CHAPTER 16 (pages ) 1.Oxidation and Reduction 2.Galvanic Cells, Half Reactions (E° anode & E° cathode ) 3.Standard Reduction Potential (E°) 4.Nernst Equation, and the dependence of Potential on Concentration 5.Relationship between Equilibrium Constant and Standard Potential 6.Driving Force, ΔG and ε 1

REDOX REACTIONS MnO HBr  MnBr 2 + Br H 2 O 3 H 2 S + 2 NO 3 – + 2 H +  S + 2 NO + 4 H 2 O 2

OBSERVED REDOX PROCESSES 3

GALVANIC CELLS 4

INERT ELECTRODES 6

STANDARD REDUCTION POTENTIALS 7

8

MEASURING STANDARD POTENTIALS 9

CALCULATING STANDARD CELL POTENTIAL Al (s) + NO 3 − (aq) + 4 H + (aq)  Al 3+ (aq) + NO (g) + 2 H 2 O (l) 10

ADDITIONAL EXAMPLE Fe (s) + Mg 2+ (aq)  Fe 2+ (aq) + Mg (s) 11

ox: Fe( s )  Fe 2+ ( aq ) + 2 e − E  = V red: Pb 2+ ( aq ) + 2 e −  Pb( s ) E  = −0.13 V tot: Pb 2+ ( aq ) + Fe( s )  Fe 2+ ( aq ) + Pb( s ) E  = V

ELECTROMOTIVE POTENTIAL 13

E° CELL, Δ G° AND K Under standard state conditions, a reaction will spontaneously proceeds in the forward direction if: –Δ G° < 1 (negative) – E° > 1 (positive) – K > 1

Design a voltaic cell with the following half cells and complete the calculations: Ag + (aq) + 1e -  Ag (s) E o = 0.80 V Pb 2+ (aq) + 2e -  Pb (s) E o = V a.Calculate the E o cell (potential at standard conditions) b.Calculate  G o. c.Calculate  d.Calculate the E cell if [Ag + ] = 2.0 M and [Pb 2+ ] = 1.0 x M.

Williams, spring 2009 stop here

Design a voltaic cell with the following half cells and complete the calculations: Ag + (aq) + 1e -  Ag (s) E o = 0.80 V Pb 2+ (aq) + 2e -  Pb (s) E o = V Calculate the E o cell (potential at standard conditions)

Design a voltaic cell with the following half cells and complete the calculations: Ag + (aq) + 1e -  Ag (s) E o = 0.80 V Pb 2+ (aq) + 2e -  Pb (s) E o = V Calculate  G o.

Design a voltaic cell with the following half cells and complete the calculations: Ag + (aq) + 1e -  Ag (s) E o = 0.80 V Pb 2+ (aq) + 2e -  Pb (s) E o = V Calculate 

Design a voltaic cell with the following half cells and complete the calculations: Ag + (aq) + 1e -  Ag (s) E o = 0.80 V Pb 2+ (aq) + 2e -  Pb (s) E o = V Calculate the E cell if [Ag + ] = 2.0 M and Pb 2+ ] = 1.0 x M.

OBJECTIVE 11.4: PROVIDE A THOROUGH OVERVIEW OF APPLICATIONS OF ELECTROCHEMICAL CELLS INCLUDING FUEL CELLS, CORROSION, AND OTHER TOPICS AS TIME PERMITS. 23

CORROSION corrosion is the spontaneous oxidation of a metal by chemicals in the environment since many materials we use are active metals, corrosion can be a very big problem

RUSTING rust is hydrated iron(III) oxide moisture must be present electrolytes promote rusting acids promote rusting – lower pH = lower E° red

Dry Cell Batteries

Lead – Acid Storage Battery

Biological Electrochemistry

Lithium Ion Battery