Chapter 17: Electrochemistry 12/26/2017 Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Electrochemistry: The area of chemistry concerned with the interconversion of chemical and electrical energy. Galvanic (Voltaic) Cell: A spontaneous chemical reaction which generates an electric current. Electrolytic Cell: An electric current which drives a nonspontaneous reaction. Galvanic and electrolytic cells are the reverse of each other. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) Oxidation half-reaction: Zn2+(aq) + 2e- Zn(s) Reduction half-reaction: Cu(s) Cu2+(aq) + 2e- Balancing redox reactions was covered in Ch 4 (Reactions in Aqueous Solution). It’s common for people to postpone redox balancing until now. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) Be careful about assumptions. Don’t always place the same electrode on the same side. For instance, in the past, textbooks would commonly placed the anode on the left side which led some students to believe that the anode always belongs on the left. Students often are confused as the purpose and need of a salt bridge. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Anode: The electrode where oxidation occurs. The electrode where electrons are produced. Is what anions migrate toward. Has a negative sign. The wire convention is used for the sign at the anode and cathode. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Anode: The electrode where oxidation occurs. The electrode where electrons are produced. Is what anions migrate toward. Has a negative sign. Cathode: The electrode where reduction occurs. The electrode where electrons are consumed. Is what cations migrate toward. Has a positive sign. “Anode” and “oxidation” both begin with vowels. “Cathode” and “reduction” both begin with consonants. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Galvanic Cells Anode half-reaction: Zn2+(aq) + 2e- Zn(s) Cathode half-reaction: Cu(s) Cu2+(aq) + 2e- Overall cell reaction: Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Shorthand Notation for Galvanic Cells Chapter 17: Electrochemistry 12/26/2017 Shorthand Notation for Galvanic Cells Anode half-reaction: Zn2+(aq) + 2e- Zn(s) Cathode half-reaction: Cu(s) Cu2+(aq) + 2e- Overall cell reaction: Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) Salt bridge Anode half-cell Cathode half-cell Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Electron flow Phase boundary Phase boundary Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Cell Potentials and Free-Energy Changes for Cell Reactions Chapter 17: Electrochemistry 12/26/2017 Cell Potentials and Free-Energy Changes for Cell Reactions Electromotive Force (emf): The force or electrical potential that pushes the negatively charged electrons away from the anode (- electrode) and pulls them toward the cathode (+ electrode). It is also called the cell potential (E) or the cell voltage. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Cell Potentials and Free-Energy Changes for Cell Reactions Chapter 17: Electrochemistry 12/26/2017 Cell Potentials and Free-Energy Changes for Cell Reactions 1 J = 1 C x 1 V joule SI unit of energy volt SI unit of electric potential coulomb Electric charge 1 coulomb is the amount of charge transferred when a current of 1 ampere flows for 1 second. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Cell Potentials and Free-Energy Changes for Cell Reactions Chapter 17: Electrochemistry 12/26/2017 Cell Potentials and Free-Energy Changes for Cell Reactions faraday or Faraday constant the electric charge on 1 mol of electrons 96,5000 C/mol e- DG = -nFE or DG° = -nFE° free-energy change The “°” refers to standard free-energy change and standard cell potential. cell potential number of moles of electrons transferred in the reaction Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Cell Potentials and Free-Energy Changes for Cell Reactions Chapter 17: Electrochemistry 12/26/2017 Cell Potentials and Free-Energy Changes for Cell Reactions The standard cell potential at 25 °C is 0.10 V for the reaction: Zn2+(aq) + Cu(s) Zn(s) + Cu2+(aq) Calculate the standard free-energy change for this reaction at 25 °C. DG° = -nFE° The sign for a spontaneous reaction for free-energy: - The sign for a spontaneous reaction for cell potential: + mol e- 96,500 C 1 C V 1 J 1000 J 1 kJ = -(2 mol e-) (1.10 V) DG° = -212 kJ Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials 2H1+(aq) + Cu(s) H2(g) + Cu2+(aq) Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials Anode half-reaction: 2H1+(aq) + 2e- H2(g) Cathode half-reaction: Cu(s) Cu2+(aq) + 2e- Overall cell reaction: 2H1+(aq) + Cu(s) H2(g) + Cu2+(aq) The standard potential of a cell is the sum of the standard half-cell potentials for oxidation at the anode and reduction at the cathode: Individual half-cell potentials can’t be measured. They must be measured in pairs. E°cell = E°ox + E°red The measured potential for this cell: E°cell = 0.34 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials The standard hydrogen electrode (S.H.E.) has been chosen to be the reference electrode. Potentials can only be measured in pairs which is why we need to establish a reference point for our half-cell potentials. Much like what was done for standard enthalpies and standard free energies. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials The standard hydrogen electrode (S.H.E.) has been chosen to be the reference electrode. H2(g, 1 atm) 2H1+(aq, 1 M) + 2e- E°ox = 0 V E°red = 0 V 2H1+(aq, 1 M) + 2e- H2(g, 1 atm) The half-cell potentials for the S.H.E. have been defined to be 0 V. Much like the atomic mass of carbon-12 has been defined to 12 u. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials Anode half-reaction: 2H1+(aq) + 2e- H2(g) Cathode half-reaction: Cu(s) Cu2+(aq) + 2e- Overall cell reaction: 2H1+(aq) + Cu(s) H2(g) + Cu2+(aq) E°cell = E°ox + E°red 0.34 V = 0 V + E°red Recall that the 0.34 V was measured for the cell with the half-cells as described above. The oxidation half-cell potential is defined to be 0 V since it is the S.H.E. Note also that the “ox” and “red” subscripts for oxidation and reduction have been dropped. While this can be very confusing to students, it is common usage. A standard reduction potential can be defined: Cu(s) Cu2+(aq) + 2e- E° = 0.34 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials H2(g) + Zn2+(aq) 2H1+(aq) + Zn(s) Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials Anode half-reaction: H2(g) 2H1+(aq) + 2e- Cathode half-reaction: Zn2+(aq) + 2e- Zn(s) Overall cell reaction: H2(g) + Zn2+(aq) 2H1+(aq) + Zn(s) E°cell = E°ox + E°red 0.76 V = E°ox + 0 V The half-cell potential sign flips for an oxidation potential (the table has reduction potentials). Zn2+(aq) + 2e- Zn(s) E° = 0.76 V As a standard reduction potential: Zn(s) Zn2+(aq) + 2e- E° = -0.76 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Standard Reduction Potentials Copyright © 2008 Pearson Prentice Hall, Inc.

Using Standard Reduction Potentials Chapter 17: Electrochemistry 12/26/2017 Using Standard Reduction Potentials Zn2+(aq) + 2e- Zn(s) E° = -(-0.76 V) Cu(s) Cu2+(aq) + 2e- E° = 0.34 V E° = 1.10 V Zn2+(g) + Cu(s) Zn(s) + Cu2+(aq) Ag(s)] 2 x [Ag1+(aq) + e- E° = 0.80 V Cu2+(aq) + 2e- Cu(s) E° = -0.34 V E° = 0.46 V 2Ag2+(g) + Cu(s) 2Ag(s) + Cu2+(aq) Half-cell potentials are intensive properties. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 The Nernst Equation DG = DG° + RT ln Q Using: DG = -nFE and DG° = -nFE° ln Q nF RT E = E° - Nernst Equation: or log Q nF 2.303RT E = E° - n is the number of moles of transferred electrons from the balanced reaction (balanced number of electrons from the half-reactions). Also, just the number is used (“2” instead of “2 mol e-”). or log Q n 0.0592 V E = E° - in volts, at 25°C Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 The Nernst Equation Consider a galvanic cell that uses the reaction: Cu2+(aq) + 2Fe2+(aq) Cu(s) + 2Fe3+(aq) What is the potential of a cell at 25 °C that has the following ion concentrations? [Fe3+] = 1.0 x 10-4 M [Cu2+] = 0.25 M [Fe2+] = 0.20 M Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 The Nernst Equation log Q n 0.0592 V E = E° - Calculate E°: Cu2+(aq) + 2e- Cu(s) E° = -0.34 V Fe2+(aq) Fe3+(aq) + e- E° = 0.77 V E°cell = -0.34 V + 0.77 V = 0.43 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 The Nernst Equation log Q n 0.0592 V E = E° - Calculate E: n 0.0592 V [Fe3+]2 [Cu2+][Fe2+]2 E = E° - log 2 0.0592 V (1.0 x 10-4)2 (0.25)(0.20)2 = 0.43 V - log E = 0.25 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Cell Potentials and Equilibrium Constants Chapter 17: Electrochemistry 12/26/2017 Standard Cell Potentials and Equilibrium Constants Using DG° = -nFE° and DG° = -RT ln K -nFE° = -RT ln K nF RT nF 2.303 RT E° = ln K = log K DG° = -RT ln K came from Ch 16 (Thermodynamics: Entropy, Free Energy, and Equilibrium). The mathematical reduction from 2.303RT/nF to 0.0592V/n is the same as what was done for the Nernst equation. n 0.0592 V E° = log K in volts, at 25°C Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Cell Potentials and Equilibrium Constants Chapter 17: Electrochemistry 12/26/2017 Standard Cell Potentials and Equilibrium Constants Copyright © 2008 Pearson Prentice Hall, Inc.

Standard Cell Potentials and Equilibrium Constants Chapter 17: Electrochemistry 12/26/2017 Standard Cell Potentials and Equilibrium Constants Three methods to determine equilibrium constants: K = [A]a[B]b [C]c[D]d K from concentration data: RT -DG° ln K = K from thermochemical data: ln K nF RT E° = K from electrochemical data: or RT nFE° ln K = Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Batteries Lead Storage Battery Anode: PbSO4(s) + H1+(aq) + 2e- Pb(s) + HSO41-(aq) Cathode: PbSO4(s) + 2H2O(l) PbO2(s) + 3H1+(aq) + HSO41-(aq) + 2e- Overall: 2PbSO4(s) + 2H2O(l) Pb(s) + PbO2(s) + 2H1+(aq) + 2HSO41-(aq) Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Batteries Dry-Cell Batteries Leclanché cell Anode: Zn2+(aq) + 2e- Zn(s) Cathode: Mn2O3(s) + 2NH3(aq)+ H2O(l) 2MnO2(s) + 2NH41+(aq) + 2e- Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Batteries Dry-Cell Batteries Alkaline cell Anode: ZnO(s) + H2O(l) + 2e- Zn(s) + 2OH1-(aq) Cathode: Mn2O3(s) + 2OH1-(aq) 2MnO2(s) + H2O(l) + 2e- Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Batteries Nickel-Cadmium (“ni-cad”) Batteries Anode: Cd(OH)2(s) + 2e- Cd(s) + 2OH1-(aq) Cathode: Ni(OH)2(s) + OH1-(aq) NiO(OH)(s) + H2O(l) + e- Nickel-Metal Hydride (“NiMH”) Batteries Anode: M(s) + H2O(l) + e- MHab(s) + OH1-(aq) Cathode: Ni(OH)2(s) + OH1-(aq) NiO(OH)(s) + H2O(l) + e- Overall: M(s) + Ni(OH)2(s) MHab(s) + NiO(OH)(s) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Batteries Lithium and Lithium Ion Batteries Lithium Anode: xLi1+(soln) + xe- xLi(s) Cathode: LixMnO2(s) MnO2(s) + xLi1+(soln) + xe- Lithium Ion Notice where lithium is on the standard reduction table. Anode: xLi1+(soln) + 6C(s) + xe- LixC6(s) Cathode: LiCoO2(s) Li1-xCoO2(s) + xLi1+(soln) + xe- Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Fuel Cells Hydrogen-Oxygen Fuel Cell Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Corrosion Corrosion: The oxidative deterioration of a metal. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Corrosion Prevention of Corrosion For some metals, oxidation protects the metal (aluminum, chromium, magnesium, titanium, zinc, and others). For other metals, there are two main techniques. Iron rust doesn’t protect iron since the rust is too porous. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Corrosion Prevention of Corrosion Galvanization: The coating of iron with zinc. Even if the zinc coating is scratched, the iron is still protected since the zinc will be selectively oxidized. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Corrosion Prevention of Corrosion Galvanization: The coating of iron with zinc. When some of the iron is oxidized (rust), the process is reversed since zinc will reduce Fe2+ to Fe: Fe(s) Fe2+(aq) + 2e- E° = -0.45 V Fe2+ is higher up in the reduction table so it will be selectively reduced when in contact with zinc. Zn(s) Zn2+(aq) + 2e- E° = -0.76 V Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Chapter 17: Electrochemistry 12/26/2017 Corrosion Prevention of Corrosion Cathodic Protection: Instead of coating the entire surface of the first metal with a second metal, the second metal is placed in electrical contact with the first metal: Anode: Mg2+(aq) + 2e- Mg(s) E° = 2.37 V Cathode: 2H2O(l) O2(g) + 4H1+(aq) + 4e- Magnesium is farther down in the reduction table (on the right) so it is more likely to be oxidized than zinc. Zinc is used as a sacrificial anode for iron sea walls for hurricane protection (an iron wall placed into salt water for a long period of time is not a good idea from a corrosion perspective!). Blocks of solid zinc (sacrificial anodes) are attached to the iron and spaced at regular intervals. The sacrificial anodes are monitored over time and replaced when needed. E° = 1.23 V Attaching a magnesium stake to iron will corrode the magnesium instead of the iron. Magnesium acts as a sacrificial anode. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chapter 17: Electrochemistry 12/26/2017 Electrolysis and Electrolytic Cells Electrolysis: The process of using an electric current to bring about chemical change. Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chapter 17: Electrochemistry 12/26/2017 Electrolysis and Electrolytic Cells Electrolysis of Molten Sodium Chloride Anode: Cl2(g) + 2e- 2Cl1-(l) Cathode: 2Na(l) 2Na1+(l) + 2e- Overall: 2Na(l) + Cl2(g) 2Na1+(l) + 2Cl1-(l) Copyright © 2008 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chapter 17: Electrochemistry 12/26/2017 Electrolysis and Electrolytic Cells Electrolysis of Aqueous Sodium Chloride Anode: Cl2(g) + 2e- 2Cl1-(aq) Cathode: H2(g) + 2OH1-(aq) 2H2O(l) + 2e- Overall: Cl2(g) + H2(g) + 2OH1-(aq) 2Cl1-(l) + 2H2O(l) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Electrolysis and Electrolytic Cells Chapter 17: Electrochemistry 12/26/2017 Electrolysis and Electrolytic Cells Electrolysis of Water Anode: O2(g) + 4H1+(aq) + 4e- 2H2O(l) Cathode: 2H2(g) + 4OH1-(aq) 4H2O(l) + 4e- Overall: 2H2(g) + O2(g) + 4H1+ + 4OH1-(aq) 6H2O(l) Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chapter 17: Electrochemistry 12/26/2017 Commercial Applications of Electrolysis Down’s Cell for the Production of Sodium Metal Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chapter 17: Electrochemistry 12/26/2017 Commercial Applications of Electrolysis A Membrane Cell for Electrolytic Production of Cl2 and NaOH Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chapter 17: Electrochemistry 12/26/2017 Commercial Applications of Electrolysis Hall-Heroult Process for the Production of Aluminum Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Commercial Applications of Electrolysis Chapter 17: Electrochemistry 12/26/2017 Commercial Applications of Electrolysis Electrorefining of copper metal Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.

Quantitative Aspects of Electrolysis Chapter 17: Electrochemistry 12/26/2017 Quantitative Aspects of Electrolysis Charge(C) = Current(A) x Time(s) 96,500 C 1 mol e- Moles of e- = Charge(A) x Faraday constant Copyright © 2008 Pearson Prentice Hall, Inc. Copyright © 2008 Pearson Prentice Hall, Inc.