1. Chem. 1B – 11/10 Lecture

2. Announcements Mastering Chemistry –Chapter 18 Assignment is due 11/17 Today’s Lecture – Electrochemistry (Ch. 18) –More Nernst Equation Applications –Specific Applications of Voltaic Cells –Electrolytic Cells –Corrosion

3. Chapter 18 Electrochemistry Nernst Equation Application Example: Determine the voltage for a Ag(s)/AgCl(s) electrode when [Cl - ] = 0.010 M if Eº = 0.222 V (at T = 25°C)? Note: this is the same as when this electrode is attached to a SHE.

4. Chapter 18 Electrochemistry Nernst Equation Application – Cont. 2 nd Example: –The following cell, Cd(s)|CdC 2 O 4 (s)|C 2 O 4 2- (aq)||Cu 2+ (aq)1 M|Cu(s) is used to determine [C 2 O 4 2- ]. If Eº for the Cd reaction is -0.522 V (reduction potential for oxidation reaction) and Eº for the Cu reaction is +0.337 V, and the measured voltage is 0.647 V, what is [C 2 O 4 2- ]?

5. Chapter 18 Electrochemistry Nernst Equation Application – Cont. 3 rd Type of Example: –While metal electrodes are commonly used, it is also possible to combine metal reactions with other aqueous phase metal ion reactions. For example, in the lab, we use electrochemical measurements to determine equilibrium constants (1/K f for Cu(NH 3 ) 4 2+ and K sp for AgCl)

6. Chapter 18 Electrochemistry Nernst Equation Application – Cont. 3 rd Type – Specific Example: –Using information given from the following cell and standard potentials, determine the K sp for Hg 2 Cl 2 (s) Hg(l)|Hg 2 Cl 2 (s)|Hg 2 2+ (X M), 0.10 M NaCl||1.0 M Cu 2+ |Cu(s) Above cell has E = 0.01 V, E°(Hg 2 2+ reduction) = 0.80 V, and E°(Cu 2+ reduction) = 0.34 V Calculate K sp and E° for Hg 2 Cl 2 (s) reduction.

7. Chapter 18 Electrochemistry Voltaic Cells – Batteries Requirements –General interest is in high energy density (large E° per unit mass or volume) –To get high E values, more highly reactive reagents are desired (provided package is stable) –Rechargeable batteries When switching from voltaic cell (using battery) to electrolytic cell (charging battery), reverse reaction must occur (vs. producing other products)

8. Chapter 18 Electrochemistry Voltaic Cells – Batteries Examples: –Lead-Acid (standard Car battery) Pb used in both oxidation and reduction to Pb 2+ (most stable form) Oxidation: Pb(s) + HSO 4 - (aq) ↔ PbSO 4 (s) + H + (aq) + 2e - Reduction: PbO 2 (s) + HSO 4 - (aq) + 3H + (aq) + 2e - ↔ PbSO 4 (s) + H 2 O(l)

9. Chapter 18 Electrochemistry Batteries Examples: –Lead-Acid (standard Car battery) Net: Pb(s) + PbO 2 (s) + 2HSO 4 - (aq) + 2H + (aq) ↔ PbSO 4 (s) + 2H 2 O(l) So E = Eº – 0.0592/2log[1/([HSO 4 - ] 2 [H + ] 2 )] As reaction proceed, voltage slowly drops (faster when nearly depleted) Rechargeable Low Energy Density (Pb is heavy) 12 V from 6 cells

10. Chapter 18 Electrochemistry Batteries Examples: –Alkaline batteries Net: Zn(s) + 2MnO 2 (s) + 2H 2 O(l) ) ↔ Zn(OH) 2 (s) + 2MnO(OH)(s) Since all reactants and products are solids, E = Eº = constant Better if V must stay constant, but harder to tell if battery is near end of lifetime Non-rechargeable

11. Chapter 18 Electrochemistry Batteries - Questions 1.If 100 g. of Pb(s) is used in a lead acid battery, what mass of PbO 2 (s) is required for best efficiency? 2.How many Amp–Hours will this provide in a 12 V battery? 3.Will the voltage generated by a lead acid battery be affected by how “depleted” it is? 4.What is the voltage when it is 90% depleted (vs. initial voltage)?

12. Chapter 18 Electrochemistry Other Voltaic Cells Examples: –Fuel Cells Voltaic cell where reactants (fuel plus oxygen) flow to electrodes to produce electricity New Toyota Fuel Cell vehicle now available (reported recently in Sacramento Bee) Reactions: 2H 2 (g) + 4OH - (aq) ↔ 4H 2 O(l) + 4e - And O 2 (g) + 2H 2 O(l) + 4e - ↔ 4OH - (aq) If H 2 is produced from electrolysis using solar energy, 100% renewable More commonly, H 2 is made from natural gas

13. Chapter 18 Electrochemistry Other Voltaic Cells Examples: –Powering Medical Devices Batteries have a limited lifetime, so pacemakers and defibrillators must be surgically removed to replace batteries Another option is to run devices off of blood glucose oxidation (C 6 H 12 O 6 (aq) + O 2 (aq) ↔ C 6 H 10 O 6 (aq) + H 2 O 2 (aq) - requires enzyme) Either attachment of enzyme to electrode or electrodes for H 2 O 2 oxidation or reduction can be used to generate electricity

14. Chapter 18 Electrochemistry Electrolytic Cells Differences with Voltaic Cells –Uses External voltage to drive unfavorable reaction –Charge at electrodes is reversed anode (note: oxidation driven by voltage, but now + charge) cathode (reduction, - charge) Power Supply + -

15. Chapter 18 Electrochemistry Electrolytic Cells Example Reactions 1.Electrolysis of water (opposite of fuel cell example) Anode: H 2 O – oxygen is oxidized to O 2 (g) Cathode: H 2 O – hydrogen is reduced to H 2 (g) 2.Industrial Use – Electroplating (Chrome, nickel, silver plating possible) – using external potential to deposit metal to electrode

16. Chapter 18 Electrochemistry Electrolytic Cells Example Reactions 3.Electrolysis of Mixtures – e.g. analysis External potential will work on easiest to oxidize/reduce pair For example, if we have a mixture of NaI and NaCl in water, electrolysis will cause the following reactions: –Na + (aq) + e - ↔ Na(s) Eº = -2.71 V –H 2 O(l) + 2e - ↔ H 2 (g) + 2OH - (aq) Eº = -0.83 V –2Cl - (aq) ↔ Cl 2 (g) + 2e - Eº = +1.36 V –2I - (aq) ↔ I 2 (aq) + 2e - Eº = +0.54 V –2H 2 O(l) ↔ O 2 (g) + 4H + (aq) + 4e - Eº = 1.23 V

17. Chapter 18 Electrochemistry Electrolytic Cells - Questions 1.Which of the following changes in switching from a voltaic to an electrolytic cell? a)Charge on anode/cathode b)Which electrode (e.g. anode) does oxidation/reduction c)Ion migration to electrode 2.An anode in an electrolytic cell is used to measure oxalate (Eº = -0.49V) in the presence of pyruvate (Eº = -0.70V). Can this be done? What if one is interested in pyruvate?

18. Chapter 18 Electrochemistry Corrosion Most metals are more stable as oxides, so oxidation of metals is common One would think that oxidation is primarily dependent upon Eº values, but oxides like Al 2 O 3 can protect further oxidation of Al metal Iron in particular is prone to rusting Galvanized metal uses a more readily oxidized metal (e.g. Zn) to protect iron

19. Chapter 18 Electrochemistry Corrosion - Questions Using the table below, which metals can be added as a sacrificial agent to prevent iron oxidation? ReactionEº (V) Ag + (aq) + e - ↔ Ag(s)+0.799 Cu 2+ (aq) + 2e - ↔ Cu(s)+0.337 Co 2+ (aq) + 2e - ↔ Co(s)-0.277 Fe 2+ (aq) + 2e - ↔ Fe(s)-0.45 Cr 3+ (aq) + 3e - ↔ Cr(s)-0.73 Mn 2+ (aq) + 2e - ↔ Mn(s)-1.18