Chapter 20: Electrochemistry Chemistry 140 Fall 2002 General Chemistry Principles and Modern Applications Petrucci • Harwood • Herring 10th Edition Chapter 20: Electrochemistry Mamdouh Abdelsalam King Faisal University Prentice-Hall © 2002 https://sites.google.com/site/engchem142/

General Chemistry: Chapter 21 Chemistry 140 Fall 2002 Contents 20-1 Electrode Potentials and Their Measurement 20-2 Standard Electrode Potentials 20-3 Ecell, ΔG, and Keq 20-4 Ecell as a Function of Concentration 20-5 Batteries: Producing Electricity Through Chemical Reactions. 20-6 Corrosion: Unwanted Voltaic Cells 20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur 20-8 Industrial Electolysis Processes Focus On Membrane Potentials Prentice-Hall © 2002 General Chemistry: Chapter 21

20-1 Electrode Potentials and Their Measurement Cu(s) + 2Ag+(aq) Cu2+(aq) + 2 Ag(s) Cu(s) + Zn2+(aq) No reaction Prentice-Hall © 2002 General Chemistry: Chapter 20

An Electrochemical Half Cell Anode Cathode Prentice-Hall © 2002 General Chemistry: Chapter 20

An Electrochemical Cell Chemistry 140 Fall 2002 An Electrochemical Cell Cu ⇢ Cu2+ + 2e- 2Ag+ + 2e-⇢ 2Ag Anode: oxidation Cathode: Reduction Section 1 Sunday Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 21 Salt bridge Salt bridge: To provide ions to the solution to keep the electro neutrality Prentice-Hall © 2002 General Chemistry: Chapter 21

General Chemistry: Chapter 20 Terminology Electromotive force, Ecell. 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. Oil (oxidation is loss), Rig (Reduction is gaining) Boundary between phases shown by |. Boundary between half cells (usually a salt bridge) shown by ||. Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Terminology Zn ⇢Zn2+ + 2e- Cu2+ + 2e- ⇢Cu Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = 1.103 V Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Terminology Galvanic cells. Convert chemical energy to electricity as a result of spontaneous reactions. Electrolytic cells. Convert electricity into chemical energy Non-spontaneous chemical change driven by electricity. Couple, M|Mn+ A pair of species related by a change in number of e-. Prentice-Hall © 2002 General Chemistry: Chapter 20

20-2 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) Prentice-Hall © 2002 General Chemistry: Chapter 20

Standard Hydrogen Electrode Chemistry 140 Fall 2002 Standard Hydrogen Electrode 2 H+(a = 1) + 2 e- ⇄ H2(g, 1 bar) E° = 0 V Pt|H2(g, 1 bar)|H+(a = 1) The two vertical lines indicate three phases are present. For simplicity we usually assume that a = 1 at [H+] = 1 M and replace 1 bar by 1 atm. Prentice-Hall © 2002 General Chemistry: Chapter 20

Standard Electrode Potential, E° Chemistry 140 Fall 2002 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). Tuesday 12/12/2017 Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Reduction Couples Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ? Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V anode cathode Standard cell potential: the potential difference of a cell formed from two standard electrodes. E°cell = E°cathode - E°anode Prentice-Hall © 2002 General Chemistry: Chapter 20

Standard Cell Potential Chemistry 140 Fall 2002 Standard Cell Potential Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = 0.340 V E°cell = E°cathode - E°anode E°cell = E°Cu2+/Cu - E°H+/H2 0.340 V = E°Cu2+/Cu - 0 V E°Cu2+/Cu = +0.340 V Wednesday (13/12/2017) H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s) E°cell = 0.340 V Prentice-Hall © 2002 General Chemistry: Chapter 20

Measuring Standard Reduction Potential Chemistry 140 Fall 2002 Measuring Standard Reduction Potential anode cathode cathode anode Prentice-Hall © 2002 General Chemistry: Chapter 20

Standard Reduction Potentials Chemistry 140 Fall 2002 Standard Reduction Potentials Section 2 Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Chemistry 140 Fall 2002 20-3 Ecell, ΔG, and Keq Cells do electrical work. Moving electric charge. Faraday constant, F = 96,485 C mol-1 elec = -nFE ΔG = -nFE ΔG° = -nFE° Thursday section 1 (14/12/2017) Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Chemistry 140 Fall 2002 Spontaneous Change ΔG < 0 for spontaneous change. Therefore E°cell > 0 because ΔGcell = -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. Prentice-Hall © 2002 General Chemistry: Chapter 20

The Behavior of Metals Toward Acids M(s) → M2+(aq) + 2 e- E° = -E°M2+/M 2 H+(aq) + 2 e- → H2(g) E°H+/H2 = 0 V 2 H+(aq) + M(s) → H2(g) + M2+(aq) E°cell = E°H+/H2 - E°M2+/M = -E°M2+/M When E°M2+/M < 0, E°cell > 0. Therefore ΔG° < 0. Metals with negative reduction potentials react with acids Prentice-Hall © 2002 General Chemistry: Chapter 20

Relationship Between E°cell and Keq Chemistry 140 Fall 2002 Relationship Between E°cell and Keq ΔG° = -RT ln Keq = -nFE°cell E°cell = nF RT ln Keq Section 1 Tuesday Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Chemistry 140 Fall 2002 Summary of Thermodynamic, Equilibrium and Electrochemical Relationships. Prentice-Hall © 2002 General Chemistry: Chapter 20

20-4 Ecell as a Function of Concentration ΔG = ΔG° -RT ln Q -nFEcell = -nFEcell° -RT ln Q Ecell = Ecell° - ln Q nF RT Convert to log10 and calculate constants Ecell = Ecell° - log Q n 0.0592 V The Nernst Equation: Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Example 20-8 Applying the Nernst Equation for Determining Ecell. What is the value of Ecell for the voltaic cell pictured below and diagrammed as follows? Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s) Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Example 20-8 Ecell = Ecell° - log Q n 0.0592 V Ecell = Ecell° - log n 0.0592 V [Fe3+] [Fe2+] [Ag+] Ecell = 0.029 V – 0.018 V = 0.011 V Pt|Fe2+(0.10 M),Fe3+(0.20 M)||Ag+(1.0 M)|Ag(s) Fe2+(aq) + Ag+(aq) → Fe3+(aq) + Ag (s) Prentice-Hall © 2002 General Chemistry: Chapter 20

20-5 Batteries: Producing Electricity Through Chemical Reactions Primary Cells (or batteries). Cell reaction is not reversible. (not rechargeable) Secondary Cells. Cell reaction can be reversed by passing electricity through the cell (rechargeable). Flow Batteries and Fuel Cells. Materials pass through the battery which converts chemical energy to electric energy. Prentice-Hall © 2002 General Chemistry: Chapter 20

Lead-Acid (Storage) Battery Chemistry 140 Fall 2002 Lead-Acid (Storage) Battery The most common secondary battery Section 1 Thursday Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Lead-Acid Battery Reduction: PbO2(s) + 3 H+(aq) + HSO4-(aq) + 2 e- → PbSO4(s) + 2 H2O(l) Oxidation: Pb (s) + HSO4-(aq) → PbSO4(s) + H+(aq) + 2 e- PbO2(s) + Pb(s) + 2 H+(aq) + 2HSO4-(aq) → 2 PbSO4(s) + 2 H2O(l) E°cell = E°PbO2/PbSO4 - E°PbSO4/Pb = 1.74 V – (-0.28 V) = 2.02 V Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Fuel Cells O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq) 2{H2(g) + 2 OH-(aq) → 2 H2O(l) + 2 e-} 2H2(g) + O2(g) → 2 H2O(l) E°cell = E°O2/OH- - E°H2O/H2 = 0.401 V – (-0.828 V) = 1.229 V  = ΔG°/ ΔH° = 0.83 Prentice-Hall © 2002 General Chemistry: Chapter 20

20-6 Corrosion: Unwanted Voltaic Cells Chemistry 140 Fall 2002 20-6 Corrosion: Unwanted Voltaic Cells In neutral solution: O2(g) + 2 H2O(l) + 4 e- → 4 OH-(aq) EO2/OH- = 0.401 V 2 Fe(s) → 2 Fe2+(aq) + 4 e- EFe/Fe2+ = -0.440 V 2 Fe(s) + O2(g) + 2 H2O(l) → 2 Fe2+(aq) + 4 OH-(aq) Ecell = 0.841 V In acidic solution: Cancelled sem 1 (2017-2018) O2(g) + 4 H+(aq) + 4 e- → 4 H2O (aq) EO2/OH- = 1.229 V Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Corrosion Protection Prentice-Hall © 2002 General Chemistry: Chapter 20

General Chemistry: Chapter 20 Corrosion Protection Prentice-Hall © 2002 General Chemistry: Chapter 20

20-7 Electrolysis: Causing Non-spontaneous Reactions to Occur Galvanic Cell: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) EO2/OH- = 1.103 V Electolytic Cell: Zn2+(aq) + Cu(s) → Zn(s) + Cu2+(aq) EO2/OH- = -1.103 V Prentice-Hall © 2002 General Chemistry: Chapter 20

Faraday’s law The amount of chemical change is proportional to the amount of current passed. m mass of substance M molar mass I current A T time for which the current passe (s) n number of electrons transferred F Faraday constant (96485 C mol-1) M Abdelsalam