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Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 1 of 54 CHEMISTRY Ninth Edition GENERAL Principles and Modern Applications Petrucci Harwood Herring.

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Presentation on theme: "Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 1 of 54 CHEMISTRY Ninth Edition GENERAL Principles and Modern Applications Petrucci Harwood Herring."— Presentation transcript:

1 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 1 of 54 CHEMISTRY Ninth Edition GENERAL Principles and Modern Applications Petrucci Harwood Herring Madura Chapter 20: Electrochemistry

2 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 2 of 54 Contents 20-1Electrode Potentials and Their Measurement 20-2Standard Electrode Potentials 20-3E cell, ΔG, and K eq 20-4 E cell as a Function of Concentration 20-5Batteries: Producing Electricity Through Chemical Reactions 20-6Corrosion: Unwanted Voltaic Cells 20-7Electrolysis: Causing Non-spontaneous Reactions to Occur 20-8Industrial Electrolysis Processes

3 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 3 of Electrode Potentials and Their Measurement Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2 Ag(s) Cu(s) + Zn 2+ (aq) No reaction

4 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 4 of 54 An Electrochemical Half Cell Anode Cathode

5 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 5 of 54 An Electrochemical Cell

6 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 6 of 54 Terminology  Electromotive force, E cell.  The cell voltage or cell potential.  Cell diagram.  Shows the components of the cell in a symbolic way.  Anode (where oxidation occurs) on the left.  Cathode (where reduction occurs) on the right. ◦Boundary between phases shown by |. ◦Boundary between half cells (usually a salt bridge) shown by ||.

7 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 7 of 54 Terminology Zn(s) | Zn 2+ (aq) || Cu 2+ (aq) | Cu(s) E cell = V

8 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 8 of 54 Terminology  Galvanic cells.  Produce electricity as a result of spontaneous reactions.  Electrolytic cells.  Non-spontaneous chemical change driven by electricity.  Couple, M|M n+  A pair of species related by a change in number of e -.

9 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 9 of Standard Electrode Potentials  Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.  The potential of an individual electrode is difficult to establish.  Arbitrary zero is chosen. The Standard Hydrogen Electrode (SHE)

10 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 10 of 54 Standard Hydrogen Electrode 2 H + (a = 1) + 2 e - H 2 (g, 1 bar) E° = 0 V Pt|H 2 (g, 1 bar)|H + (a = 1)

11 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 11 of 54 Standard Electrode Potential, E°  E° defined by international agreement.  The tendency for a reduction process to occur at an electrode.  All ionic species present at a=1 (approximately 1 M).  All gases are at 1 bar (approximately 1 atm).  Where no metallic substance is indicated, the potential is established on an inert metallic electrode (ex. Pt).

12 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 12 of 54 Reduction Couples Cu 2+ (1M) + 2 e - → Cu(s)E° Cu 2+ /Cu = ? Pt|H 2 (g, 1 bar)|H + (a = 1) || Cu 2+ (1 M)|Cu(s) E° cell = V Standard cell potential: the potential difference of a cell formed from two standard electrodes. E° cell = E° cathode - E° anode cathodeanode

13 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 13 of 54 Standard Cell Potential Pt|H 2 (g, 1 bar)|H + (a = 1) || Cu 2+ (1 M)|Cu(s) E° cell = V E° cell = E° cathode - E° anode E° cell = E° Cu 2+ /Cu - E° H + /H V = E° Cu 2+ /Cu - 0 V E° Cu 2+ /Cu = V H 2 (g, 1 atm) + Cu 2+ (1 M) → H + (1 M) + Cu(s) E° cell = V

14 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 14 of 54 Measuring Standard Reduction Potential anode cathode

15 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 15 of 54 Standard Reduction Potentials

16 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 16 of 54

17 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 17 of 54

18 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 18 of E cell, ΔG, and K eq  Cells do electrical work.  Moving electric charge.  Faraday constant, F = 96,485 C mol -1  elec = -nFE ΔG = -nFE ΔG° = -nFE° Michael Faraday

19 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 19 of 54 Combining Half Reactions Fe 3+ (aq) + 3e - → Fe(s) E° Fe 3+ /Fe = ? Fe 2+ (aq) + 2e - → Fe(s) E° Fe 2+ /Fe = V Fe 3+ (aq) + 3e - → Fe 2+ (aq) E° Fe 3+ /Fe 2+ = V Fe 3+ (aq) + 3e - → Fe(s) ΔG° = J ΔG° = J ΔG° = VE° Fe 3+ /Fe = V ΔG° = V = -nFE° E° Fe 3+ /Fe = V /(-3F) = V

20 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 20 of 54 Spontaneous Change  ΔG < 0 for spontaneous change.  Therefore E° cell > 0 because ΔG cell = -nFE° cell  E° cell > 0  Reaction proceeds spontaneously as written.  E° cell = 0  Reaction is at equilibrium.  E° cell < 0  Reaction proceeds in the reverse direction spontaneously.

21 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 21 of 54 The Behavior or Metals Toward Acids M(s) → M 2+ (aq) + 2 e - E° = -E° M 2+ /M 2 H + (aq) + 2 e - → H 2 (g) E° H + /H 2 = 0 V 2 H + (aq) + M(s) → H 2 (g) + M 2+ (aq) E° cell = E° H + /H 2 - E° M 2+ /M = -E° M 2+ /M When E° M 2+ /M 0. Therefore ΔG° < 0. Metals with negative reduction potentials react with acids.

22 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 22 of 54 Relationship Between E° cell and K eq ΔG° = -RT ln K eq = -nFE° cell E° cell = nF RT ln K eq

23 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 23 of 54 Summary of Thermodynamic, Equilibrium and Electrochemical Relationships.

24 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 24 of E cell as a Function of Concentration ΔG = ΔG° -RT ln Q -nFE cell = -nFE cell ° -RT ln Q E cell = E cell ° - ln Q nF RT Convert to log 10 and calculate constants. E cell = E cell ° - log Q n V The Nernst Equation:

25 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 25 of 54 Pt|Fe 2+ (0.10 M),Fe 3+ (0.20 M)||Ag + (1.0 M)|Ag(s) Applying the Nernst Equation for Determining E cell. What is the value of E cell for the voltaic cell pictured below and diagrammed as follows? EXAMPLE 20-8

26 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 26 of 54 E cell = E cell ° - log Q n V Pt|Fe 2+ (0.10 M),Fe 3+ (0.20 M)||Ag + (1.0 M)|Ag(s) E cell = E cell ° - log n V [Fe 3+ ] [Fe 2+ ] [Ag + ] Fe 2+ (aq) + Ag + (aq) → Fe 3+ (aq) + Ag (s) E cell = V – V = V EXAMPLE 20-8

27 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 27 of 54 Concentration Cells Two half cells with identical electrodes but different ion concentrations. 2 H + (1 M) → 2 H + (x M) Pt|H 2 (1 atm)|H + (x M)||H + (1.0 M)|H 2 (1 atm)|Pt(s) 2 H + (1 M) + 2 e - → H 2 (g, 1 atm) H 2 (g, 1 atm) → 2 H + (x M) + 2 e -

28 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 28 of 54 Concentration Cells E cell = E cell ° - log n V x2x E cell = 0 - log V x2x2 1 E cell = V log x E cell = ( V) pH 2 H + (1 M) → 2 H + (x M) E cell = E cell ° - log Q n V

29 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 29 of 54 Measurement of K sp Ag + (0.100 M) → Ag + (sat’d M) Ag|Ag + (sat’d AgI)||Ag + (0.10 M)|Ag(s) Ag + (0.100 M) + e - → Ag(s) Ag(s) → Ag+(sat’d) + e -

30 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 30 of 54 Using a Voltaic Cell to Determine K sp of a Slightly Soluble Solute. With the date given for the reaction on the previous slide, calculate K sp for AgI. AgI(s) → Ag + (aq) + I - (aq) Let [Ag+] in a saturated Ag + solution be x: Ag + (0.100 M) → Ag + (sat’d M) E cell = E cell ° - log Q = n V E cell ° - log n V [Ag + ] 0.10 M soln [Ag + ] sat’d AgI EXAMPLE 20-10

31 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 31 of 54 E cell = E cell ° - log n V [Ag + ] 0.10 M soln [Ag + ] sat’d AgI E cell = E cell ° - log n V x = 0 - (log x – log 0.100) V log log x == -1 – 7.04 = x = = 9.1  K sp = x 2 = 8.3  EXAMPLE 20-10

32 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 32 of Batteries: Producing Electricity Through Chemical Reactions  Primary Cells (or batteries).  Cell reaction is not reversible.  Secondary Cells.  Cell reaction can be reversed by passing electricity through the cell (charging).  Flow Batteries and Fuel Cells.  Materials pass through the battery which converts chemical energy to electric energy.

33 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 33 of 54 The Leclanché (Dry) Cell

34 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 34 of 54 Dry Cell Zn(s) → Zn 2+ (aq) + 2 e - Oxidation: 2 MnO 2 (s) + H 2 O(l) + 2 e - → Mn 2 O 3 (s) + 2 OH - Reduction: NH OH - → NH 3 (g) + H 2 O(l)Acid-base reaction: NH 3 + Zn 2+ (aq) + Cl - → [Zn(NH 3 ) 2 ]Cl 2 (s)Precipitation reaction:

35 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 35 of 54 Alkaline Dry Cell Zn 2+ (aq) + 2 OH - → Zn (OH) 2 (s) Zn(s) → Zn 2+ (aq) + 2 e - Oxidation reaction can be thought of in two steps: 2 MnO 2 (s) + H 2 O(l) + 2 e - → Mn 2 O 3 (s) + 2 OH - Reduction: Zn (s) + 2 OH - → Zn (OH) 2 (s) + 2 e -

36 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 36 of 54 Lead-Acid (Storage) Battery  The most common secondary battery.

37 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 37 of 54 Lead-Acid Battery PbO 2 (s) + 3 H + (aq) + HSO 4 - (aq) + 2 e - → PbSO 4 (s) + 2 H 2 O(l) Oxidation: Reduction: Pb (s) + HSO 4 - (aq) → PbSO 4 (s) + H + (aq) + 2 e - PbO 2 (s) + Pb(s) + 2 H + (aq) + HSO 4 - (aq) → 2 PbSO 4 (s) + 2 H 2 O(l) E° cell = E° PbO 2 /PbSO 4 - E° PbSO 4 /Pb = 1.74 V – (-0.28 V) = 2.02 V

38 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 38 of 54 The Silver-Zinc Cell: A Button Battery Zn(s),ZnO(s)|KOH(sat’d)|Ag 2 O(s),Ag(s) Zn(s) + Ag 2 O(s) → ZnO(s) + 2 Ag(s) E cell = 1.8 V

39 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 39 of 54 The Nickel-Cadmium Cell Cd(s) + 2 NiO(OH)(s) + 2 H 2 O(L) → 2 Ni(OH) 2 (s) + Cd(OH) 2 (s)

40 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 40 of 54 Fuel Cells O 2 (g) + 2 H 2 O(l) + 4 e - → 4 OH - (aq) 2{H 2 (g) + 2 OH - (aq) → 2 H 2 O(l) + 2 e - } 2H 2 (g) + O 2 (g) → 2 H 2 O(l) E° cell = E° O 2 /OH - - E° H 2 O/H 2 = V – ( V) = V  = ΔG°/ ΔH° = 0.83

41 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 41 of 54 Air Batteries 4 Al(s) + 3 O 2 (g) + 6 H 2 O(l) + 4 OH - → 4 [Al(OH) 4 ](aq)

42 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 42 of Corrosion: Unwanted Voltaic Cells O 2 (g) + 2 H 2 O(l) + 4 e - → 4 OH - (aq) 2 Fe(s) → 2 Fe 2+ (aq) + 4 e - 2 Fe(s) + O 2 (g) + 2 H 2 O(l) → 2 Fe 2+ (aq) + 4 OH - (aq) E cell = V E O 2 /OH - = V E Fe/Fe 2+ = V In neutral solution: In acidic solution: O 2 (g) + 4 H + (aq) + 4 e - → 4 H 2 O (aq) E O 2 /OH - = V

43 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 43 of 54 Corrosion

44 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 44 of 54 Corrosion Protection

45 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 45 of 54 Corrosion Protection

46 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 46 of Electrolysis: Causing Non-spontaneous Reactions to Occur Galvanic Cell: Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s) E O 2 /OH - = V Electolytic Cell: Zn 2+ (aq) + Cu(s) → Zn(s) + Cu 2+ (aq) E O 2 /OH - = V

47 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 47 of 54 Predicting Electrolysis Reaction  An Electrolytic Cell  e - is the reverse of the voltaic cell.  Battery must have a voltage in excess of V in order to force the non-spontaneous reaction.

48 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 48 of 54 Complications in Electrolytic Cells  Overpotential.  Competing reactions.  Non-standard states.  Nature of electrodes.

49 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 49 of 54 Quantitative Aspects of Electrolysis 1 mol e - = C Charge (C) = current (C/s)  time (s) n e - = I  t F

50 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 50 of Industrial Electrolysis Processes

51 Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 51 of 54 Chlor-Alkali Process


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