1. ELECTROCHEMISTRY
Chap 20

2. Types of Electrochemical Cells
Voltaic (galvanic) ⇒ spontaneously produces electricity
Electrolytic ⇒ consumes electricity
Reversible ⇒ reversing current reverses reaction
e.g., auto lead storage battery, hydrogen fuel cell
Irreversible ⇒ cannot be reversed due to different
½-rxn occurring at one or both electrodes
e.g., dry cell, alkaline cell, Hg cell

3. Fig 20.3 A spontaneous oxidation-reduction reaction

4. Fig 20.4 a Voltaic Cell based on Equation 20.7
Zn (s) + Cu2+ (aq, 1 M) → Zn2+ (aq, 1 M) + Cu (s)
1.10

5. Components of a Cell All cells contain 2 electrodes and an electrolyte
Anode – electrode where oxidation occurs
e.g.: Zn (s) ⇌ Zn2+ (aq) + 2e−
Cathode – electrode where reduction occurs
e.g.: Cu2+ (aq) + 2e− ⇌ Cu (s)
Liquid junction - maintains balance of ion charges
e.g.: salt bridge or porous disk

6. Fig 20. 5 A voltaic cell that uses a salt bridge to complete the
Fig A voltaic cell that uses a salt bridge to complete the electrical circuit

7. Fig 20.6 A summary of the terminology used to describe voltaic cells

8. Fig 20.7 Atomic-level depiction of the Zn (s) / Cu2+ (aq) rxn

9. Fig 20.8 Atomic-level depiction of the voltaic cell in Fig 20.52+ 2+

10. Cell EMF under standard conditions
Fig Water analogy for electron flow
Volt (V) - potential difference (E) between two points
1 V = 1 J/C

11. Also called the cell potential, Ecell
Electromotive force (emf) ≡ potential difference between the anode and cathode in a cell
Also called the cell potential, Ecell
Fig 20.5

12. Standard Reduction (Half-Cell) Potentials, E°
Only differences in potential measurable:
∴ measure E° in combination with a standard
half-reaction in a cell
Standard Hydrogen Electrode (SHE)

13. Standard Hydrogen Electrode (SHE)
By convention, reaction written as a reduction:
2H+(aq; 1.00 M) + 2e− ⇌ H2 (g, 1.00 atm)
By definition: E°SHE ≡ 0.00V as a cathode
When used as an anode:
H2 (g, 1.00 atm) ⇌ 2H+(aq; 1.00 M) + 2e−
By definition: E°SHE ≡ 0.00V as a anode

14. Standard Hydrogen Electrode (SHE)

15. Standard Electrode Potential
E° ≡ potential of a cell with electrode of interest acting as a cathode and SHE as the anode
e.g., Cu/Cu2+ couple E° = 0.34 V
e.g., Zn/Zn2+ couple E° = − 0.76 V
E° describes ½-rxn written as a reduction
with respect to SHE
E° > 0: spontaneous w.r.t. SHE
E° < 0: nonspontaneous w.r.t. SHE

16. Fig 20.11 A voltaic cell using a SHE to measure the Eo of
a Zn/Zn2+ electrode

17. Standard Cell Potentials
Ecell 
= Ered (cathode) − Ered (anode)
In the previous example:
Ecell 
= Ered (SHE) − Ered (Zn2+/Zn)
Ecell 
= 0.00 V − 0.76 V = − 0.76 V
Ered (Zn2+/Zn) = − 0.76 V 

18. Table 20.1 Standard Reduction Potentials in Water at 25 oC
Very good
oxidizing agent
(poor reducing
agent)
Very poor
oxidizing agent
(very good
reducing agent)

19. Calculating Standard Cell Potentials
Ecell 
= Ered (cathode) − Ered (anode)
Fig 20.5
For the oxidation:
For the reduction:
Ered = −0.76 V 
Ered = V 
Ecell  =
Ered
(cathode) −
(anode)
= V − (−0.76 V)
= V
© 2009, Prentice-Hall, Inc.

20.   Ecell = Ered (cathode) − Ered (anode)
Sample Exercise 20.6 Calculating E°cell from E°red
Using the standard reduction potentials listed in Table 20.1 (p 857), calculate the standard emf for the voltaic cell, which is based on the reaction:
Ecell 
= Ered (cathode) − Ered (anode)
Cathode: Cr2O72−(aq) H+(aq) + 6 e− → 2 Cr3+(aq) + 7 H2O(l)
Eo = V
Anode: 3 I2(s) e− → 6 I− (aq)
Eo = V
Ecell 
= Ered (cathode) − Ered (anode)
= 1.33 V − V = V