1. Balancing Redox Equations

2. 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

3. 1. Change in Oxidation-Number Method: based on equal total increases and decreases in oxidation #’s Steps: page 645 in textbook 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.

4. Example S + HNO 3  SO 2 + NO + H 2 O Step 1: Oxidation numbers 0+1 +5 -2 +4-2 +2 -2+1 -2 Step 2: Which was oxidized? S. By how much? 0 to +4 = change of +4 Which was reduced?N. By how much? +5 to +2 = change of -3 +4 -3 3 4 3 4 2 So, 3 coefficient x +4 = +12 So, 4 coefficient x -3 = -12

5. 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

6. Example: Balance the following equation, assuming it takes place in acidic solution. Page 648 in textbook ClO 4 - +I -  Cl - +I 2 Step 1: Oxidation numbers +7 -2 0 Step 2: Which was oxidized? Iodine, -1 to 0 = +1 Which was reduced?Chlorine, +7 to -1 = -8 +1 -8 8 4 Step 5: Balance acid soln with water… + H 2 O +8 H + 4

7. Example: Balance the following equation, assuming it takes place in basic solution. ClO 4 - +I -  Cl - +I 2 Cancel H + by adding OH - to both sides. H + + OH - = H 2 O Cancel H 2 O

8. 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.

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

10. Example: Balance the following using the half-reaction method: HNO 3 +H 2 S  NO +S+H 2 O Step 1: oxidation numbers +1+5 +2 +1 -2 0 Step 2: Which was oxidized? S. -2 to 0 = +2 Which was reduced? N. +5 to +2 = -3 Step 3: unbalanced half-rxns S 2-  S N 5+  N 2+ Step 4: balance the half – rxns by adding electrons + 2 e- + 3 e- Step 5: balance electron transfer by multiplying by appropriate integers x3 x2 3S 2-  3S + 6e- 2 N 5+ + 6e-  2 N 2+ Step 6: Add half-rxns and cancel any common terms to obtain a balanced eq. ---------------------------------- 3S 2- + 2N 5+  3S + 2N 2+ Now, balance the eq. w/coefficients 2 3 234

11. 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

12. 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

13. 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 -