2. Chapter 211 Electrochemistry Chemistry 12

3. Chapter 212 Overview Redox reactions Oxidation Numbers Balancing redox reactions Electrochemical Cells Standard Reduction Potentials E° Spontaneity of redox reactions Applied Electrochemistry – Corrosion of iron Electrolysis Batteries

4. Chapter 213 Electrochemistry defined Branch of chemistry which is concerned with the relationships between electricity and chemical reactions Conversion of electrical into chemical energy and vice versa Reduction-Oxidation (REDOX) reactions are primarily involved

5. Chapter 214 Reduction-Oxidation Defined oxidizing agent = substance that cause oxidation by being reduced reducing agent = substance that cause oxidation by being reduced LEO goes GER Loss of Electrons is Oxidation Gain of Electrons is Reduction

6. Chapter 215 A Redox Reaction Example With time, Ag plates out onto Cu metal strip, and Cu strip “disappears.” Cu is oxidized and is the reducing agent Cu(s)  Cu 2+ (aq) + 2e - Ag + is reduced and is the oxidizing agent Ag + (aq) + 2e -  Ag(s) Cu metal Ag + ions

7. Chapter 216 Redox Reaction Example Overall reaction is Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2Ag(s) Ox. #: 0 +1 +2 0 oxidation reduction Ox. # = Oxidation Number Thus: Oxidation is ALSO gain in oxidation number And Reduction is ALSO loss in oxidation number

8. Chapter 217 Oxidation Numbers The charge that an atom would have if the compound in which it were found were ionic Can be used as an “electron-bookkeeping system” to keep track of electrons transferred in redox reactions Oxidation numbers are written as positive or negative numbers and are determined according to the following rules (which you must memorize – no joke!!) See “Oxidation Number or State” handout

9. Chapter 218 Oxidation-Reduction Reactions Zn added to HCl yields the spontaneous reaction Zn(s) + 2H + (aq)  Zn 2+ (aq) + H 2 (g). The oxidation number of Zn has increased from 0 to 2+. The oxidation number of H has reduced from 1+ to 0. Therefore, Zn is oxidized to Zn 2+ while H + is reduced to H 2. H + causes Zn to be oxidized and is the oxidizing agent. Zn causes H + to be reduced and is the reducing agent. Note that the reducing agent is oxidized and the oxidizing agent is reduced.

10. Chapter 219 Balancing Oxidation-Reduction Equations Law of conservation of mass: the amount of each element present at the beginning of the reaction must be present at the end. Conservation of charge: electrons are not lost in a chemical reaction. In complicated redox reactions, we need to look at the transfer of electrons carefully.Half-Reactions Half-reactions are a convenient way of separating oxidation and reduction reactions.

11. Chapter 2110 Balancing Oxidation-Reduction Equations Half-Reactions The half-reactions for Sn 2+ (aq) + 2Fe 3+ (aq)  Sn 4+ (aq) + 2Fe 3+ (aq) are Sn 2+ (aq)  Sn 4+ (aq) +2e - 2Fe 3+ (aq) + 2e -  2Fe 2+ (aq) Oxidation: electrons are products. Reduction: electrons are reagents.

12. Chapter 2111 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions Consider the titration of an acidic solution of Na 2 C 2 O 4 (sodium oxalate, colorless) with KMnO 4 (deep purple). MnO 4 - is reduced to Mn 2+ (pale pink) while the C 2 O 4 2- is oxidized to CO 2. The equivalence point is given by the presence of a pale pink color. If more KMnO 4 is added, the solution turns purple due to the excess KMnO 4.

13. Chapter 2112 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions

14. Chapter 2113 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions What is the balanced chemical equation? 1. Write down the two half reactions. 2. Balance each half reaction: a. First with elements other than H and O. b. Then balance O by adding water. c. Then balance H by adding H +. d. Finish by balancing charge by adding electrons.

15. Chapter 2114 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions 3. Multiply each half reaction to make the number of electrons equal. 4. Add the reactions and simplify. 5. Check! For KMnO 4 + Na 2 C 2 O 4 :

16. Chapter 2115 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions 1. The two incomplete half reactions are MnO 4 - (aq)  Mn 2+ (aq) C 2 O 4 2- (aq)  2CO 2 (g) 2. Adding water and H + yields 8H + + MnO 4 - (aq)  Mn 2+ (aq) + 4H 2 O There is a charge of 7+ on the left and 2+ on the right. Therefore, 5 electrons need to be added to the left: 5e - + 8H + + MnO 4 - (aq)  Mn 2+ (aq) + 4H 2 O

17. Chapter 2116 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions In the oxalate reaction, there is a 2- charge on the left and a 0 charge on the right, so we need to add two electrons: C 2 O 4 2- (aq)  2CO 2 (g) + 2e - 3. To balance the 5 electrons for permanganate and 2 electrons for oxalate, we need 10 electrons for both. Multiplying gives:

18. Chapter 2117 Balancing Oxidation-Reduction Equations Balancing Equations by the Method of Half- Reactions 10e - + 16H + + 2MnO 4 - (aq)  2Mn 2+ (aq) + 8H 2 O 5C 2 O 4 2- (aq)  10CO 2 (g) + 10e - 4. Adding gives: 16H + (aq) + 2MnO 4 - (aq) + 5C 2 O 4 2- (aq)  2Mn 2+ (aq) + 8H 2 O(l) + 10CO 2 (g) 5. Which is balanced!

19. Chapter 2118 Balancing Oxidation-Reduction Equations Balancing Equations for Reactions Occurring in Basic Solution We use OH - and H 2 O rather than H + and H 2 O. The same method as above is used, but OH - is added to “neutralize” the H + used.

20. Chapter 2119 Voltaic Cells The energy released in a spontaneous redox reaction is used to perform electrical work. Voltaic or galvanic cells are devices in which electron transfer occurs via an external circuit. Voltaic cells are spontaneous. If a strip of Zn is placed in a solution of CuSO 4, Cu is deposited on the Zn and the Zn dissolves by forming Zn 2+. Zn is spontaneously oxidized to Zn 2+ by Cu 2+. The Cu 2+ is spontaneously reduced to Cu 0 by Zn. The entire process is spontaneous.

21. Chapter 2120 Voltaic Cells

22. Chapter 2121 Voltaic Cells Voltaic cells consist of –Anode: Zn(s)  Zn 2+ (aq) + 2e 2 –Cathode: Cu 2+ (aq) + 2e -  Cu(s) –Salt bridge (used to complete the electrical circuit): cations move from anode to cathode, anions move from cathode to anode. The two solid metals are the electrodes (cathode and anode). As oxidation occurs, Zn is converted to Zn 2+ and 2e -. The electrons flow towards the anode where they are used in the reduction reaction.

23. Chapter 2122 Voltaic Cells We expect the Zn electrode to lose mass and the Cu electrode to gain mass. “Rules” of voltaic cells: 1. At the anode electrons are products. (Oxidation) 2. At the cathode electrons are reagents. (Reduction) 3. Electrons cannot swim. Electrons flow from the anode to the cathode. Therefore, the anode is negative and the cathode is positive. Electrons cannot flow through the solution, they have to be transported through an external wire. (Rule 3.)

24. Chapter 2123 Voltaic Cells

25. Chapter 2124 Voltaic Cells Anions and cations move through a porous barrier or salt bridge. Cations move into the cathodic compartment to neutralize the excess negatively charged ions (Cathode: Cu 2+ + 2e -  Cu, so the counterion of Cu is in excess). Anions move into the anodic compartment to neutralize the excess Zn 2+ ions formed by oxidation.

26. Chapter 2125 Voltaic Cells A Molecular View of Electrode Processes Consider the spontaneous redox reaction between Zn(s) and Cu 2+ (aq). During the reaction, Zn(s) is oxidized to Zn 2+ (aq) and Cu 2+ (aq) is reduced to Cu(s). On the atomic level, a Cu 2+ (aq) ion comes into contact with a Zn(s) atom on the surface of the electrode. Two electrons are directly transferred from the Zn(s) (forming Zn 2+ (aq)) to the Cu 2+ (aq) (forming Cu(s)).

27. Chapter 2126 Cell EMF The flow of electrons from anode to cathode is spontaneous. Electrons flow from anode to cathode because the cathode has a lower electrical potential energy than the anode. Potential difference: difference in electrical potential. Measured in volts. One volt is the potential difference required to impart one joule of energy to a charge of one coulomb:

28. Chapter 2127 Cell EMF Electromotive force (emf) is the force required to push electrons through the external circuit. Cell potential: E cell is the emf of a cell. For 1M solutions at 25  C (standard conditions), the standard emf (standard cell potential) is called E  cell. Standard Reduction Potentials Convenient tabulation of electrochemical data. Standard reduction potentials, E  red are measured relative to the standard hydrogen electrode (SHE).

29. Chapter 2128 Cell EMF Standard Reduction Potentials

30. Chapter 2129 Cell EMF Standard Reduction Potentials The SHE is the cathode. It consists of a Pt electrode in a tube placed in 1 M H + solution. H 2 is bubbled through the tube. For the SHE, we assign 2H + (aq, 1M) + 2e -  H 2 (g, 1 atm) E  red of zero. The emf of a cell can be calculated from standard reduction potentials: E  cell = E  red (cathode) - E  red (anode)

31. Chapter 2130 Cell EMF Standard Reduction Potentials

32. Chapter 2131 Cell EMF Standard Reduction Potentials Consider Zn(s)  Zn 2+ (aq) + 2e -. We measure E cell relative to the SHE (cathode): E  cell = E  red (cathode) - E  red (anode) 0.76 V = 0 V - E  red (anode). Therefore, E  red (anode) = -0.76 V. Standard reduction potentials must be written as reduction reactions: Zn 2+ (aq) + 2e -  Zn(s), E  red = -0.76 V. Since E  red = -0.76 V we conclude that the reduction of Zn 2+ in the presence of the SHE is not spontaneous.

33. Chapter 2132 Cell EMF Standard Reduction Potentials The oxidation of Zn with the SHE is spontaneous. Changing the stoichiometric coefficient does not affect E  red. Therefore, 2Zn 2+ (aq) + 4e -  2Zn(s), E  red = -0.76 V. Reactions with E  red > 0 are spontaneous reductions relative to the SHE. Reactions with E  red < 0 are spontaneous oxidations relative to the SHE.

34. Chapter 2133 Cell EMF Standard Reduction Potentials The larger the difference between E  red values, the larger E  cell. In a voltaic (galvanic) cell (spontaneous) E  red (cathode) is more positive than E  red (anode). Recall E  cell = E  red (cathode) - E  red (anode)

35. Chapter 2134 Cell EMF Oxidizing and Reducing Agents The more positive E  red the stronger the oxidizing agent on the left. The more negative E  red the stronger the reducing agent on the right. A species on the higher to the left of the table of standard reduction potentials will spontaneously oxidize a species that is lower to the right in the table. That is, F 2 will oxidize H 2 or Li; Ni 2+ will oxidize Al(s). Any species on the right will spontaneously reduce anything that is higher to the left in the series.

36. Chapter 2135 Cell EMF Oxidizing and Reducing Agents

37. Chapter 2136 Spontaneity of Redox Reactions In a voltaic (galvanic) cell (spontaneous) E  red (cathode) is more positive than E  red (anode) since E  cell = E  red (cathode) - E  red (anode) More generally, for any electrochemical process E  = E  red (reduction process) - E  red (oxidation process). A positive E  indicates a spontaneous process (galvanic cell). A negative E  indicates a nonspontaneous process. The above equation is used to understand the activity series.

38. Chapter 2137 Spontaneity of Redox Reactions Consider the displacement of silver by nickel: Ni(s) + 2Ag + (aq)  Ni 2+ (aq) + 2Ag(s) has E  = E  red (Ag + /Ag) - E  red (Ni 2+ /Ni) = (0.80 V) - (-0.28 V) = 1.08 V, which indicates a spontaneous process. EMF and Free-Energy Change We can show that  G = -nFE

39. Chapter 2138 Spontaneity of Redox Reactions EMF and Free-Energy Change  G is the change in free-energy, n is the number of moles of electrons transferred, F is Faraday’s constant, and E is the emf of the cell. We define Since n and F are positive, if  G > 0 then E < 0.

40. Chapter 2139 Effect of Concentration on Cell EMF A voltaic cell is functional until E = 0 at which point equilibrium has been reached. The point at which E = 0 is determined by the concentrations of the species involved in the redox reaction. The Nernst Equation The Nernst equation relates emf to concentration using and noting that

41. Chapter 2140 Effect of Concentration on Cell EMF The Nernst Equation This rearranges to give the Nernst equation: The Nernst equation can be simplified by collecting all the constants together using a temperature of 298 K: (Note that change from natural logarithm to base-10 log.) Remember that n is number of moles of electrons.

42. Chapter 2141 Effect of Concentration on Cell EMF Concentration Cells We can use the Nernst equation to generate a cell that has an emf based solely on difference in concentration. One compartment will consist of a concentrated solution, while the other has a dilute solution. Example: 1.00 M Ni 2+ (aq) and 1.00  10 -3 M Ni 2+ (aq). The cell tends to equalize the concentrations of Ni 2+ (aq) in each compartment. The concentrated solution has to reduce the amount of Ni 2+ (aq) (to Ni(s)), so must be the cathode.

43. Chapter 2142 Effect of Concentration on Cell EMF Concentration Cells Since the two half-reactions are the same, Eº will be zero.

44. Chapter 2143 Effect of Concentration on Cell EMF Cell EMF and Chemical Equilibrium A system is at equilibrium when  G = 0. From the Nernst equation, at equilibrium:

45. Chapter 2144 Batteries Lead-Acid Battery A 12 V car battery consists of 6 cathode/anode pairs each producing 2 V. Cathode: PbO 2 on a metal grid in sulfuric acid: PbO 2 (s) + SO 4 2- (aq) + 4H + (aq) + 2e -  PbSO 4 (s) + 2H 2 O(l) Anode: Pb: Pb(s) + SO 4 2- (aq)  PbSO 4 (s) + 2e -

46. Chapter 2145 Batteries Lead-Acid Battery The overall electrochemical reaction is PbO 2 (s) + Pb(s) + 2SO 4 2- (aq) + 4H + (aq)  2PbSO 4 (s) + 2H 2 O(l) for which E  cell = E  red (cathode) - E  red (anode) = (+1.685 V) - (-0.356 V) = +2.041 V. Wood or glass-fiber spacers are used to prevent the electrodes form touching.

47. Chapter 2146 Batteries Lead-Acid Battery

48. Chapter 2147 Batteries Alkaline Battery Anode: Zn cap: Zn(s)  Zn 2+ (aq) + 2e - Cathode: MnO 2, NH 4 Cl and C paste: 2NH 4 + (aq) + 2MnO 2 (s) + 2e -  Mn 2 O 3 (s) + 2NH 3 (aq) + 2H 2 O(l) The graphite rod in the center is an inert cathode. For an alkaline battery, NH 4 Cl is replaced with KOH. Anode: Zn powder mixed in a gel: Zn(s)  Zn 2+ (aq) + 2e - Cathode: reduction of MnO 2.

49. Chapter 2148 Batteries Alkaline Battery

50. Chapter 2149 Batteries Fuel Cells Direct production of electricity from fuels occurs in a fuel cell. On Apollo moon flights, the H 2 -O 2 fuel cell was the primary source of electricity. Cathode: reduction of oxygen: 2H 2 O(l) + O 2 (g) + 4e -  4OH - (aq) Anode: 2H 2 (g) + 4OH - (aq)  4H 2 O(l) + 4e -

51. Chapter 2150 Batteries Fuel Cells

52. Chapter 2151 Corrosion Corrosion of Iron Since E  red (Fe 2+ ) < E  red (O 2 ) iron can be oxidized by oxygen. Cathode: O 2 (g) + 4H + (aq) + 4e -  2H 2 O(l). Anode: Fe(s)  Fe 2+ (aq) + 2e -. Dissolved oxygen in water usually causes the oxidation of iron. Fe 2+ initially formed can be further oxidized to Fe 3+ which forms rust, Fe 2 O 3.xH 2 O(s). Oxidation occurs at the site with the greatest concentration of O 2.

53. Chapter 2152 Corrosion Corrosion of Iron

54. Chapter 2153 Corrosion Preventing the Corrosion of Iron Corrosion can be prevented by coating the iron with paint or another metal. Galvanized iron is coated with a thin layer of zinc. Zinc protects the iron since Zn is the anode and Fe the cathode: Zn 2+ (aq) +2e -  Zn(s), E  red = -0.76 V Fe 2+ (aq) + 2e -  Fe(s), E  red = -0.44 V With the above standard reduction potentials, Zn is easier to oxidize than Fe.

55. Chapter 2154 Corrosion Preventing the Corrosion of Iron

56. Chapter 2155 Corrosion Preventing the Corrosion of Iron To protect underground pipelines, a sacrificial anode is added. The water pipe is turned into the cathode and an active metal is used as the anode. Often, Mg is used as the sacrificial anode: Mg 2+ (aq) +2e -  Mg(s), E  red = -2.37 V Fe 2+ (aq) + 2e -  Fe(s), E  red = -0.44 V

57. Chapter 2156 Corrosion Preventing the Corrosion of Iron

58. Chapter 2157 Electrolysis Electrolysis of Aqueous Solutions Nonspontaneous reactions require an external current in order to force the reaction to proceed. Electrolysis reactions are nonspontaneous. In voltaic and electrolytic cells: –reduction occurs at the cathode, and –oxidation occurs at the anode. –However, in electrolytic cells, electrons are forced to flow from the anode to cathode. –In electrolytic cells the anode is positive and the cathode is negative. (In galvanic cells the anode is negative and the cathode is positive.)

59. Chapter 2158 Electrolysis Electrolysis of Aqueous Solutions

60. Chapter 2159 Electrolysis Electrolysis of Aqueous Solutions Example, decomposition of molten NaCl. Cathode: 2Na + (l) + 2e -  2Na(l) Anode: 2Cl - (l)  Cl 2 (g) + 2e -. Industrially, electrolysis is used to produce metals like Al. Electrolysis with Active Electrodes Active electrodes: electrodes that take part in electrolysis. Example: electrolytic plating.

61. Chapter 2160 Electrolysis Electrolysis with Active Electrodes

62. Chapter 2161 Electrolysis Electrolysis with Active Electrodes Consider an active Ni electrode and another metallic electrode placed in an aqueous solution of NiSO 4 : Anode: Ni(s)  Ni 2+ (aq) + 2e - Cathode: Ni 2+ (aq) + 2e -  Ni(s). Ni plates on the inert electrode. Electroplating is important in protecting objects from corrosion.

63. Chapter 2162 Electrolysis Quantitative Aspects of Electrolysis We want to know how much material we obtain with electrolysis. Consider the reduction of Cu 2+ to Cu. –Cu 2+ (aq) + 2e -  Cu(s). –2 mol of electrons will plate 1 mol of Cu. –The charge of 1 mol of electrons is 96,500 C (1 F). –Since Q = It, the amount of Cu can be calculated from the current (I) and time (t) taken to plate.

64. Chapter 2163 Electrolysis Electrical Work Free-energy is a measure of the maximum amount of useful work that can be obtained from a system. We know If work is negative, then work is performed by the system and E is positive. The emf can be thought about as a measure of the driving force for a redox process.

65. Chapter 2164 Electrolysis Electrical Work In an electrolytic cell and external source of energy is required to force the reaction to proceed. In order to drive the nonspontaneous reaction the external emf must be greater than E cell. From physics: work has units watts: 1 W = 1 J/s. Electric utilities use units of kilowatt-hours:

66. Chapter 2165 End of Chapter 20 Electrochemistry