1. Electrochemistry Generating Voltage (Potential)

2. Historically Historically oxidation involved reaction with O2.
i.e., Rusting
4 Fe(s) + 3O2 (g) Fe2O3 (s)
Other example
Zn(s) + Cu2+(aq) g Zn2+(aq) + Cu(s)
In this reaction:
Zn(s) g Zn2+(aq) Oxidation
Cu2+(aq) g Cu(s) Reduction
In a redox reaction, one process can’t occur without the other. Oxidation-Reduction reaction must simultaneously occurs.

3. Redox Between
If Zn(s) and Cu2+(aq) is in the same solution, then the electron is a transferred directly between the Zn and Cu.
No useful work is obtained. However if the reactants are separated and the electrons shuttle through an external path...

4. Electrochemical Cells
Voltaic / Galvanic Cell Apparatus which produce electricity
Electrolytic Cell Apparatus which consumes electricity
Consider: Zn Cu
Initially there is a flow of e-
After some time the process stops
Electron transport stops because of charge build up
Build up of positive charge
Build up of negative charge
The charge separation will lead to process where it cost too much energy to transfer electron.

5. Completing the Circuit
Electron transfer can occur if the circuit is closed
Parts:
Two conductors
Electrolyte solution
Salt Bridge / Porous membrane
3 process must happen if e- is to flow.
A. e- transport through external circuit
B. In the cell, ions a must migrate
C. Circuit must be closed (no charge build up)
Anode (-)
Black
Negative electrode generates electron
Oxidation Occur
Cathode (+)
Red
Positive electrode accepts electron
Reduction Occur A C B
Anode/Anion (-)
Cathode/Cation(+)

6. Voltaic Cell Electron transfer can occur if the circuit is closed
Parts:
Two conductors
Electrolyte solution
Salt Bridge / Porous membrane
3 process must happen if e- is to flow.
A. e- transport through external circuit
B. In the cell, ions a must migrate
C. Circuit must be closed (no charge build up)
Anode (-)
Black
Negative electrode generates electron
Oxidation Occur
Cathode (+)
Red
Positive electrode accepts electron
Reduction Occur
Anode/Anion (-)
Cathode/Cation(+)

7. Completing the Circuit: Salt Bridge
In order for electrons to move through an external wire, charge must not build up at any cell. This is done by the salt bridge in which ions migrate to different compartments neutralize any charge build up.

8. Sign Convention of Voltaic Cell
@ Anode: Negative Terminal (anions).
Source of electron then repels electrons. Oxidation occurs.
Zn(s) g Zn+2(aq) + 2e- : Electron source
@ Cathode: Positive Terminal (cation)
Attracts electron and then consumes electron. Reduction occurs.
Electron target: 2e- + Cu+2(aq) g Cu(s)
Overall:
Zn(s) + Cu+2(aq) g Zn+2(aq) + Cu(s) E° = 1.10 V
Note when the reaction is reverse: Cu(s) + Zn+2(aq) g Cu+2(aq) + Zn(s)
Sign of E ° is also reversed E° = V
Oxidation: Zn(s) g Zn+2(aq) E° = 0.76 V
Reduction: Cu+2(aq) g Cu(s) E° = 0.34 V
1.10 V = E°CELL
or E°CELL = E°red (Red-cathode) - E°red (Oxid-anode)

9. Other Voltaic Cell Zn(s) + 2H+ (aq) g Zn+2(aq) + H2 (g) E° = 0.76 V
@ Anode: Negative Terminal (anions): Zn(s) g Zn+2(aq) + 2e- :
Source of electron then repels electrons. Oxidation occurs.
@ Cathode: Positive Terminal (cation): 2e H+(aq) g H2 (g)
Attracts electron and then consumes electron. Reduction occurs.
Net: Zn(s) + 2H+ (aq) g Zn2+ (aq) + H2 (g)

10. Other Voltaic Cell Zn(s) + 2H+ (aq) g Zn+2(aq) + H2 (g) E° = 0.76 V
@ Anode: Negative Terminal (anions): Zn(s) g Zn+2(aq) + 2e- :
Source of electron then repels electrons. Oxidation occurs.
@ Cathode: Positive Terminal (cation): 2e H+(aq) g H2 (g)
Attracts electron and then consumes electron. Reduction occurs.
Net: Zn(s) + 2H+ (aq) g Zn2+ (aq) + H2 (g)

11. Line Notation Convention
Line notation: Convenient convention for electrochemical cell
Schematic Representation
1. Anode g Cathode
[oxidation (-) ] [reduction (+)]
2. “ “ phase boundary
(where potential may develop)
3. “ “ Liquid junction
4. Concentration of component
Zn(s) ZnSO4 (aq,1.0M) CuSO4 (aq,1.0M) Cu(s) 4 1 3 2

12. Line Notation Examples
Consider : Zn(s) + Cu+2(aq) g Zn+2(aq) + Cu(s
Anode: Zn g Zn e-
Cathode: Cu e- g Cu
Shorthand “Line” notation
Zn (s) Zn+2 (aq)(1.0M) Cu+2(aq) (1.0M) Cu(s)
2nd Example : Zn(s) + 2H+ (aq) g Zn+2(aq) + H2(g)
Cathode: 2H e- g H2 (g)
Zn (s) Zn+2 (aq)(1.0M) H+(aq) (1.0M), H2(g, 1atm) Pt(s)

13. Other Voltaic Cell & Their Line Notation
Oxidation half-reaction
Zn(s) g Zn+2(aq) e-
Oxidation half-reaction
2I- (aq) g I2 (s) + 2e-
Oxidation half-reaction
Cr(s) g Cr+3(aq) e-
Reduction half-reaction
MnO4-(aq) + 8H+(aq) + 5e-
g Mn 2+(aq) +4H2O(l)
Reduction half-reaction
Ag+(aq) + e- g Ag (s)
Zn(s) | Zn+2 (aq)||H+(aq) , H2 (g,1atm)|Pt
Cr(s) | Cr+3 (aq)||Ag+(aq) | Ag(s)
C(s)| I-(aq) , I2 (g,1atm) || MnO4-(aq) , Mn+2 (aq)| C(s)

14. Line Notation Examples
Example 1: B&L 20.13
Zn(s) + Ni2+(aq) g Zn+2(aq) + Ni (aq)
Example 2: B&L 20.19
Tl+3(aq) + 2Cr2+(aq) g Tl+(aq) + Cr+3(aq)

15. Voltage of Galvanic / Voltaic Cell
Transport of any object requires a net force.
Consider water flowing through pipes. This occurs because of pressure gradient.
Flow (Fluid Transport)
Pressure (h)
Pressure (i) Or
Similarly, electron are transported through wires because of the electromotive force EMF or Ecell.
Object falling or transport down due to Dh
(+)
(-) e -

16. EMF - ElectroMotive Force
Potential energy of electron is higher at the anode. This is the driving force for the reaction (e- transfer) e
Anode (-)
D P.E. = V = J
e - C
e- flow toward cathode
(+) Cathode
Larger the gap, the greater the potential (Voltage)

17. ElectroMotive Force EMF - Electro Motive Force
Potential energy difference between the two electrodes
The larger the DP.E. the larger EMF value.
The magnitude of P.E. for the reaction (half reaction) is an intensive property)
i.e., Size independent: r, Tbpt, Cs.
Therefore EMF is also an intensive property.
Analogy:
Size of rock not important, only the height from ground.
(Electron all have the same mass)
Unit: EMF: V - Volts :
1V - 1 Joule / Coulomb
1 Joule of work per coulomb of charge transferred.

18. Stoichiometry Relationship to E°
EMF - Intensive Property
E°cell Standard state conditions 25°C, 1atm, 1.0 M
E°cell Intensive property, Size Independent
Consider:
Li+ + e- g Li (s) E°Cell = V
x Li e- g 2 Li (s) E°Cell = ( V) x 2 = ??
But E° = Voltage per electron
E° ‘ = E° x 2 = ? g V • 2 = V
2 e-
\ Stoichiometry does not change E°, but reversing the reaction does change the sign of E°.

19. Standard Reduction Potential
Cell Potential is written as a reduction equation.
M+ + e- g M E° = std red. potential
Most spontaneous <Reduction occurs> Oxidizing Agent
Written as reduction
Most non-spontaneous Spontaneous in the reverse direction. <Oxidation occurs> Reducing Agent

20. Zoom View of Std. Reduction Potential
Cell Potential is written as a reduction equation.
M+ + e- g M E° =
F2 (g) e- g 2 F - (aq) V
Ce e- g Ce3+ (aq) V
2H e- g H2 (g) V
Li+(aq) + e- g Li(s) V
Most spontaneous Reduction Oxidizing Agent
Written as reduction
Most non-spontaneous Spontaneous in the reverse direction. Oxidation Reducing Agent
All reaction written as reduction reaction. But in electrochemistry, there can’t be just a reduction reaction. It must be coupled with an oxidation reaction.

21. E°Cell Evaluation E°Cell Function of the reaction
g Oxidation Process (Anode reaction)
g Reduction Process (Cathode reaction) or
E°Cell = E°Cathode & E°Anode
Cathode (+)
Anode (-) Most Negative Reduction reaction
Therefore,
E°Cell = E°red (Cathode) - E°red (anode)
Neg Minus (Large negative)
(Very Positive Value)
Very Positive
\ Very Spontaneous

22. Standard Reduction Potential
How is E°red (Cathode) and E°red (Anode) determine.
E° (EMF) - State Function; there is no absolute scale
Absolute E° value can’t be measured experimentally
The method of establishing a scale is to measure the difference in potential between two half-cells.
Consider:
Zn g Zn e E°=?
Can’t determine because the reaction must be coupled
How can a scale of reduction potential be determine ?
Use a half reaction as reference and assign it a potential of zero.
Electrochemical reaction more spontaneous than this reference will have positive E°, and those less spontaneous will have negative E°.

23. Side-Bar: Relative Scale
Consider a baby whose weight is to be determine but will not remain still on top of a scale. How can the parents determine the babies weight?
Carry the child in arms and weight both child and parent then subtract the weight of the parent from the total to yield the baby weight.

24. Reference Potential Selected half reaction is:
H+ / H2 (g) couple half reaction: 2H+ (aq, 1.0M) + 2e- g H2 (g,1atm)
by definition c E° = 0.0 V, the reverse is also 0.0 V
H+/H2 couple - Standard Hydrogen Electrode (SHE)
To determine E° for a another half reaction, the reaction of interest needs to be coupled to this SHE. The potential measured is then assigned to the half-reaction under investigation.
E°Cell = V = E°red (Cat) - E°red (Anode)
0.0 V (?)
E°red (Anode) = V
\Zn+2/Zn E° = V Reduction rxn

25. Determining Other Half-Cell Potential
Now consider the reaction:
Zn(s)|Zn+2 (1.0 M)||Cu+2(1.0 M)|Cu(s)
E°Cell = V
E°Cell = E°red (Cat) - E°red (Anode)
recall, E° Zn+2/Zn = V
Therefore,
E°Cell = E°Cu+2/Cu - E° Zn+2/Zn
1.10 V = (?) ( V)
E°Cu+2/Cu = V

26. Example: Half-Cell Potential
Example BBL20.19:
For the reaction: Tl Cr2+  Tl Cr3+ E°Cell = 1.19 V
i) Write both half reaction and balance
ii) Calculate the E°Cell Tl+3  Tl+
iii) Sketch the voltaic cell and line notation
i) Tl e-  Tl+
(Cr2+  2Cr3+ + 2e- ) x E° = 0.41 V
ii) E°Cell = 1.19 V = E°red (Cat) - E°red (Anode)
1.19 V= E°red (Cat) V
for Tl e-  Tl+ :
1.19 V = E°red (Cat) = V Pt
Cr2+ g Cr3+
Tl3+ g Tl+
1.19 V

27. Voltaic Vs. Electrolytic Cells
Voltaic Cell
Energy is released from spontaneous redox reaction
Electrolytic Cell
Energy is absorbed to drive nonspontaneous redox reaction
General characteristics of voltaic and electrolytic cells. A voltaic cell generates energy from a spontaneous reaction (DG<0), whereas an electrolytic cell requires energy to drive a nonspontaneous reaction (DG>0). In both types of cell, two external circuits provides the means or electrons to flow. Oxidation takes place all the anode, and reduction takes place at the cathode, but the relative electrode changes are opposite in the two cells.
System does work on load (surroundings)
Surrounding (power supply) do work on system (cell)
Anode (Oxidation)
Oxidation Reaction
X g X+ + e-
Oxidation Reaction
A- g A + e-
Reduction Reaction
e- + Y+ g Y
Reduction Reaction
e- + B+ g B
Overall (Cell) Reaction
A- + B+ g A + B, DG> 0
Overall (Cell) Reaction
X + Y+ g X+ + Y, DG = 0