Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 1 of 54 Juana Mendenhall, Ph.D. Assistant Professor Lecture 4 March 22 Chapter 20: Electrochemistry

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 2 of 54 Objectives 1.Compare and contrast the difference between a dry cell battery, lead storage battery, silver-zinc battery, and fuel cell. 2.Define corrosion of metals, define how corrosion of metals occurs in electrochemical cells, and define how methods of corrosion protection

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 3 of 54 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.

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 4 of 54 The Dry Cell Zn(s) → Zn 2+ (aq) + 2 e - Anode/Oxidation: 2 MnO 2 (s) + 2NH 4 (aq) + 2 e - → Mn 2 O 3 (s) + 2NH 3 (aq) 2 H 2 O(l) Cathode/Reductn: Used in flashlights & transistor radios Cell not completely dry; contains a moist electrolyte paste Voltage 1.5 V

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 5 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:

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 6 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 -

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 7 of 54 Lead-Acid (Storage) Battery  The most common secondary battery.

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 8 of 54 Lead-Acid Battery Each cell produces 2 V, a total of 12 V from six cells

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 9 of 54 Lead-Acid Battery PbO 2 (s) + 3 H + (aq) + HSO 4 - (aq) + 2 e - → PbSO 4 (s) + 2 H 2 O(l) Anode/Oxidation: Cathode/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

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 10 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

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 11 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 Hydrogen Fuel Cell Fuel + oxygen → oxidation products

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 12 of 54 Corrosion in electrochemical systems Fe (s) → Fe 2+ (aq) + 2e - Corrosion: the deterioration of metals by an electrochemical process. A region of the metal’s surface serves as the anode: The electrons given up by iron reduce the atmospheric oxygen to water at the cathode, which is the other region of the same metal’s surface: O 2 (g) + 4H + (aq) + 4e - → 2H 2 O (l) The overall reaction is: Fe (s) + O 2 (g) + 4H + (aq) → Fe 2+ (aq) + 2H2O

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 13 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

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 14 of 54 Corrosion (A) Aluminum foil was placed in copper chloride (CuCl M ) and (B) Aluminum foil was placed in olive oil first then placed in CuCl M (A) (B) What do you think will happen?

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 15 of 54 Corrosion Protection

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 16 of 54 End of Chapter Summary  All electrochemical reactions involve the transfer of electrons and are therefore redox reactions.  In a galvanic cell, electricity is produced by a spontaneous chemical reaction. The oxidation at the anode and the reduction at the cathode take place seprately, while the electrons flow through an external circuit.  The electromotive force (emf) of a cell is the voltage difference b/w the two electrodes. In the external circuit, electrons flow from the anode to the cathode in a galvanic cell. In solution, the anions move toward the anode and the cations move toward the cathode.

Prentice-Hall © 2007 General Chemistry: Chapter 20 Slide 17 of 54 End of Chapter Summary  The quantity of electricity carried by 1 mole of electrons is called a faraday, which is equal to 95,500 coulombs.  Standard reduction potentials show the relative likelihood of half-cell reduction reactions and can be used to predict the products, direction, and spontaneity of redox reactions between various substances.  The decrease in free energy of the system in a spontaneous redox reaction is equal to the electrical work done by the system on the surrounds, or  G = -nF E