1. Chem. 1B – 11/3 Lecture

2. Announcements I Exam 2 - Results Average about 65%
Broader Distribution than Exam 1
Time was more of an issue
Several questions with low % correct were easy conceptual questions
Questions with lowest % correct:
Score Range # 7
80s 22
70s 20
60s 38
50s 28
<50 21
Version
24%
27%
31%
37%
40%
41% A 23 21 12 9 19 17 B 10 11 18 22

3. Announcements II Lab Mastering – Some problems recently?
Starting Experiment 9 (Wed./Thurs.)
Next Week: Quiz 8 (Experiment Questions + Electrochem Questions)
Mastering – Some problems recently?
Today’s Lecture
Electrochemistry (Ch. 18)
Some review + basic concepts today
Definitions
Standard Half-Cells and Cells
Standard Reduction Potential

4. Chapter 18 Electrochemistry
Electrochemical Reactions
Balancing Redox Reactions:
6 step method:
Assign oxidation states
Separate overall reaction into oxidation and reduction reactions
Balance each half reaction with respect to mass in order a) mass all elements other than H, O, b) O by adding H2O, c) by adding H+, d) Add OH- to both side if in alkaline sol’n
Balance each half reaction for charge by adding electrons
Use common multiplier to get equal numbers of electrons for each half-reaction
Add each half reaction together to get net reaction without electrons as reactants or products
Note: steps 5 and 6 are skipped if stopping at half reactions

5. Chapter 18 Electrochemistry
Electrochemical Reactions
Balancing Redox Reactions:
Examples (unbalanced):
AgNO3(aq) + Zn(s) ↔ Ag(s) + Zn(NO3)2(aq)
HClO(aq) + Fe2+(aq) ↔ Cl2(g) + Fe3+(aq)
MnO4- (aq) + C2O42-(aq) ↔ Mn2+(aq) + CO2(g)

6. Chapter 18 Electrochemistry
Electrochemical Reactions – Different Forms
“Beaker” Reactions
Products form along with heat (assuming DH < 0)
Little control of reaction
Products co-mingled (from reduction and oxidation)
Example: nail “rusts” (oxidation of Fe, reduction of O2)
Voltaic (Galvanic) Cells
Oxidation and reduction reactions may be divided into different parts (half-cells sometimes physically separated through two reaction cells)
Two electrodes are also needed
Reaction can be “harnessed” through voltage/power production
Examples: batteries, pH measuring electrodes

7. Chapter 18 Electrochemistry
Electrochemical Reactions – Different Forms
Electrolytic Cell
In this type of cell, external electrical energy is used to force unfavorable reactions (e.g. 2H2O(l) ↔ 2H2(g) + O2(g)) to occur
Also requires two electrodes – but some differences from electrodes of voltaic cells
Examples: Production of Cl2 gas from NaCl(aq), production of H2 gas from water (above), instruments that measure degree of oxidation/reduction at specific voltages (analogous to spectrometers)

8. Chapter 18 Electrochemistry
Voltaic 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. Chapter 18 Electrochemistry
Basic Electrical Quantities
Current: the flow of electrons (although defined where a positive current has electrons moving backwards)
Current units: Amperes (A) with 1 A = 1 C/s and 1 C = 1 Coulomb where 1 electron (elementary charge) has a value of 1.60 x C
Potential or Voltage: The potential energy associated with the movement of charge (e.g. to electrode of opposite sign)
Potential units: Volts (V) = 1 J/C

10. Chapter 18 Electrochemistry
Basic Electrical Quantities – From Voltaic Cells
Current: related to the flow of electrons
Potential: related to the reaction occurring (more energetic means higher potential)
The ability of a metal (or other elements) to reduce can be measured under standard conditions
Example: Zn(s) + 2Ag+(aq) ↔ Zn2+(aq) + 2Ag(s)
If [Ag+] and [Zn2+] = 1 M, Ecellº = 1.56 V

11. Chapter 18 Electrochemistry Voltaic Cells
Cell notation
Example Cell:
Zn(s)|Zn2+(aq)||Ag+ (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

12. Chapter 18 Electrochemistry
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

13. Chapter 18 Electrochemistry
Given the following cell, write the cell notation:
GALVANIC CELL
voltmeter – reads V
Pt(s)
Ag(s) – +
Note: In this case the Pt(s) is an “inert” electrode (provides electrons but doesn’t react
AgCl(s)
NaCl(aq)
FeSO4 (aq), Fe2(SO4)3(aq)
Salt Bridge 13

14. Chapter 18 Electrochemistry Standard Reduction Potential
A cell used to determine the standard reduction potential consists of two half cells
One half-cell, the anode, is the standard hydrogen electrode (2H+(aq) + 2e- ↔ H2(g))
Eanodeº = 0 (defined)
Other is the test cell (compound being reduced when half-cell is coupled to standard hydrogen electrode (oxidation electrode)
Both cells under standard conditions (1 M, 1 atm)
Ecellº = Ecathodeº
The SHE is not actually used much any more (just a reference for relative potential)
Pt(s)
Ag(s)
H2(g)
AgNO3(aq)
H+(aq) 14

15. Chapter 18 Electrochemistry Standard Reduction Potential
Meaning of Values
Half-cells that exhibit positive values have electrodes with compounds that easily reduce (e.g. Ag+(aq), MnO4-, PbO2(s))
Half-cells that exhibit negative values have electrodes that easily oxidize (e.g. alkali metals)
What if we have two half-cells (neither SHE), can we find Ecellº?
Example: Zn(s)|Zn2+(aq)||Ag+ (aq)|Ag(s) Ecellº = ?
Ag+ reduction
Eº = V
Eº = 0
Zn2+ reduction
Eº = V

16. Chapter 18 Electrochemistry
Example Question
An Ag/AgCl electrode is a common reference electrode. What is the standard potential of a cell made up of a Cu2+ solution being reduced to Cu(s) and AgCl(s) being reduced to Ag(s)?
E°(Cu2+ + 2e- ↔ Cu(s)) = 0.34 V
E°(AgCl(s) + e- ↔ Ag(s) + Cl- (aq)) = 0.22 V
What is the balanced reaction and what species must be present at 1 M?