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balancing-redox-equations-handout

balancing-redox-equations-handout - Chapter 4.11 Balancing...

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OFB Chapter 12 1 11/13/09 Chapter 4.11 Balancing Redox Equations This section is intended to be a supplemental resources for students. A stepwise process to balance oxidation-reduction (redox) reactions is introduced. Student’s should use these slides to practice balancing redox reactions. The number of electrons transferred will be used for the Electrochemical cell reactions in Chapter 16
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OFB Chapter 12 2 11/13/09 This section 4.11deals with Oxidation, Reduction, and electron transfer Balancing Redox Reactions All reactions take place in either Acid (H + ) or Basic (OH - ) conditions Very important to remember the problem you are working. Is it acidic or basic conditions? We will be adding H 2 O and H + for acid reactions or H 2 O and OH- for basic reaction in order to properly balance the reactions Practice, Practice, Practice
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OFB Chapter 12 3 11/13/09 Gains Electrons Decrease Substance Reduced Loses Electrons Increase Substance Oxidized Supplies Electrons Increase Reducing Agent, does the reducing Picks Up electrons Decrease Oxidizing Agent, does the oxidizing Gain of Electrons Decrease Reduction Loss of Electrons Increase Oxidation Electron Change Oxidation Number Change Term The Oxidation States Method Cl 2 Na NaCl +
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OFB Chapter 12 4 11/13/09 Step 1 Write two unbalanced half- equations , one for the species that is oxidized and its product and one for the species that is reduced and its product Step 2 Insert coefficients to make the numbers of atoms of all elements except oxygen and hydrogen equal on the two sides of each half-equation. Step 3 Balance oxygen by adding H 2 O to the side deficient in O in each half- equation Step 4 Balance hydrogen . For half-reaction in acidic solution, add H + on to the side deficient in hydrogen. For a half-reaction in basic solution, add H 2 O to the side that is deficient in hydrogen and an equal amount of OH - to the other side
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OFB Chapter 12 5 11/13/09 Step 5 Balance charge by inserting e - (electrons) as a reactant or product in each half-reaction. Step 6 Multiply the two half-equations by numbers chosen to make the number of electrons given off by the oxidation equal to the number taken up by the reduction. Then add the two half-equations and cancel out the electrons. If H + ion, OH - ion, or H 2 O appears on both sides of the final equation, cancel out the duplication Step 7 Check for balance
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OFB Chapter 12 6 11/13/09 Balancing Redox Reactions Dithionate ion reacting with chlorous acid in aqueous acidic conditions S 2 O 6 2- (aq) +HClO 2 (aq) → SO 4 2- (aq) + Cl 2 (g) Will have: Oxidation half reaction Reduction half reaction
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OFB Chapter 12 7 11/13/09 Step 1 Write two unbalanced half-equations , one for the species that is oxidized and its product and one for the species that is reduced and its product S 2 O 6 2- → SO 4 2- Dithionate (we will see eventually that this is the oxidation reaction) HClO 2 → Cl 2 Chlorous acid (we will see eventually that this is the reduction reaction) S 2 O 6 2- (aq) +HClO 2 (aq) → SO 4 2- (aq) + Cl 2 (g)
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