1. Lecture 254/4/05 Seminar today

2. Standard Reduction Potentials 1. Each half-reaction is written as a reduction 2. Each half-reaction can occur in either direction 3. The more positive E° red, the more easily the substance on the left side of the half-reaction can be reduced 4. The more negative the E° red, the less likely the reaction will occur as reduction, and the more likely oxidation will occur 5. Under standard conditions, any species on the left side of a half-reaction will oxidize any species on the right that is farther down the table 6. Electrode potentials depend on the identity and concentration of the reactants and products, not the total quantity of each

3. Each half-reaction is written as a reduction

4. Each half-reaction can occur in either direction

5. F 2 + 2e -  2F - reduction 2F -  F 2 + 2e - oxidation Each half-reaction can occur in either direction

6. The more positive E°red, the more easily the substance on the left side of the half-reaction can be reduced

7. The more negative the E°red, the less likely the reaction will occur as reduction, and the more likely oxidation will occur

8. Under standard conditions, any species on the left side of a half-reaction will oxidize any species on the right that is farther down the table

9. Examples: What are the reactions for: F 2 and Cl - F 2 and Cl 2 F - and Cl 2 F - and Cl - Need to consider 2 questions: 1)Are both reduction and oxidation reactions possible? 2)Does the reduction reaction have a higher reduction potential than the oxidation reaction?

10. Practice E° red (V) Cl 2 + 2e -  2Cl - +1.36 I 2 + 2e -  I - +0.535 Pb 2+ + 2e -  Pb (s)+0.126 V 2+ + 2e -  V (s)-1.18 a)What is the weakest oxidation agent? b)What is the strongest oxidation agent? c)What is the weakest reduction agent? d)What is the strongest reducing agent? e)Will Pb(s) reduce V 2+ ? f)Will I 2 oxidize Cl - ? g)Which molecule(s) can be reduced by Pb(s)

11. Gibb’s free Energy (∆G) Measure of the spontaneity of the reaction at constant temperature and pressure If negative, reaction is spontaneous Units are J/mol ∆G = -nFE n = # of electrons transferred F = Faraday’s constant (96,500 J/V-mol) E = reduction potential ∆G° = -nFE°

12. Nernst Equation Remember E° is at standard conditions AND E is at all other conditions As the reactants are consumed and the products formed, E eventually drops to zero This is why batteries don’t last forever Use Nernst equation to calculate E cell generated under non- standard conditions

13. Nernst Equation Calculate the voltage delivered by a voltaic cell using the following reaction if all dissolved species are 0.015M. Fe 2+ + H 2 O 2  Fe 3+ + H 2 O Want this Calculate using the standard reduction potential table Get from half reactions Get from reaction quotient equation and concentrations

14. Calculate the voltage delivered by a voltaic cell using the following reaction if all dissolved species are 0.015M. Fe 2+ + H 2 O 2  Fe 3+ + H 2 O Get from half reactions Oxidation: Fe 2+  Fe 3+ + e - Reduction: 2e - + 2H + + H 2 O 2  2H 2 O Net: 2Fe 2+ + 2H + + H 2 O 2  2Fe 3+ + 2H 2 O n = 2

15. Calculate the voltage delivered by a voltaic cell using the following reaction if all dissolved species are 0.015M. 2Fe 2+ + 2H + + H 2 O 2  2Fe 3+ + 2H 2 O Calculate using the standard reduction potential table E° cell = E° cathode - E° anode = 1.77 – 0.771 = 1.0 V

16. Get from reaction quotient equation and concentrations Calculate the voltage delivered by a voltaic cell using the following reaction if all dissolved species are 0.015M. 2Fe 2+ + 2H + + H 2 O 2  2Fe 3+ + 2H 2 O

17. Calculate the voltage delivered by a voltaic cell using the following reaction if all dissolved species are 0.015M. 2Fe 2+ + 2H + + H 2 O 2  2Fe 3+ + 2H 2 O