1. Redox Reactions Chapter 18 + O 2 

2. Oxidation-Reduction (Redox) Reactions “redox” reactions: rxns in which electrons are transferred from one species to another oxidation & reduction always occur simultaneously we use OXIDATION NUMBERS to keep track of electron transfers

3. Rules for Assigning Oxidation Numbers: 1) the ox. state of any free (uncombined) element is zero. Ex: Na, S, O 2, H 2, Cl 2, O 3

4. Rules for Assigning Oxidation Numbers: 2) The ox. state of an element in a simple ion is the charge of the ion. Mg 2+  oxidation of Mg is +2

5. Rules for Assigning Oxidation Numbers: 3) the ox. # for hydrogen is +1 (unless combined with a metal, then it has an ox. # of –1) Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)

6. Rules for Assigning Oxidation Numbers: 4) the ox. # of fluorine is always –1.

7. Rules for Assigning Oxidation Numbers: 5) the ox. # of oxygen is usually –2. Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide: H2O2H2O2

8. Rules for Assigning Oxidation Numbers: 6) in any neutral compound, the sum of the oxidation #’s = zero.

9. Rules for Assigning Oxidation Numbers: 7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.

10. Rules for Assigning Oxidation Numbers: **use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:

11. Examples: Assign oxidation #’s to each element: a) NaNO 3

12. Examples: Assign oxidation #’s to each element: b) SO 3 2-

13. Examples: Assign oxidation #’s to each element: c) HCO 3 -

14. Examples: Assign oxidation #’s to each element: d) H 3 PO 4

15. Examples: Assign oxidation #’s to each element: e) Cr 2 O 7 2-

16. Examples: Assign oxidation #’s to each element: f) K 2 Sn(OH) 6

17. Definitions Oxidation: the process of losing electrons (ox # increases) Reduction: the process of gaining electrons (ox # decreases) Oxidizing agents: species that cause oxidation (they are reduced) Reducing agents: species that cause reduction (they are oxidized)

18. To help you remember… OIL RIG Oxidation Is Loss Reduction Is Gain

19. Are all rxns REDOX rxns? a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

20. Examples MgCO 3  MgO +CO 2

21. Examples Zn + CuSO 4  ZnSO 4 + Cu

22. Examples NaCl + AgNO 3  AgCl + NaNO 3

23. Examples CO 2 + H 2 O  C 6 H 12 O 6 + O 2

24. Balancing Redox Equations

25. In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #) There are 2 methods for balancing redox equations: 1. Change in Oxidation-Number Method 2. The Half-Reaction Method

26. 1. Change in Oxidation-Number Method: based on equal total increases and decreases in oxidation #’s Steps: 1. Write equation and assign oxidation #’s. 2. Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each. 3. Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket. 4. Choose coefficients that make the total increase in ox. # = the total decrease in ox. #. 5. Balance the remaining elements by inspection.

27. Example S + HNO 3  SO 2 + NO + H 2 O

28. If needed, reactions that take place in acidic or basic solutions can be balanced as follows: Acidic:Basic: add H 2 O to the side needing oxygen balance as if in acidic sol’n then add H + to balance the hydrogen add enough OH - to both sides to cancel out each H + (making H 2 O) & then cancel out water as appropriate

29. Example: Balance the following equation, assuming it takes place in acidic solution. ClO 4 - +I -  Cl - +I 2

30. Example: Balance the following equation, assuming it takes place in basic solution. ClO 4 - +I -  Cl - +I 2

31. 2. The Half-Reaction Method: separate and balance the oxidation and reduction half-reactions. Steps: 1. Write equation and assign oxidation #’s. 2. Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each. 3. Construct unbalanced oxidation and reduction half reactions. 4. Balance the elements and the charges (by adding electrons as reactants or products) in each half-reaction. 5. Balance the electron transfer by multiplying the balanced half- reaction by appropriate integers. 6. Add the resulting half-reaction and eliminate any common terms to obtain the balanced equation.

32. Example: Balance the following using the half-reaction method: HNO 3 +H 2 S  NO +S+H 2 O

33. If needed, reactions that take place in acidic solutions can be balanced as follows: Acidic: 1. Write separate eq’ns for the oxidation & reduction half-rxns 2. For each half-rxn: a) Balance all the elements except H and O b) Balance O using H 2 O c) Balance H using H + d) Balance charge using elections 3. If necessary, multiply one or both balanced half-rxns by an integer to equalize the number of electrons transferred in the two half-rxns. 4. Add the half-reactions and cancel the identical species (those appearing in reactants and products) 5. Check that the elements and charges are balanced

34. If needed, reactions that take place in basic solutions can be balanced as follows: Basic: 1. Balance as if in acidic sol’n (follow ALL steps for acidic redox) 2. Add enough OH - to both sides to cancel out each H + (making H 2 O) & then cancel out water as appropriate 3. Check that the elements and charges are balanced

35. HOMEWORK: Balance the following using the half-rxn method… In acidic sol’n: a) Cu + NO 3 -  Cu 2+ + NO b) Cr 2 O 7 2- + Cl -  Cr 3+ + Cl 2 c) Pb + PbO 2 + H 2 SO 4  PbSO 4 In basic sol’n: a) Al + MnO 4 -  MnO 2 + Al(OH) 4 - b) Cl 2  Cl - + OCl - c) NO 2 - + Al  NH 3 + AlO 2 -