1. Prentice Hall ©2004 Chapter 18Slide 1 Redox reaction are those involving the oxidation and reduction of species. OIL – Oxidation Is Loss of electrons. RIG – Reduction Is Gain of electrons. Oxidation and reduction must occur together. They cannot exist alone. Redox Reactions01

2. Prentice Hall ©2004 Chapter 18Slide 2 Oxidation Half-Reaction: Zn(s)  Zn 2+ (aq) + 2 e –. The Zn loses two electrons to form Zn 2+. Redox Reactions02

3. Prentice Hall ©2004 Chapter 18Slide 3 Redox Reactions03 Reduction Half-Reaction: Cu 2+ (aq) + 2 e –  Cu(s) The Cu 2+ gains two electrons to form copper.

4. Prentice Hall ©2004 Chapter 18Slide 4 Overall: Zn(s) + Cu 2+ (aq)  Zn 2+ (aq) + Cu(s) Redox Reactions04

5. Prentice Hall ©2004 Chapter 18Slide 5 Electrochemical Cells01 Electrodes: are usually metal strips/wires connected by an electrically conducting wire. Salt Bridge: is a U-shaped tube that contains a gel permeated with a solution of an inert electrolyte. Anode: is the electrode where oxidation takes place. Cathode: is the electrode where reduction takes place.

6. Prentice Hall ©2004 Chapter 18Slide 6 Electrochemical Cells02 Convention for expressing the cell: Anode Half-Cell || Cathode Half-Cell Electrode | Anode Soln || Cathode Soln | Electrode Zn(s) | Zn 2+ (1 M) || Cu 2+ (1 M) | Cu(s) Electrons flow from anode to cathode. Anode is placed on left by convention.

7. Prentice Hall ©2004 Chapter 18Slide 7 Electrochemical Cells05

8. Prentice Hall ©2004 Chapter 18Slide 8 Electrochemical Cells06

9. Prentice Hall ©2004 Chapter 18Slide 9 Electrochemical Cells07 When selecting two half-cell reactions the more negative value will form the oxidation half-cell. Consider the reaction between zinc and silver: Ag + (aq) + e –  Ag(s)E° = 0.80 V Zn 2+ (aq) + 2 e –  Zn(s)E° = – 0.76 V Therefore, zinc forms the oxidation half-cell: Zn(s)  Zn 2+ (aq) + 2 e – E° = – (–0.76 V)

10. Prentice Hall ©2004 Chapter 18Slide 10 Batteries01 Batteries are the most important practical application of galvanic cells. Single-cell batteries consist of one galvanic cell. Multicell batteries consist of several galvanic cells linked in series to obtain the desired voltage.

11. Prentice Hall ©2004 Chapter 18Slide 11 Batteries02 Lead Storage Battery: A typical 12 volt battery consists of six individual cells connected in series. Anode: Lead grid packed with spongy lead. Pb(s) + HSO 4 – (aq)  PbSO 4 (s) + H + (aq) + 2 e – Cathode: Lead grid packed with lead oxide. PbO 2 (s) + 3 H + (aq) + HSO 4 – (aq) + 2 e –  PbSO 4 (s) + 2 H 2 O(l) Electrolyte: 38% by mass sulfuric acid. Cell Potential: 1.924 V

12. Prentice Hall ©2004 Chapter 18Slide 12 Batteries03 Zinc Dry-Cell: Also called a Leclanché cell, uses a viscous paste rather than a liquid solution. Anode: Zinc metal can on outside of cell. Zn(s)  Zn 2+ (aq) + 2 e – Cathode: MnO 2 and carbon black paste on graphite. 2 MnO 2 (s) + 2 NH 4 + (aq) + 2 e –  Mn 2 O 3 (s) + 2 NH 3 (aq) + 2 H 2 O(l) Electrolyte: NH 4 Cl and ZnCl 2 paste. Cell Potential: 1.5 V but deteriorates to 0.8 V with use.

13. Prentice Hall ©2004 Chapter 18Slide 13 Batteries04 Alkaline Dry-Cell: Modified Leclanché cell which replaces NH 4 Cl with NaOH or KOH. Anode: Zinc metal can on outside of cell. Zn(s) + 2 OH – (aq)  ZnO(s) + H 2 O(l) + 2 e – Cathode: MnO 2 and carbon black paste on graphite. 2 MnO 2 (s) + H 2 O(l) + 2 e –  Mn 2 O 3 (s) + 2 OH – (aq) Electrolyte: NaOH or KOH, and Zn(OH) 2 paste. Cell Potential: 1.5 V but longer lasting, higher power, and more stable current and voltage.

14. Prentice Hall ©2004 Chapter 18Slide 14 Batteries05 Mercury Dry-Cell: Modified Leclanché cell which replaces MnO 2 with HgO and uses a steel cathode. Anode: Zinc metal can on outside of cell. Zn(s) + 2 OH – (aq)  ZnO(s) + H 2 O(l) + 2 e – Cathode: HgO in contact with steel. 2 HgO (s) + H 2 O(l) + 2 e –  Hg(l) + 2 OH – (aq) Electrolyte: KOH, and Zn(OH) 2 paste. Cell Potential: 1.3 V with small size, longer lasting, and stable current and voltage.

15. Prentice Hall ©2004 Chapter 18Slide 15 Batteries06 Nickel–Cadmium Battery: Modified Leclanché cell which is rechargeable. Anode: Cadmium metal. Cd(s) + 2 OH – (aq)  Cd(OH) 2 (s) + 2 e – Cathode: Nickel(III) compound on nickel metal. NiO(OH) (s) + H 2 O(l) + e –  Ni(OH) 2 (s) + OH – (aq) Electrolyte: Nickel oxyhydroxide, NiO(OH). Cell Potential: 1.30 V

16. Prentice Hall ©2004 Chapter 18Slide 16 Batteries07 Nickel–Metal–Hydride (NiMH): Replaces toxic Cd anode with a hydrogen atom impregnated ZrNi 2 metal alloy. During oxidation at the anode, hydrogen atoms are released as H 2 O. Recharging reverses this reaction.

17. Prentice Hall ©2004 Chapter 18Slide 17 Batteries08 Lithium Ion (Li–ion): The newest rechargeable battery is based on the migration of Li + ions. Anode: Li metal, or Li atom impregnated graphite. Li(s)  Li + + e – Cathode: Metal oxide or sulfide that can accept Li +. MnO 2 (s) + Li + (aq) + e –  LiMnO 2 (s) Electrolyte: Lithium-containing salt such as LiClO 4, in organic solvent. Solid state polymers can also be used. Cell Potential: 3.0 V

18. Prentice Hall ©2004 Chapter 18Slide 18 Batteries09 Fuel Cell: Uses externally fed CH 4 or H 2, which react to form water. Most common is H 2. Anode: Porous carbon containing metallic catalysts. 2 H 2 (s) + 4 OH – (aq)  4 H 2 O(l) + 4 e – Cathode: Porous carbon containing metallic catalysts. O 2 (s) + 2 H 2 O(l) + 4 e –  4 OH – (aq) Electrolyte: Hot aqueous KOH solution. Cell Potential: 1.23 V, but is only 40% of cell capacity.

19. Prentice Hall ©2004 Chapter 18Slide 19 Batteries10 Fuel cells are not batteries because they are not self- contained. Fuel cells typically have about 40% conversion to electricity; the remainder is lost as heat. Excess heat can be used to drive turbine generators.

20. Prentice Hall ©2004 Chapter 18Slide 20 Corrosion01 Corrosion is the oxidative deterioration of metal. 25% of steel produced in USA goes to replace steel structures and products destroyed by corrosion. Rusting of iron requires the presence of BOTH oxygen and water. Rusting results from tiny galvanic cells formed by water droplets.

21. Prentice Hall ©2004 Chapter 18Slide 21 Corrosion02 Oxidation:Fe(s)  Fe 2+ (aq) + 2 e – Reduction:O 2 (g) + 4 H + (aq) + 4 e –  2 H 2 O(l) Overall: 2 Fe(s) + O 2 (g) + 4 H + (aq)  2 Fe 2+ (aq) + 2 H 2 O(l)

22. Prentice Hall ©2004 Chapter 18Slide 22 Corrosion03 Galvanizing: is the coating of iron with zinc. Zinc is more easily oxidized than iron, which protects and reverses oxidation of the iron. Cathodic Protection: is the protection of a metal from corrosion by connecting it to a metal (a sacrificial anode) that is more easily oxidized. All that is required is an electrical connection to the sacrificial anode (usually magnesium or zinc).

23. Prentice Hall ©2004 Chapter 18Slide 23 Corrosion04 Cathodic Protection with Zinc Layer

24. Prentice Hall ©2004 Chapter 18Slide 24 Corrosion05 Cathodic Protection with Magnesium Anode

25. Prentice Hall ©2004 Chapter 18Slide 25 Electrolysis: is the process in which electrical energy is used to drive a nonspontaneous chemical reaction. An electrolytic cell is an apparatus for carrying out electrolysis. Processes in an electrolytic cell are the reverse of those in a galvanic cell. Electrolysis01

26. Prentice Hall ©2004 Chapter 18Slide 26 Electrolysis05 Electrolysis of Water: Requires an electrolyte species, that is less easily oxidized and reduced than water, to carry the current. Anode: Water is oxidized to oxygen gas. 2 H 2 O(l)  O 2 (g) + 4 H + (aq) + 4 e – Cathode: Water is reduced to hydrogen gas. 4 H 2 O(l) + 4 e –  2 H 2 (g) + 4 OH – (aq)

27. Prentice Hall ©2004 Chapter 18Slide 27 Electrolysis02 Electrolysis of Molten Sodium Chloride:Electrolysis of Aqueous Sodium Chloride: uses different processes to molten sodium chloride. Based on cell potentials, water (–0.83 V) would be preferentially reduced over sodium ions (–2.71 V). Based on cell potentials, water (+1.23 V) should be preferentially oxidized over chloride ions (+1.36 V). The observed product at the anode is Cl 2, not O 2, because of a phenomenon called overvoltage.

28. Prentice Hall ©2004 Chapter 18Slide 28 Electrolysis04 Overvoltage: Additional voltage needed to maintain the rate of electron transfer at the electrode–solution interface. Overvoltage is required when a half-reaction has a significant activation energy, and so a slow rate. Overvoltage for formation of O 2 or H 2 is much greater than for formation of Cl 2.

29. Prentice Hall ©2004 Chapter 18Slide 29 Electrolysis07 Quantitative Electrolysis: The amount of substance produced at an electrode by electrolysis depends on the quantity of charge passed through the cell. Reduction of 1 mol of sodium ions requires 1 mol of electrons to pass through the system. The charge on 1 mol of electrons is 96,500 coulombs.

30. Prentice Hall ©2004 Chapter 18Slide 30 Electrolysis Applications01 Manufacture of Sodium (Downs Cell):

31. Prentice Hall ©2004 Chapter 18Slide 31 Electrolysis Applications02 Manufacture of Cl 2 and NaOH (Chlor–Alkali):

32. Prentice Hall ©2004 Chapter 18Slide 32 Electrolysis Applications03 Manufacture of Aluminum (Hall–Heroult):

33. Prentice Hall ©2004 Chapter 18Slide 33 Electrolysis Applications04 Electrorefining and Electroplating: