Chapter 17 Electrochemistry

1. Chapter 17Electrochemistry

2. Voltaic Cells • In spontaneous reduction-oxidation reactions, electrons are transferred and energy is released. • The energy released can be used to do work if the electrons flow through an external pathway, such as a voltaic or galvanic cell.

3. Anode

4. Cathode

5. Salt bridge-is a U-shaped tube that contains an electrolytic solution whose ions will not react with the ions in the cell or with the electrode material.

6. Molecular Perception • At the anode, electrons leave (driven or pushed) to flow through the wire to the cathode. • As the electrons leave the anode the cations formed dissolve into the solution in the anode compartment.

7. Molecular Perception • As the electrons reach the cathode, cations in the cathode compartment solution are attracted to the negative cathode. • The electrons are taken by the cation, reducing the ion to a neutral atom that is deposited on the cathode.

8. Sample Problem The two half-reactions in a voltaic cell are (a) Indicate which reaction occurs at the anode and which at the cathode. (b) Which electrode is consumed in the cell reaction? (c) Which electrode is positive?

9. Sample Problem The oxidation-reduction reaction is spontaneous. A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with either solution (such as platinum foil) is suspended in each solution, and the two conductors are connected with wires through a voltmeter or some other device to detect an electric current. The resultant voltaic cell generates an electric current. Indicate the reaction occurring at the anode, the reaction at the cathode, the direction of electron migration, the direction of ion migration, and the signs of the electrodes.

10. Cell Potential • What causes water to flow over a waterfall? • The flow of electrons from the anode to the cathode is spontaneous. • Electromotive force (emf, cell potential, cell voltage) is the difference in electrical potential energy between anode and the cathode.

11. Standard Reduction Potentials • By convention, the potentials for many electrodes have been measured and tabulated as standard reduction potentials. • All standard reduction potentials are referenced to a standard hydrogen electrode (SHE) or normal hydrogen electrode (NHE).

12. Standard Cell Potentials • The standard cell potential can be found by • Consider the spontaneous reaction between Zn and H1+. E°cell = E°red(cathode) - E°red(anode)

13. Consider the reaction between Zn and Cu2+, the cell potential can be calculated by using the equation.

14. Sample Problem A voltaic cell is based on the half-reactions The standard emf for this cell is 1.46 V. Using the standard reduction potentials, calculate E°red for the reduction of In3+ to In+.

15. Sample Problem Using the standard reduction potentials listed in the standard reduction potentials, calculate the standard emf for the voltaic cell made based on the reaction

16. Strengths of Oxidizing/Reducing Agents • The strongest oxidizers have the most positive reduction potentials. • The strongest reducers have the most negative reduction potentials.

17. Strengths of Oxidizing/Reducing Agents • Use the standard reduction potential table to predict if one reactant can spontaneously oxidize/reduce another. • The greater the difference between the two reduction potentials the greater the voltage of the cell.

18. Sample Problem Rank the following ions in order of increasing strength as oxidizing agents: NO3–(aq), Ag+(aq), Cr2O72–(aq). Rank the following species from the strongest to the weakest reducing agent: I–(aq), Fe(s), Al(s).

19. Sample Problem A voltaic cell is based on the following two standard half-reactions Determine (a) the half-reactions that occur at the cathode and the anode, and (b) the standard cell potential.

20. Sample Problem A voltaic cell is based on a Co2+/Co half-cell and an AgCl/Ag half-cell. (a) What reaction occurs at the anode? (b) What is the standard cell potential?

21. Activity Series • For any electrochemical process the cell potential can be found by E°cell= E°red(reduction) - E°red(oxidation) • Equation can be used to understand the activity series of metals.

22. Sample Problem Using standard reduction potentials, determine whether the following reactions are spontaneous under standard conditions.

23. Free Energy and EMF • The ΔG of a redox reaction can be determined by ΔG = - nFE • A negative value of ΔG and a positive E indicates a spontaneous reaction. • Under standard conditions, the standard emf is related to the equilibrium constant for a reaction through ΔG.

24. Sample Problem For the reaction (a) What is the value of n? (b) Use the data in Appendix E to calculate G°. (c) Calculate K at T = 298 K.

25. RT nF E = E − lnQ 0.0592 n log Q E = E − Cell EMF Under Nonstandard Conditions • Cell EMF at nonstandard conditions is determined using the Nernst Equation • The equation may be simplified by combining all the constants and using a temperature of 298 K to get

26. Sample Problem Calculate the emf at 298 K generated by the cell described below, when [Cr2O72–] = 2.0 M, [H+] = 1.0 M, [I–] = 1.0 M, and [Cr3+] = 1.0 10–5M.

27. Sample Problem What is the pH of the solution in the cathode compartment of the cell pictured below, when PH2 = 1.0 atm, [Zn2+] in the anode compartment is 0.10 M, and cell emf is 0.542 V?

29. Concentration Cells • A cell whose emf is generated because of a concentration difference.

30. Sample Problem A concentration cell is constructed with two Zn(s)-Zn2+(aq) half-cells. The first half-cell has [Zn2+] = 1.35 M, and the second half-cell has [Zn2+] = 3.75  10–4M. (a) Which half-cell is the anode of the cell? (b) What is the emf of the cell?

31. Electrolysis • Electrolysis reactions are nonspontaneous reactions that require an external current in order to force the reaction to proceed. • In voltaic and electrolytic cells, the role of the electrodes are the same; reduction at the cathode and oxidation at the anode. • In electrolytic cells, an external source of electrical current is necessary to drive the reaction forward. This source pumps electrons into one electrode (cathode) and pulls them out of the other (anode).

32. Electrolytic Cell

33. Electroplating • An electrolytic process that deposits a thin layer of one metal on another metal in order to improve its characteristics, such as appearance or resistance to corrosion.

34. Quantitative Aspects of Electrolysis • Determination of the amount of material produced in electrolysis. • The amount of substance that is reduced or oxidized in an electrolytic cell is directly proportional to the number of electrons (quantity of charge) passed into the cell.

35. The quantity of the charge passing through an electrical circuit may be found by I = q/t

36. Sample Problem Calculate the number of grams of aluminum produced in 1.00 h by the electrolysis of molten AlCl3 if the electrical current is 10.0 A.

37. Sample Problem (a) The half-reaction for formation of magnesium metal upon electrolysis of molten MgCl2 is Mg2+ + 2 e- Mg. Calculate the mass of magnesium formed upon passage of a current of 60.0 A for a period of 4.00 103s. (b) How many seconds would be required to produce 50.0 g of Mg from MgCl2 if the current is 100.0 A?

38. Electrical Work • Free energy is a measure of the maximum useful work that can be obtained from a spontaneous system. • The emf of a cell is a way of measuring the driving force for a redox reaction. • In voltaic cells, the system is able to do work on its surroundings.

39. In electrolytic cells, the system must have work done on it by the surroundings to cause the redox reaction. • To force the process to occur, an external potential (Eext) must be larger than the cell emf. • The amount of work done on the system is given by

40. Sample Problem Calculate the number of kilowatt-hours of electricity required to produce 1.0  103 kg of aluminum by electrolysis of Al3+ if the applied voltage is 4.50 V.

41. Sample Problem Calculate the number of kilowatt-hours of electricity required to produce 1.00 kg of Mg from electrolysis of molten MgCl2 if the applied emf is 5.00 V. Assume that the process is 100% efficient.