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CHAPTER 16 (pages 776-792) 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

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

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OBSERVED REDOX PROCESSES 3

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GALVANIC CELLS 4

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INERT ELECTRODES 6

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STANDARD REDUCTION POTENTIALS 7

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MEASURING STANDARD POTENTIALS 9

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

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

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ox: Fe( s ) Fe 2+ ( aq ) + 2 e − E = +0.45 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 = +0.32 V

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ELECTROMOTIVE POTENTIAL 13

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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

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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 = -0.13 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 10 -4 M.

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Williams, spring 2009 stop here

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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 = -0.13 V Calculate the E o cell (potential at standard conditions)

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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 = -0.13 V Calculate G o.

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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 = -0.13 V Calculate

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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 = -0.13 V Calculate the E cell if [Ag + ] = 2.0 M and Pb 2+ ] = 1.0 x 10 -4 M.

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

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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

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

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Dry Cell Batteries

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Lead – Acid Storage Battery

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Biological Electrochemistry

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Lithium Ion Battery

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Electrochemistry. #13 Electrochemistry and the Nernst Equation Goals: To determine reduction potentials of metals To measure the effect of concentration.

Electrochemistry. #13 Electrochemistry and the Nernst Equation Goals: To determine reduction potentials of metals To measure the effect of concentration.

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