1. Chem. 133 – 2/23 Lecture

2. Announcements Exam 1 coming up (Mar. 7th)
Friday’s Seminar – Gallo Winery (Internship opportunity)
Today’s Lecture - Electrochemistry
Redox Reactions
Fundamental Equations (relating moles of electrons to charge and DG to electrical energy
Galvanic Cells
Standard Potential
Nernst Equation (if time)

3. Electrochemistry Redox Reactions
Reduction = loss of charge
e.g. Fe3+ + e- → Fe2+
Oxidation = gain in charge
e.g. Pb2+ + 2H2O → PbO2(s) + 4H+ + 2e- (Pb goes from +2 to +4)
Balancing reactions
review steps in general chemistry book
example: Zn(s) + Cr2O72- → Zn2+ + Cr3+
note: based on methods used for problems in this book, full cell balancing may not be needed

4. Electrochemistry Fundamental Equations
Relationship between charge, energy and current
redox reactions involve the exchange of electrons
when the exchange occurs on an electrode surface, current can be measured
Total charge transfer = q = nF, where
n = moles of electrons in reaction and
F = Faraday’s constant = C/moles e-
F = NAvogadro·e (e = elementary charge = 1.6 x C)
Current Produced = I = dq/dt
or q = ∫I·dt (or = I·t under constant current conditions)
can be used to determine battery lifetime

5. Electrochemistry Fundamental Equations
Relationship between charge, energy and current (continued)
Electrical work (units = J) = E·q (E = potential in volts and q in C)
and ΔG = -E·q = -nFE
under standard conditions (1 M reactant/product conc., 298K, etc.), ΔGº = -nFEº
ΔGº are given in Tables and allows calculation of K values
Eº, standard reduction potential, also given in Tables (see Appendix H), but for “half-reactions”

6. Electrochemistry Fundamental Equations
Example problem:
A NiCad battery contains 12.0 g of Cd that is oxidized to Cd(OH)2. How long should the battery last if a motor is drawing 421 mA? Assume 100% efficiency.

7. Electrochemistry Galvanic Cells
What are galvanic cells?
Cells that use chemical reactions to generate electrical energy
Batteries are examples of useful galvanic cells
Example reaction
If reactants are placed in a beaker, only products + heat are produced
When half reactions are isolated on electrodes, electrical work can be produced
GALVANIC CELL
voltmeter
Zn(s)
Ag(s)
Zn(s) + 2Ag+ → Zn2+ + 2Ag(s)
AgNO3(aq)
ZnSO4(aq)
Salt Bridge

8. Electrochemistry Galvanic Cells
Description of how example cell works
Reaction on anode = oxidation
Anode = Zn electrode (as the Eº for Zn2+ is less than for that for Ag+)
So, reaction on cathode must be reduction and involve Ag
Oxidation produces e-, so anode has (–) charge (galvanic cells only); current runs from cathode to anode
Salt bridge allows replenishment of ions as cations migrate to cathode and anions toward anodes
GALVANIC CELL
voltmeter
Ag+ + e- → Ag(s)
Zn(s)
Ag(s) + –
AgNO3(aq)
ZnSO4(aq)
Zn(s) → Zn2+ + 2e-
Salt Bridge

9. Electrochemistry Galvanic Cells
Cell notation
Example Cell:
Zn(s)|ZnSO4(aq)||AgNO3(aq)|Ag(s)
GALVANIC CELL
voltmeter
Zn(s)
Ag(s)
“|” means phase boundary
left side for anode (right side for cathode)
“||” means salt bridge
AgNO3(aq)
ZnSO4(aq)
Salt Bridge

10. Electrochemistry Galvanic Cells
Example Questions
Given the following cell, answer the following question:
MnO2(s)|Mn2+(aq)||Cr3+(aq)|Cr(s)
What compound is used for the anode?
What compound is used for the cathode?
Write out both half-cell reactions and a net reaction

11. Electrochemistry Standard Reduction Potential
A half cell or electrode, is half of a galvanic cell
A standard electrode is one under standard conditions (e.g. 1 M AgNO3(aq))
Standard reduction potential (Eº) is cell potential when reducing electrode is coupled to standard hydrogen electrode (oxidation electrode)
Large + Eº means easily reduced compounds on electrode
Large – Eº means easily oxidized compounds on anode
Pt(s)
Ag(s)
H2(g)
AgNO3(aq)
H+(aq) 11

12. Electrochemistry Electrolytic Cells
Used in more advanced electrochemical analysis (not covered in detail)
Uses voltage to drive (unfavorable) chemical reactions
Example: use of voltage to oxidize phenol in an HPLC electrochemical detector (E° of 0 to 0.5 V needed)
anode (note: oxidation driven by voltage, but now + charge)
cathode (reduction, - charge)

13. Electrochemistry The Nernst Equation
The Nernst Equation relates thermodynamic quantities to electrical quantities for a cell reaction
Thermodynamics:
ΔG = ΔGº + RTlnQ ΔG = free energy, Q = reaction quotient
so, -nFE = -nFEº + RTlnQ, or E = Eº – (RT/nF)lnQ
more often seen as: E = Eº – ( /n)logQ (although only valid at T = 298K)
Note: in calculations, E is for reductions (even if oxidation actually occurs at that electrode)
Equation for electrodes or full cells, although text uses Ecell = E+ – E- where + and – refer to voltmeter leads
Best to use activities in Q (even though we will just use concentrations)

14. Electrochemistry The Nernst Equation
Example: Determine the voltage for a Ag/AgCl electrode when [Cl-] = M if Eº = V (at T = 25°C)? 14

15. Electrochemistry Applications of The Nernst Equation
Examples:
The following electrode, Cd(s)|CdC2O4(s)|C2O42- is used to determine [C2O42-]. It is paired with a reference electrode that has an E value of V (vs. the S.H.E.) with the reference electrode connected to the + end of the voltmeter. If Eº for the above reduction reaction is V, and the measured voltage is V, what is [C2O42-]?