1. 1 Types of Reactions  Precipitation reactions l When aqueous solutions of ionic compounds are poured together a solid forms. l A solid that forms from mixed solutions is a precipitate l If you’re not a part of the solution, your part of the precipitate

2. 2 Precipitation reactions NaOH(aq) + FeCl 3 (aq)   NaCl(aq) + Fe(OH) 3 (s) l is really Na + (aq)+OH - (aq) + Fe +3 + Cl - (aq)   Na + (aq) + Cl - (aq) + Fe(OH) 3 (s) l So all that really happens is OH - (aq) + Fe +3  Fe(OH) 3 (s) l Double replacement reaction

3. 3 Precipitation reaction l We can predict the products l Can only be certain by experimenting l The anion and cation switch partners AgNO 3 (aq) + KCl(aq)  Zn(NO 3 ) 2 (aq) + BaCr 2 O 7 (aq)  CdCl 2 (aq) + Na 2 S(aq) 

4. 4 Precipitations Reactions l Only happen if one of the products is insoluble l Otherwise all the ions stay in solution- nothing has happened. l Need to memorize the rules for solubility (pg 145)

5. 5 Solubility Rules  All nitrates are soluble  Alkali metals ions and NH 4 + ions are soluble  Halides are soluble except Ag +, Pb +2, and Hg 2 +2  Most sulfates are soluble, except Pb +2, Ba +2, Hg +2,and Ca +2

6. 6 Solubility Rules  Most hydroxides are slightly soluble (insoluble) except NaOH and KOH  Sulfides, carbonates, chromates, and phosphates are insoluble  Lower number rules supersede so Na 2 S is soluble

7. 7 Three Types of Equations l Molecular Equation- written as whole formulas, not the ions. K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq)  l Complete Ionic equation show dissolved electrolytes as the ions. 2K + + CrO 4 -2 + Ba +2 + 2 NO 3 -  BaCrO 4 (s) + 2K + + 2 NO 3 - l Spectator ions are those that don’t react.

8. 8 Three Type of Equations l Net Ionic equations show only those ions that react, not the spectator ions Ba +2 + CrO 4 -2  BaCrO 4 (s) l Write the three types of equations for the reactions when these solutions are mixed. l iron (III) sulfate and potassium sulfide Lead (II) nitrate and sulfuric acid.

9. 9 Stoichiometry of Precipitation l Exactly the same, except you may have to figure out what the pieces are. l What mass of solid is formed when 100.00 mL of 0.100 M Barium chloride is mixed with 100.00 mL of 0.100 M sodium hydroxide? l What volume of 0.204 M HCl is needed to precipitate the silver from 50.ml of 0.0500 M silver nitrate solution ?

10. 10 Types of Reactions  Acid-Base l For our purposes an acid is a proton donor. l a base is a proton acceptor usually OH - l What is the net ionic equation for the reaction of HCl(aq) and KOH(aq)? Acid + Base  salt + water H + + OH -  H 2 O

11. 11 Acid - Base Reactions l Often called a neutralization reaction Because the acid neutralizes the base. l Often titrate to determine concentrations. l Solution of known concentration (titrant), l is added to the unknown (analyte), l until the equivalence point is reached where enough titrant has been added to neutralize it.

12. 12 Titration l Where the indicator changes color is the endpoint. l Not always at the equivalence point. l A 50.00 mL sample of aqueous Ca(OH) 2 requires 34.66 mL of 0.0980 M Nitric acid for neutralization. What is [Ca(OH) 2 ]? l # of H + x M A x V A = # of OH - x M B x V B

13. 13 Acid-Base Reaction l 75 mL of 0.25M HCl is mixed with 225 mL of 0.055 M Ba(OH) 2. What is the concentration of the excess H + or OH - ?

14. 14 Types of Reaction  Oxidation-Reduction called Redox l Ionic compounds are formed through the transfer of electrons. l An Oxidation-reduction reaction involves the transfer of electrons. l We need a way of keeping track.

15. 15 Oxidation States l A way of keeping track of the electrons. l Not necessarily true of what is in nature, but it works. l need the rules for assigning (memorize).  The oxidation state of elements in their standard states is zero.  Oxidation state for monoatomic ions are the same as their charge.

16. 16 Oxidation states  Oxygen is assigned an oxidation state of - 2 in its covalent compounds except as a peroxide.  In compounds with nonmetals hydrogen is assigned the oxidation state +1.  In its compounds fluorine is always –1.  The sum of the oxidation states must be zero in compounds or equal the charge of the ion.

17. 17 Oxidation States l Assign the oxidation states to each element in the following. l CO 2 l NO 3 - l H 2 SO 4 l Fe 2 O 3 l Fe 3 O 4

18. 18 Oxidation-Reduction l Transfer electrons, so the oxidation states change. Na + 2Cl 2  2NaCl CH 4 + 2O 2  CO 2 + 2H 2 O l Oxidation is the loss of electrons. l Reduction is the gain of electrons. l OIL RIG l LEO GER

19. 19 Oxidation-Reduction l Oxidation means an increase in oxidation state - lose electrons. l Reduction means a decrease in oxidation state - gain electrons. l The substance that is oxidized is called the reducing agent. l The substance that is reduced is called the oxidizing agent.

20. 20 Redox Reactions

21. 21 Agents l Oxidizing agent gets reduced. l Gains electrons. l More negative oxidation state. l Reducing agent gets oxidized. l Loses electrons. l More positive oxidation state.

22. 22 Identify the l Oxidizing agent l Reducing agent l Substance oxidized l Substance reduced l in the following reactions Fe (s) + O 2 (g)  Fe 2 O 3 (s) Fe 2 O 3 (s)+ 3 CO(g)  2 Fe(l) + 3 CO 2 (g) SO 3 2- + H + + MnO 4 -  SO 4 2- + H 2 O + Mn 2+

23. 23 Half-Reactions l All redox reactions can be thought of as happening in two halves. l One produces electrons - Oxidation half. l The other requires electrons - Reduction half. l Write the half reactions for the following. Na + Cl 2  Na + + Cl - SO 3 2- + H + + MnO 4 -  SO 4 2- + H 2 O + Mn +2

24. 24 Balancing Redox Equations l In aqueous solutions the key is the number of electrons produced must be the same as those required. l For reactions in acidic solution an 8 step procedure.  Write separate half reactions  For each half reaction balance all reactants except H and O  Balance O using H 2 O

25. 25 Acidic Solution  Balance H using H +  Balance charge using e -  Multiply equations to make electrons equal  Add equations and cancel identical species  Check that charges and elements are balanced.

26. 26 Practice l The following reactions occur in aqueous solution. Balance them MnO 4 - + Fe +2  Mn +2 + Fe +3 Cu + NO 3 -  Cu +2 + NO(g) Pb + PbO 2 + SO 4 -2  PbSO 4 Mn +2 + NaBiO 3  Bi +3 + MnO 4 -

27. 27 Now for a tough one Fe(CN) 6 -4 + MnO 4 -   Mn +2 + Fe +3 + CO 2 + NO 3 -

28. 28 Basic Solution l Do everything you would with acid, but add one more step. l Add enough OH - to both sides to neutralize the H + l Makes water CrI 3 + Cl 2  CrO 4 2- + IO 4 - + Cl - Fe(OH) 2 + H 2 O 2  Fe(OH) - Cr(OH) 3 + OCl - + OH -   CrO 4 2- + Cl - + H 2 O

29. 29 Redox Titrations l Same as any other titration. l the permanganate ion is used often because it is its own indicator. MnO 4 - is purple, Mn +2 is colorless. When reaction solution remains clear, MnO 4 - is gone. l Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect.

30. 30 Example l The iron content of iron ore can be determined by titration with standard KMnO 4 solution. The iron ore is dissolved in excess HCl, and the iron reduced to Fe +2 ions. This solution is then titrated with KMnO 4 solution, producing Fe +3 and Mn +2 ions in acidic solution. If it requires 41.95 mL of 0.205 M KMnO 4 to titrate a solution made with 0.6128 g of iron ore, what percent of the ore was iron?

31. 31 Extra Credit l Nuclear Power l Write a paper that describes l 1. How does it work? l 2. What are the advantages? l 3. What are the disadvantages? l 4. Using your information to support your conclusion, answer the question, “What role should nuclear power play in future energy generation for the United States?” l 5-7 pages l Researched using MLA style with in text citations. l Due Oct. 26- no exceptions.