate, dropper bottles containing the following: Sodium chloride, ammonium hydroxide, silver nitrate, sodium carbonate, potassium iodide, copper sulfate, cobalt chloride, lead nitrate; chemistry reference table.

26. Procedure Take a well plate and clean it completely. Place 1 drop of the following in a well. Be sure not to mix any of the chemicals accidentally! To avoid this, do not put the dropper directly on the plastic well – drop the drop a half an inch above the well plate!

27. No.ReactantsReactant 1 Formula Reactant 2 Formula Precipitate Color 1Sodium chloride and silver nitrate 2cobalt (II) chloride and silver nitrate 3sodium carbonate and silver nitrate 4potassium iodide and silver nitrate 5 lead (II) nitrate and ammonium hydroxide 6lead (II) nitrate and sodium chloride 7lead (II) nitrate and potassium iodide 8copper (II) sulfate and sodium carbonate 9copper (II) sulfate and ammonium hydroxide 10copper (II) sulfate and lead nitrate 11cobalt (II) chloride and sodium carbonate 12cobalt (II) chloride and ammonium hydroxide

28. Try some new combinations! Choose a combination of any two solutions that are not combined above. Describe what happens (or doesn’t happen) below: There is one other combination that produces a precipitate. Use your reference tables to figure out what it is for extra credit!

29. Results Writing a balanced chemical equation is not that easy. There are a number of simple steps that you must take. The key is figuring out the formulas of the products. Write the oxidation numbers of the elements over the chemical formulas of the reactants. Do the same for the charges on the polyatomic ions – use tables E and/or for this. Draw the arrow to the products, and switch partners – the metal ion of one compound goes with the non metal ion or polyatomic ion of the other compound. Write the charges on the product side – they do not change in double replacement reactions! Using the criss-cross method, determine the chemical formula of the products. Use parenthesis when you have more than one polyatomic ion. Check with the teacher if you’re not sure. Balance the equations using coefficients. Determine the identity of the precipitate using Table F. Write a downward arrow ↓ or (s) after the chemical formula of the precipitate.

30. Reactant NamesSodium chloride and silver nitrate Oxidation #s Chemical Equation Reactant Namescobalt chloride and silver nitrate Oxidation #s Chemical Equation Reactant Namessodium carbonate and silver nitrate Oxidation #s Chemical Equation

31. Reactant Namespotassium iodide and silver nitrate Oxidation #s Chemical Equation Reactant Nameslead nitrate and ammonium hydroxide Oxidation #s Chemical Equation Reactant Nameslead nitrate and sodium chloride Oxidation #s Chemical Equation Reactant Nameslead nitrate and potassium iodide Oxidation #s Chemical Equation Reactant Namescopper sulfate and sodium carbonate Oxidation #s Chemical Equation

32. Reactant Namescopper sulfate and ammonium hydroxide Oxidation #s Chemical Equation Reactant Namescopper sulfate and lead nitrate Oxidation #s Chemical Equation Reactant Namescobalt chloride and sodium carbonate Oxidation #s Chemical Equation Reactant Namescobalt chloride and ammonium hydroxide Oxidation #s Chemical Equation

33. Questions What are some possible sources of error in this experiment? One student looked at a precipitate and identified the precipitate as pink. Her lab partner insisted it was white. Sometimes it isn’t easy to tell what the color of a precipitate is. Devise a method to find out! Why do some combinations of solutions not produce precipitates?

34. Summary of Results Using Table F, identify the precipitate and give two justifications based on the table. The first one is done for you. No.PrecipitateJustification #1 (soluble product)Justification #2 (insoluble product) 1Silver ChlorideSodium is always soluble (Group 1); Nitrate is always soluble Silver is insoluble with halides 2 3