1. Chem. 1B – 11/8 Lecture

2. Announcements Today’s Lecture Electrochemistry
The Nernst Equation (for non-standard conditions, including determining Ksp and Kf values)
Batteries – applied voltaic cells
Electrolytic Cells

3. Chapter 18 Electrochemistry
Nernst Equation Application – Cont.
2nd Examples:
The following cell,
Cd(s)|CdC2O4(s)|C2O42-(aq)||Cu2+(aq)1 M|Cu(s) is used to determine [C2O42-]. If Eº for the Cd reaction is V (reduction potential for oxidation reaction) and Eº for the Cu reaction is V, and the measured voltage is V, what is [C2O42-]?

4. Chapter 18 Electrochemistry
Nernst Equation Application – Cont.
3rd 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/Kf for Cu(NH3)42+ and Ksp for AgCl)

5. Chapter 18 Electrochemistry
Nernst Equation Application – Cont.
3rd Type – Specific Example:
Using information given from the following cell and standard potentials, determine the Ksp for Hg2Cl2(s)
Hg(l)|Hg2Cl2(s)|Hg22+(X M), 0.10 M NaCl||1.0 M Cu2+|Cu(s)
Above cell has E = 0.01 V , E°(Hg22+ reduction) = 0.80 V, and E°(Cu2+ reduction) = 0.34 V
Calculate Ksp and E° for Hg2Cl2(s) reduction.

6. 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)

7. Chapter 18 Electrochemistry
Voltaic Cells – Batteries
Examples:
Lead-Acid (standard Car battery)
Pb used in both oxidation and reduction to Pb2+ (most stable form)
Oxidation: Pb(s) + HSO4-(aq) ↔ PbSO4(s) + H+(aq) + 2e-
Reduction: PbO2(s) + HSO4-(aq) + 3H+(aq) + 2e- ↔ PbSO4(s) + H2O(l)

8. Chapter 18 Electrochemistry Batteries
Examples:
Lead-Acid (standard Car battery)
Net: Pb(s) + PbO2(s) + 2HSO4-(aq) + 2H+(aq) ↔ PbSO4(s) + 2H2O(l)
So E = Eº – /2log[1/([HSO4-]2[H+]2)]
As reaction proceed, voltage slowly drops (faster when nearly depleted) – due to loss of HSO4-, H+
Rechargeable
Low Energy Density (Pb is heavy)
12 V from 6 cells

9. Chapter 18 Electrochemistry Batteries
Examples:
Alkaline batteries
Net: Zn(s) + 2MnO2(s) + 2H2O(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

10. Chapter 18 Electrochemistry Batteries - Questions
If 100 g. of Pb(s) is used in a lead acid battery, what mass of PbO2(s) is required for best efficiency?
How many Amp–Hours will this provide in a 12 V battery?
Will the voltage generated by a lead acid battery be affected by how “depleted” it is?
What is the voltage when it is 90% depleted (vs. initial voltage)?

11. 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 available (reported last year in Sacramento Bee)
Reactions: 2H2(g) + 4OH-(aq) ↔ 4H2O(l) + 4e-
And O2(g) + 2H2O(l) + 4e- ↔ 4OH-(aq)
If H2 is produced from electrolysis using solar energy, 100% renewable
More commonly, H2 is made from natural gas

12. 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 (C6H12O6(aq) + O2(aq) ↔ C6H10O6(aq) + H2O2(aq) - requires enzyme)
Either attachment of enzyme to electrode or electrodes for H2O2 oxidation or reduction can be used to generate electricity

13. Chapter 18 Electrochemistry Electrolytic Cells
Example Reactions
Electrolysis of water (opposite of fuel cell example)
Anode: H2O – oxygen is oxidized to O2(g)
Cathode: H2O – hydrogen is reduced to H2(g)
Industrial Use – Electroplating (Chrome, nickel, silver plating possible) – using external potential to deposit metal to electrode

14. Chapter 18 Electrochemistry Electrolytic Cells
Example Reactions
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 = V
H2O(l) + 2e- ↔ H2(g) + 2OH-(aq) Eº = V
2Cl-(aq) ↔ Cl2(g) + 2e- Eº = V
2I-(aq) ↔ I2(aq) + 2e- Eº = V
2H2O(l) ↔ O2(g) + 4H+(aq) + 4e- Eº = 1.23 V

15. Chapter 18 Electrochemistry Electrolytic Cells - Questions
Which of the following changes in switching from a voltaic to an electrolytic cell?
Charge on anode/cathode
Which electrode (e.g. anode) does oxidation/reduction
Ion migration to electrode
An anode in an electrolytic cell is used to measure oxalate (Eº = -0.49V) in the presence of pyruvate (Eº = -0.70V). Which will oxidize first in a mixture?

16. Chapter 18 Electrochemistry Electrolytic Cells – Questions – Cont.
When a battery is being recharged, which of the following happens?
Electrode charge changes
Electrode reaction changes (anode – oxidation becomes cathode)
Used as voltaic cell – battery powers light bulb +
External Power Generator – to charge battery - + -
Pb(IV) → Pb(II)
Pb(II) → Pb(IV)
Reactions reverse (products to reactants to replenish charge)

17. Chapter 18 Electrochemistry Electrolytic Cells – Questions – cont.
A steel rod is being chrome plated. If the rod has a diameter of 10 cm and a length of 150 cm, how long does it take to chrome plate the rod to a thickness of 1 mm, if a current of 20 A is applied during the chrome plating process?
V(rod) = pd2L/4 and AW(Cr) = g/mol