1. C H E M I S T R Y Chapter 4 Reactions in Aqueous Solution

2. Oxidation-Reduction (Redox) Reactions 2Fe 2 O 3 (s)4Fe(s) + 3O 2 (g) Rusting of iron: an oxidation of Fe 4Fe(s) + 3CO 2 (g)2Fe 2 O 3 (s) + 3C(s) Manufacture of iron: a reduction of Fe

3. Oxidation-Reduction (Redox) Reactions 1.An atom in its elemental state has an oxidation number of 0. Rules for Assigning Oxidation Numbers Oxidation Number (State): A value which indicates whether an atom is neutral, electron-rich, or electron-poor. NaH2H2 Br 2 SNe Oxidation number 0

4. Oxidation-Reduction (Redox) Reactions 2.A monatomic ion has an oxidation number identical to its charge. Na 1+ +1 Ca 2+ +2 Al 3+ +3 Cl 1- O 2- -2

5. Oxidation-Reduction (Redox) Reactions b)Oxygen usually has an oxidation number of -2. HO 1- -2+1 HHCa +2 3.An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. a)Hydrogen can be either +1 or -1. OHO +1 HHO +1 -2 H +1

6. Oxidation-Reduction (Redox) Reactions HCl +1 c)Halogens usually have an oxidation number of -1. 3. Cl O +1 -2

7. Oxidation-Reduction (Redox) Reactions Cr 2 O 7 2- -2+1x 4.The sum of the oxidation numbers is 0 for a neutral compound and is equal to the net charge for a polyatomic ion. x = +6 2(x) + 7(-2) = -2 (net charge) H2SO3H2SO3 x-2 x = +4 2(+1) + x + 3(-2) = 0 (net charge)

8. Example Determine the oxidation number for each atom in the following compounds/molecules ◦ CO 2 ◦ CCl 4 ◦ CoSO 4 ◦ K 2 O 2

9. Identifying Redox Reactions Oxidation: losing one or more electrons increasing in oxidation number  Reducing agent 2Fe 2 3O2(g)3O2(g)+4Fe(s)O3O3 (s)(s) +30 oxidation reduction -2 0 Reduction: gaining one or more electron decreasing in oxidation number  Oxidizing agent

10. Identifying Redox Reactions Reducing Agent Causes reduction Loses one or more electrons Undergoes oxidation Oxidation number of atom increases Oxidizing Agent Causes oxidation Gains one or more electrons Undergoes reduction Oxidation number of atom decreases

11. Example Identify each of half reactions as 1) oxidation or 2) reduction. __A. Sn(s) Sn 4+ (aq) + 4e − __B. Fe 3+ (aq) + 1e − Fe 2+ (aq) __C. Cl 2 (g) + 2e − 2Cl - (aq) 11

12. Writing Oxidation and Reduction Reactions Write the separate half oxidation and reduction reactions for the following equation. 2Cs(s) + F 2 (g) 2CsF(s) 3 Na(l) + AlCl 3 (l)  3 NaCl(l) + Al(l) 12

13. The Activity Series of the Elements 2Ag 1+ (aq) + Cu(s)2Ag(s) + Cu 2+ (g) Cu 2+ (aq) + 2Ag(s)Cu(s) + 2Ag 1+ (g) Which one of these reactions will occur?

14. The Activity Series of the Elements Elements that are higher up in the table are more likely to be oxidized. Thus, any element higher in the activity series will reduce the ion of any element lower in the activity series.

15. The Activity Series of the Elements 2Ag 1+ (aq) + Cu(s)2Ag(s) + Cu 2+ (g) Cu 2+ (aq) + 2Ag(s)Cu(s) + 2Ag 1+ (g) Which one of these reactions will occur?

16. Example Predict whether the following redox reactions will occurred or not. If so, predict the products ◦ Zn(s) + FeCl 2 (aq)  ◦ Ni(s) + Mg(NO 3 ) 2 (aq) 

17. Redox Titrations If the unknown concentration is the potassium permanganate solution, MnO 4 1-, it can be slowly added to a known amount of oxalic acid, H 2 C 2 O 4, until a faint purple color persists. Titration: A procedure for determining the concentration of a solution by allowing a carefully measured volume to react with a solution of another substance (the standard solution) whose concentration is known. 5H 2 C 2 O 4 (aq) + 2MnO 4 1- (aq) + 6H 1+ (aq) 10CO 2 (g) + 2Mn 2+ (aq) + 8H 2 O(l)

18. Redox Titrations 5H 2 C 2 O 4 (aq) + 2MnO 4 1- (aq) + 6H 1+ (aq) 10CO 2 (g) + 2Mn 2+ (aq) + 8H 2 O(l) A solution is prepared with 0.2585 g of oxalic acid, H 2 C 2 O 4. 22.35 mL of an unknown solution of potassium permanganate are needed to titrate the solution. What is the concentration of the potassium permanganate solution?

19. Calculation Set up Moles of H 2 C 2 O 4 Mass of H 2 C 2 O 4 Moles of KMnO 4 Molarity of KMnO 4 Mole Ratio Molarity of KMnO 4 Molar Mass of H 2 C 2 O 4

20. Example A 0.0484M standard solution of potassium permanganate was titrated against 25.00mL of an iron (II) sulfate solution. The equivalence point, as indicated by a faint pink color, was reached when 15.50mL of potassium permanganate solution had been added. Calculate the concentration of the iron (II) sulfate solution