Electrochemistry

1. Chapter 21 Electrochemistry

2. Electrochemistry the study of the production of electricity during chemical rxns and the changes produced by electrical current. Electrochemical reactions are oxidation-reduction reactions. Oilrig: oxidation loss of electrons, reduction gaining of electrons 1. Oxidation = loss of electrons a. the substance oxidized is the reducing agent 2. Reduction = gain of electrons a. the substance reduced is the oxidizing agent

3. Electrical Conduction • Metals conduct electric currents well. • metallic conduction • Positively charged ions, cations, move toward the negative electrode. • Negatively charged ions, anions, move toward the positive electrode.

4. There are two kinds electrochemical cells. Electrolytic cells - nonspontaneous chemical reactions Voltaic or galvanic cells - spontaneous chemical reactions The two parts of the reaction are physically separated. –oxidation occurs at one cell –reduction occurs in the other cell

5. Conventions for electrodes: Cathode - electrode at which reduction occurs (red cat) Anode - electrode at which oxidation occurs (an ox) Inert electrodes do not react with the liquids or products of the electrochemical reaction. Graphite and Platinum are common inert electrodes.

6. Electrolytic Cells Use electrical energy to force nonspontaneous(non thermodynamically favored) chemical reactions to occur. Process called electrolysis. Used in: –plating of jewelry and auto parts –electrolysis of chemical compounds Electrolytic cells consist of a: –container for reaction mixture –electrodes immersed in the reaction mixture –source of direct current

7. In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode). In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

8. Faraday’s Law of Electrolysis The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell. During electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent. Corresponds to the passage of one mole of electrons through the electrolytic cell.

9. Faraday Amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode. • 1 faraday of electricity = 6.022x1023 e- • 1 faraday = 6.022x1023 e- = 96487 coulombs • 1 eq. of oxidizing agent= gain of 6.022x1023 e- • 1 eq. of reducing agent = loss of 6.022x1023 e-

10. Ex. 1) Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes. (1 ampere = 1 coulomb per second)

11. Ex. 2) Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in Ex. 1.

12. Voltaic or Galvanic Cells • •Electrochemical cells in which a spontaneous chemical reaction produces electrical energy. • •In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode). • •Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference Examples: Car & flashlight batteries

13. The Construction of Simple Voltaic Cells • Half-cell contains the oxidized and reduced forms of an element (or other chemical species) in contact with each other. • Simple cells consist of: • two pieces of metal immersed in solutions of their ions • wire to connect the two half-cells • salt bridge to • complete circuit • maintain neutrality • prevent solution mixing

14. The Zinc-Copper Cell • Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Zn strip immersed in 1.0 M zinc (II) sulfate wire and a salt bridge to complete circuit • Initial voltage is 1.10 volts

15. The Zinc-Copper Cell

16. The Zinc-Copper Cell • In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

17. The Zinc-Copper Cell • Short hand notation for voltaic cells • Zn-Cu cell example

18. The Copper - Silver Cell • Cell components: Cu strip immersed in 1.0 M copper (II) sulfate Ag strip immersed in 1.0 M silver (I) nitrate wire and a salt bridge to complete circuit • Initial voltage is 0.46 volts

19. The Copper - Silver Cell

20. The Copper - Silver Cell • Compare the Zn-Cu cell to the Cu-Ag cell Cu electrode is cathode in Zn-Cu cell Cu electrode is anode in Cu-Ag cell • Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

21. The Copper - Silver Cell • Demonstrates that Cu2+ is a stronger oxidizing agent than Zn2+ Cu2+ oxidizes metallic Zn to Zn2+ • Ag+ is is a stronger oxidizing agent than Cu2+ Ag+ oxidizes metallic Cu to Cu2+ • Arrange these species in order of increasing strengths

22. Standard Electrode Potential • Establish an arbitrary standard to measure potentials of a variety of electrodes • Standard Hydrogen Electrode (SHE) • assigned an arbitrary voltage of 0.000000… V

23. The Electromotive (Activity) Series of the Elements • Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials. • Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials. • Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.

24. Uses of the Electromotive Series • Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written. • For example, the half-reaction for the standard potassium electrode is: • The large negative value tells us that this reaction will occur only under extreme conditions.

25. Uses of the Electromotive Series • Compare the potassium half-reaction to fluorine’s half-reaction: • The large positive value denotes that this reaction occurs readily as written. • Positive E0 values tell us that the reaction tends to occur to the right • larger the value, greater tendency to occur to the right • Opposite for negative values

26. Uses of the Electromotive Series 1. Choose the appropriate half-reactions from a table of standard reduction potentials. 2. Write the equation for the half-reaction with the more positive E0 value first, along with its E0 value. 3. Write the eqn for the other half-reaction as an oxidation with its oxidation potential, reverse the tabulated reduction half-reaction and change the sign of the tabulated E0. 4. Balance the electron transfer. 5. Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous.

27. Uses of the Electromotive Series • Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously. • Ex. 3) Will silver ions, Ag+, oxidize metallic zinc to Zn2+ ions, or will Zn2+ ions oxidize metallic Ag to Ag+ ions? What is the overall value for Eo? • E0 values are not multiplied by any stoichiometric relationships in this procedure.

28. Electrode Potentials for Other Half-Reactions • Ex.4) Will tin(IV) ions oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize tin(II) ions to tin(IV) ions in acidic solution? What is the overall value for rxn?

29. Effect of Conc. (or Partial Pressures) on Electrode Potentials - Nernst Equation • Standard electrode potentials are determined at thermodynamic standard conditions. 1 M solutions 1 atm of pressure for gases liquids and solids in their standard states temperature of 250C • Potentials change if conditions are nonstandard. • Nernst equation describes the electrode potentials at nonstandard conditions.

30. Effect of Conc. (or Partial Pressures) on Electrode Potentials - Nernst Equation

31. Substitution of the values of the constants into the Nernst equation at 250 C gives:

32. The Nernst Equation • used to calculate electrode potentials and cell potentials for concentrations other than standard –state values. (pg 877) • Ex. 5) Calculate the cell potential for the following electrochemical cell. If the [Sn2+] = 4.5 x 10-1 M and [Ag+] = 0.110 M • Sn(s) + 2Ag+ Sn2+ + 2Ag(s)

33. The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes. • Ex. 6) Calculate the initial potential of a cell that consists of an Fe3+/Fe2+ electrode in which [Fe3+]=1.0 x 10-2M and [Fe2+]=0.1 M connected to a Sn4+/Sn2+ electrode in which [Sn4+]=1.0 M and [Sn2+]=0.10 M . A wire and salt bridge complete the circuit.

34. Relationship of E0cell to DG0 and K • From previous chapters we know the relationship of DG0 and K for a reaction.

35. Relationship of E0cell to DG0 and K • The relationship between DG0 and E0cell is also a simple one.

36. Relationship of E0cell to DG0 and K • You can combine these two relationships into a single relationship to relate E0cell to K.

37. Relationship of E0cell to DG0 and K • Ex. 7) Calculate the standard Gibbs free energy change, DG0 , at 250C for the following reaction.

38. Relationship of E0cell to DG0 and K • Calculate E0cell using the appropriate half-reactions.

39. Relationship of E0cell to DG0 and K • Now that we know E0cell , we can calculate DG0 . • The negative value tells us that the reaction is spontaneous as written.

40. Relationship of E0cell to DG0 and K • Ex. 8) Calculate the thermodynamic equilibrium constant for the reaction in Ex. 7 at 250C.

41. Corrosion • Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO2, O2, and H2O.

42. Corrosion = Oxidation of a metal • The oxidation of most metals by oxygen is spontaneous. Many metals develop a thin coating of metal oxide on the outside that prevents further oxidation • The presence of a salt accelerates the corrosion process by increasing the ease with which electrons are conducted from anodic to cathodic regions

43. Corrosion Protection • Some examples of corrosion protection. • Plate a metal with a thin layer of a less active (less easily oxidized) metal. 2. Connect the metal to a sacrificial anode, a piece of a more active metal.

44. Corrosion Protection • Allow a protective film to form naturally.

45. Corrosion Protection • Galvanizing, coating steel with zinc, a more active metal. 5. aint or coat with a polymeric material such as plastic or ceramic.

46. Primary Voltaic Cells • As a voltaic cell discharges, its chemicals are consumed. • Once chemicals are consumed, further chemical action is impossible. • Electrodes and electrolytes cannot be regenerated by reversing current flow through cell.

47. The Dry Cell (LeClanche’ Cell) • One example is flashlight, radio, etc. batteries. • Container is made of zinc • acts as an electrode • Graphite rod is in center of cell • acts as the other electrode • Space between electrodes is filled with a mixture of: • ammonium chloride, NH4Cl • manganese (IV) oxide, MnO2 • zinc chloride, ZnCl2 • porous inactive solid

48. The Dry Cell (LeClanche’ Cell) • As current is produced, Zn dissolves and goes into solution as Zn2+ ions. • Zn electrode is negative (anode).

49. Secondary Voltaic Cells • Secondary cells are reversible, rechargeable. • Electrodes can be regenerated • One example is the lead storage or car battery.

50. Lead Storage Battery • Electrodes are two sets of lead alloy grids (plates). • Holes in one grid are filled with lead (IV) oxide, PbO2. • Other holes are filled with spongy lead. • Electrolyte is dilute sulfuric acid. • When battery is discharging, spongy lead is oxidized to lead ions and the plate becomes negatively charged.