# Balancing Redox Equations. In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained.

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Balancing Redox Equations

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

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.

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

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

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

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.

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

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