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Redox Reactions and Electrochemistry

Redox Reactions and Electrochemistry. Chapter 19. Applications of Oxidation-Reduction Reactions. Zn (s) Zn 2+ ( aq ) + 2e -. +. 2NH 4 ( aq ) + 2MnO 2 ( s ) + 2e - Mn 2 O 3 ( s ) + 2NH 3 ( aq ) + H 2 O ( l ).

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Redox Reactions and Electrochemistry

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  1. Redox Reactions and Electrochemistry Chapter 19

  2. Applications of Oxidation-Reduction Reactions

  3. Zn (s) Zn2+ (aq) + 2e- + 2NH4(aq) + 2MnO2(s) + 2e- Mn2O3(s) + 2NH3(aq) + H2O (l) Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s) Batteries Dry cell Leclanché cell Anode: Cathode: 19.6

  4. Batteries Most common nonrechargable battery; provides far superior performance over older “dry cells” that were also based on Zn and MnO2 as the electrochemically active substances Alkaline battery (1.5 V) Anode: Zn (s) + 2 OH-(aq) Zn(OH)2(s) + 2e- 2 MnO2(s)+2 H2O (l)+ 2e- 2 MnO(OH) (s) + 2OH-(aq) Cathode:

  5. Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e- HgO (s) + H2O (l) + 2e- Hg (l) + 2OH-(aq) Zn(Hg) + HgO (s) ZnO (s) + Hg (l) Batteries Used in pacepakers and hearing aids Mercury Battery Anode: Cathode: 19.6

  6. Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4 PbO2(s) + 4H+(aq) + SO2-(aq) + 2e- PbSO4(s) + 2H2O (l) 4 Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) 2PbSO4 (s) + 2H2O (l) 4 Batteries Used in cars and trucks Lead storage battery Anode: Cathode: 19.6

  7. Batteries Used in laptops and cell phones Solid State Lithium Battery 19.6

  8. 2H2 (g) + 4OH- (aq) 4H2O (l) + 4e- O2(g) + 2H2O (l) + 4e- 4OH-(aq) 2H2 (g) + O2 (g) 2H2O (l) Batteries Used in space vehicles: liquid H2 and O2 are stored as fuel, and the product of the reaction is consumed by the spacecraft crew A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning (not batteries because they are not self-contained systems) Anode: Cathode: 19.6

  9. Hydrogen Fuel Cells

  10. Corrosion • Deterioration of metals by an electrochemical process • Rust on iron • Tarnish on silver • Green patina formed on copper and brass

  11. Corrosion • Oxygen gas and water must be present for iron to rust • Reactions are quite complex and not completely undersood, but the main steps are outlined here • Note that reaction occurs in acidic medium 19.7

  12. In cold climates, salts spread on roadways to melt ice and snow are a major cause of rust formation on automobiles Electrical circuit is completed by the migration of electrons and ions Therefore, rusting occurs rapidly in salt water Corrosion

  13. Corrosion Prevention Employing a sacrificial anode to prevent Fe oxidation

  14. Cathodic Protection of an Iron Storage Tank Connecting to a metal that oxidizes more-readily 19.7

  15. Corrosion prevention Coating the surface with a metal that oxidizes less-readily • A “tin” can is made by applying a thin layer of tin over iron • Rust is prevented as long as the tin layer remains intact • However, once the surface has been scratched, rusting occurs rapidly

  16. Electrolysis Electricity can be used to decompose molten NaCl into its component elements • Voltaic cells are based on spontaneous oxidation-reduction reactions • Conversely, it is possible to use electrical energy to cause nonspontaneous redox reactions to occur • Such processes, which are driven by an outside source of electrical energy, are called electrolysis reactions and take place in an electrolytic cell

  17. Electrolysis Electricity can be used to decompose molten NaCl into its component elements • This is the reason manufacturers of automotive batteries caution against immersing the battery in salt water the standard 12-V car battery has more than enough electromotive force to produce harmful products, such as poisonous Cl2 gas!

  18. A battery (or some other source of direct electrical current) acts as an electron pump, pushing electrons into one electrode and pulling them from the other. Downs cell Simplified schematic 19.8

  19. The electrodes are inert; they do not undergo a reaction but merely serve as the surface where oxidation and reduction occur. Downs cell Simplified schematic 19.8

  20. Electrolysis Electricity can be used to decompose molten NaCl into its component elements • Note that the cathode of the voltage source is connected to the anode of the electrolytic cell • And that the anode of the voltage source is connected to the cathode of the electrolytic cell • Thus the circuit is complete

  21. Electrolysis Electricity can be used to decompose molten NaCl into its component elements • Na is not found free in nature due to its great reactivity • It is obtained commercially by the electrolysis of dry molten sodium chloride • Sodium is a soft, silvery-white metal which is generally stored in paraffin, as it oxidises rapidly when cut.

  22. Electrolysis Electricity can be used to decompose molten NaCl into its component elements. Why MOLTEN NaCl?

  23. Water undergoes electrolysis 19.8

  24. Water undergoes electrolysis 19.8

  25. Quantitative aspects of electrolysis Electroplating uses electrolysis to deposit a thin layer of one metal onto another metal in order to improve beauty or resistance to corrosion. 19.8

  26. Nickel dissolves from the anode to form Ni2+(aq) At the cathode, Ni2+(aq) is reduced and forms a nickel “plate” on the cathode Quantitative aspects of electrolysis An example of electroplating would be depositing a thin layer of nickel onto steel. 19.8

  27. These Ernie Ball strings are made from nickel-plated steel wire wrapped around tin plated hex shaped steel core wire. Their nickel-wound sets are by far the most popular, producing a well balanced and all around good sound. Quantitative aspects of electrolysis An example of electroplating would be depositing a thin layer of nickel onto steel. 19.8

  28. Quantity of charge passing through an electrical circuit, such as that in an electrolytic cell, is generally measured in coulombs The charge on 1 mole of electrons is 96,485 C (1 faraday) A coulomb is the quantity of charge passing a point in a circuit in 1 s when the current is 1 ampere (A) Therefore, number of coulombs passing through a cell can be obtained by multiplying the amperage and the elapsed time in seconds. Quantitative aspects of electrolysis For any half-reaction, the amount of a substance that is reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons passed into the cell. 19.8

  29. Electrolysis and Mass Changes charge (C) = current (A) x time (s) 1 mole e- = 96,500 C 19.8

  30. How much Ca will be produced in an electrolytic cell of molten CaCl2 if a current of 0.452 A is passed through the cell for 1.5 hours? 2 mole e- = 1 mole Ca mol Ca = 0.452 x 1.5 hr x 3600 C s 2Cl- (l) Cl2 (g) + 2e- hr s Ca2+(l) + 2e- Ca (s) 1 mol Ca 1 mol e- x x 96,500 C 2 mol e- Ca2+ (l) + 2Cl- (l) Ca (s) + Cl2 (g) Anode: Cathode: = 0.0126 mol Ca = 0.50 g Ca 19.8

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