Unit 12 Electrochemistry Day 2

1. Lecture Presentation Unit 12Electrochemistry Day 2

2. Warm Up SHOW ME: Cornell Notes, Video Notes, and Practice Problems FIND: Balancing Redox Reactions Worksheet 1 at desk and take out notebook paper to work on TIME:4 minutes WHEN DONE: Get calculator from blue bin

3. Agenda • Review Balancing Redox Rxn Practice (12 min) • Voltaic Cells (12 min) • Cell Potential (9 min) • Multiple Choice Questions • Pass Back Work and Work Time: ChemActivities and Set Up Lab

4. Voltaic Cells In spontaneous redox reactions, electrons are transferred and energy is released. That energy can do work if the electrons flow through an external device. This is a voltaic cell.

5. Voltaic Cells The oxidation occurs at the anode. The reduction occurs at the cathode. When electrons flow, charges aren’t balanced. So, a salt bridge, usually a U-shaped tube that contains a salt/agar solution, is used to keep the charges balanced.

6. Voltaic Cells In the cell, electrons leave the anode and flow through the wire to the cathode. Cations are formed in the anode compartment. As the electrons reach the cathode, cations in solution are attracted to the now negative cathode. The cations gain electrons and are deposited as metal on the cathode.

7. Electromotive Force (emf) Water flows spontaneously one way in a waterfall. Comparably, electrons flow spontaneously one way in a redox reaction, from high to low potential energy.

8. Electromotive Force (emf) The potential difference between the anode and cathode in a cell is called the electromotive force (emf). It is also called the cell potential and is designated Ecell. It is measured in volts (V). One volt is one joule per coulomb (1 V = 1 J/C).

9. Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated. The values are compared to the reduction of hydrogen as a standard.

10. Standard Hydrogen Electrode Their reference is called the standard hydrogen electrode (SHE). By definition as the standard, the reduction potential for hydrogen is 0 V: 2 H+(aq, 1M) + 2e– H2(g, 1 atm)

11. Standard Cell Potentials The cell potential at standard conditions can be found through this equation: ° ° ° Ecell = Ered (cathode) –Ered (anode) Because cell potential is based on the potential energy per unit of charge, it is an intensive property.

12. Cell Potentials For the anode in this cell, E°red = –0.76 V For the cathode, E°red= +0.34 V So, for the cell, E°cell= E°red (anode) – E°red (cathode) = +0.34 V – (–0.76 V) = +1.10 V

13. Oxidizing and Reducing Agents The more positive the value of E°red, the greater the tendency for reduction under standard conditions. The strongest oxidizers have the most positive reduction potentials. The strongest reducers have the most negative reduction potentials.

14. Free Energy and Redox Spontaneous redox reactions produce a positive cell potential, or emf. E° = E°red (reduction) – E°red (oxidation) Note that this is true for ALL redox reactions, not only for voltaic cells. Since Gibbs free energy is the measure of spontaneity, positive emf corresponds to negative ΔG. How do they relate? ΔG= –nFE (F is the Faraday constant, 96,485 C/mol.)

15. Free Energy, Redox, and K How is everything related? ΔG°= –nFE°= –RT ln K

16. Nernst Equation Remember, ΔG = ΔG°+ RT ln Q So, –nFE = nFE°+ RT ln Q Dividing both sides by –nF, we get the Nernst equation: E= E°– (RT/nF) ln Q OR E = E°– (2.303 RT/nF) log Q Using standard thermodynamic temperature and the constants R and F, E = E°– (0.0592/n) log Q

17. Concentration Cells Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes, called a concentration cell. Ecell ° • For such a cell, would be 0, but Q would not. • Therefore, as long as the concentrations are different, E will not be 0.

18. Some Applications of Cells • Electrochemistry can be applied as follows: • Batteries: a portable, self-contained electrochemical power source that consists of one or more voltaic cells. • Batteries can be primary cells (cannot be recharged when “dead”—the reaction is complete) or secondary cells (can be recharged). • Prevention of corrosion (“rust-proofing”) • Electrolysis

19. Some Examples of Batteries • Lead–acid battery: reactants and products are solids, so Q is 1 and the potential is independent of concentrations; however, made with lead and sulfuric acid (hazards). • Alkaline battery: most common primary battery. • Ni–Cd and Ni–metal hydride batteries: lightweight, rechargeable; Cd is toxic and heavy, so hydrides are replacing it. • Lithium-ion batteries: rechargeable, light; produce more voltage than Ni-based batteries.

20. Some Batteries Lead–Acid Battery Alkaline Battery

21. Lithium-Ion Battery

22. Fuel Cells • When a fuel is burned, the energy created can be converted to electrical energy. • Usually, this conversion is only 40% efficient, with the remainder lost as heat. • The direct conversion of chemical to electrical energy is expected to be more efficient and is the basis for fuel cells. • Fuel cells are NOT batteries; the source of energy must be continuously provided.

23. Hydrogen Fuel Cells • In this cell, hydrogen and oxygen form water. • The cells are twice as efficient as combustion. • The cells use hydrogen gas as the fuel and oxygen from the air.

24. Corrosion • Corrosion is oxidation. • Its common name is rusting.

25. Preventing Corrosion • Corrosion is prevented by coating iron with a metal that is more readily oxidized. • Cathodic protection occurs when zinc is more easily oxidized, so that metal is sacrificed to keep the iron from rusting.

26. Preventing Corrosion • Another method to prevent corrosion is used for underground pipes. • A sacrificial anode is attached to the pipe. The anode is oxidized before the pipe.

27. Electrolysis Nonspontaneous reactions can occur in electrochemistry IF outside electricity is used to drive the reaction. Use of electrical energy to create chemical reactions is called electrolysis.

28. Electrolysis and “Stoichiometry” • 1 coulomb = 1 ampere × 1 second • Q = It = nF • Q = charge (C) • I = current (A) • t = time (s) • n = moles of electrons thattravel through the wire inthe given time • F = Faraday’s constant NOTE: n is different than thatfor the Nernst equation!

29. The oxidation state of nitrogen in the ammonium ion (NH41+) is _______. +1 0 −1 −3

30. The oxidation state of manganese in the permanganate ion (MnO41−) is _______. a. −1 +2 +4 d. +7

31. Zn + Cu2+ Zn2+ + CuThe reducing agent in the reaction above is _______. Zn Cu2+ Zn2+ d. Cu

32. Zn + Cu2+ Zn2+ + CuThe oxidizing agent in the reaction above is _______. Zn Cu2+ Zn2+ d. Cu

33. To balance the half-reactionMnO4− Mn2+ in acidic solution, ___ electrons must be added on the ___ side. 5; product 2; product 7; reactant 5; reactant

34. MnO4− Mn2+ To balance this reactionin acidic solution, ___ H+ must be added on the ___ side. 8; product 4; product 8; reactant 4; reactant

35. If the value of the standard cell potential for a reaction is large and positive, then the reaction is at equilibrium. spontaneous. nonspontaneous. d. very fast.

36. The purpose of the salt bridge in a voltaic cell is to provide H+ ions needed to balance charges. maintain neutrality by allowing the flow of ions. serve as the site for oxidation to occur. serve as the site for reductionto occur.

37. Which of the following is the strongest oxidizing agent? (If necessary, consult a table of standard reduction potentials.) F2 Cl2 Br2 d. I2

38. Which of the following is the strongest reducing agent? (If necessary, consult a table of standard reduction potentials.) Zn Al Na d. Li

39. Cl2 2 Cl−εred= + 1.36 VI2 2 I−εred= + 0.54 VSelect the true statement. Cl2 will reduce I−to I2. Cl2 will oxidize I−to I2. I2 will reduce Cl−to Cl2. I2 will oxidize Cl2 to Cl−.

40. G = Gibbs free energy. n = moles of electrons. F = Faraday constant. E = cell potential. Change in G = ? F + nE F −nE c.−nFE d. nF−E

41. The Nernst Equation is most useful for determining cell potentials when _______ are nonstandard. oxidizing agents reducing agents ion concentrations d. temperatures

42. In a concentration cell, the half-reactions are the same. at equilibrium. acid–base reactions. d. different colors.

43. The electrolyte in a lead-acid automobile battery is _______. Pb PbO2 PbSO4 d. H2SO4

44. A cell that uses external energy to produce an oxidation–reduction reaction is called _______ cell. a galvanic a voltaic an electrolytic d. a prison

45. Reduction occurs at the anode, in both galvanic and electrolytic cells. cathode, in both galvanic and electrolytic cells. anode in galvanic cells and cathode in electrolytic cells. cathode in galvanic cells and anode in electrolytic cells.

46. Oxidation occurs at the anode, in both galvanic and electrolytic cells. cathode, in both galvanic and electrolytic cells. anode in galvanic cells and cathode in electrolytic cells. d. cathode in galvanic cells and anode in electrolytic cells.

47. Q = charge in coulombsI = current in amperest = time in secondsWhich is true? Q = I + t Q = I −t Q = It d. Q = I / t

48. Corrosion of metals can be prevented by all of the following methods except a sacrificial anode. a salt bridge. formation of an oxide coating. paint.