Fig. 22-1a (p.629) A galvanic electrochemical cell at open circuit
1.1 components
•
Conducting electrodes (metal, carbon)
•
External wires (electrons carry current)
•
Ion electrolyte solution (ions carry current)
•
Interfaces or junctions
•
Complete electrical circuit
conduction:
electrons moving from one electrode to another thru. External wire
within solutions, migration of ions carry current at electrode surface,
oxidation or reduction reactions occur
coupling the electron conduction of electrodes w/ ion conduction of
solutions
1.2 Galvanic Cells – cell develops spontaneous potential difference
Overall:
2Ag+(aq) + Cu(s)  2Ag(s) + Cu2+(aq)
Half reactions: Cu(s)  Cu2+ (aq) + 2eAg+ + e-  Ag (s)
Oxidation
Reduction
Convention:
Cathode where reduction occurs
Anode where oxidation occurs
Galvanic cell:
Cu anode (negative)
Ag cathode (positive)
Cu2+ = 0.02 M, Ag+ = 0.02 M,
E = 0.412 V  0 V
reaction
 equilibrium
Potential difference (voltage) is measure of tendency to move to equilibrium
Fig. 22-1b (p.629) A galvanic cell doing work
1.3 electrolytic cells – require potential difference greater than galvanic
potential difference ( to drive away from equilibrium)
Overall:
2Ag(s) + Cu2+(aq)  2Ag+(aq) + Cu(s)
[chemically reversible cell]
Half reactions: Ag(s)  Ag+ (aq) + 2e- Oxidation
Cu2+(aq) + 2e-  Cu(s) Reduction
electrolytic cell: Ag anode (negative)
Cu cathode (positive)
Fig. 22-1c (p.629) An electrolytic cell
1.4 Cell w/o liquid junctions
AgCl(s)  Ag+ (aq) + Cl- (aq)
H2 (g)  H2 (aq)
Cathode:
AgCl(s)  Ag+ + ClAg+ + e-  Ag (s)-
Anode:
Overall:
H2 (aq)  2H+ (aq) + 2e2AgCl (s) + H2(g)  2Ag(s) + 2H+ + 2Cl-
direct reaction of AgCl + H2 is very slow
Fig. 22-2 (p.631) A galvanic cell without a liquid junction
1.5 Schematic representation of cells
Short-hand cell notation
Cu|CuSO4(Cu2+ = 0.0200) ||AgNO3 (Ag+ = 0.0200)| Ag
| liquid-electrode interface
|| two phase boundaries, one at each end of the salt bridge
convention: anode on left
Galvanic cell as written
Electrolytic cell if reversed
Pt, H2 (p=1atm)|H+ (0.01M), Cl-(0.01M), AgCl (sat’d) | Ag
2.1 Cell potential
(difference between anode and cathode potential)
Ecell = Ecathode – Eanode
when half-reactions written as reduction (electron on left)
Example:
2AgCl(s) + H2(g)  2Ag (s) + 2H+ + 2ClCathode: 2AgCl(s) + 2e- 2Ag (s) + 2ClAnode: 2H+ + 2e-  H2(g)
Galvanic cell Ecell = Ecathode – Eanode = EAgCl/Ag – EH+/H2 = +0.46 V
Can’t measure potential on each electrode independently – only
differences
2.2 Standard reference electrode
Standard hydrogen electrode (SHE)
Pt, H2(p=1.00atm) | H+ (H+ = 1.00M)||…
-
-
SHE assigned 0.000V
can be cathode or anode {depending on the half cell which it couples with}
Pt does not take part in reaction, coated with a finely divided layer of
platinum to provide large surface area
controlled activity of reactants
rarely used for routine measurement
Fig. 22-5 (p.637) measurement of the electrode potential for M electrode
Alternative references electrodes {which are simple to prepare}
a. Calomel electrode
Hg|Hg2Cl2(sat’d), KCl(xM)||
Hg2Cl2 (s) + 2e- 2Cl- + 2Hg (l)
Ereference depends on Cl-1
Ereference = +0.24V vs. SHE for saturated calomel electrode (SCE)
b. Silver-silver chloride electrode
Ag|AgCl(sat’d), KCl(xM)||
AgCl(s) + e-  Cl- + Ag (s)
-
-
-
Similar construction to calomel electrode
Ag wire coated with AgCl
Solution of KCl sat’d with AgCl
Again, Ereference depends on Cl-, and = + 0.22 V vs. SHE
Can be used for uncontrolled temperature (lower temperature coefficient,
see Table 23-1)
Can be used for temp > 60C
But Ag reacts with more ions (e.g, proteins), while Hg reacts with few
sample components,
Fig. 23-2 (p.661) Typical commerical reference electrode a) A saturated
calomel electrode, and b) a silver-silver chloride electrode
2.3 electrode and standard electrode potential (E and E0)
- Definition: potential of electrodes vs. SHE
- Electrode potential varies with activity of ion
Activity x = x[X]
x: activity coefficient, varies with presence of other ions (ionic strength)
(see Appendix 2)
Note: activity of pure liquid or solid in excess = 1.00
Use pressure (atm) for gases
If  = 1.00M, the electrode potential E, becomes standard electrode potential E0
Appendix 3
Cu2+ + 2e-  Cu(s) E0 = +0.337 V
2H2+ + 2e-  H2(g) E0 = +0.000 V
Cd2+ + 2e-  Cd(s) E0 = - 0.403 V
Zn2+ + 2e-  Zn(s) E0 = -0.763 V
2.4 Nernst equation
In general, E and Ecell can be calculated for any activity
using Nernst equation: pP + qQ + ne-  rR + sS
E  E 
0
RT
In
nF
RT
(a R )r (aS )
s
p
q
(a P ) (aQ )
 2 . 568  10
2
F
E  E 
0
2 . 568  10
n
E  E 
0
0 . 0592
n
2
In
(a R )r (aS )
s
p
q
(a P ) (aQ )
log
(a R )r (aS )
s
p
q
(a P ) (aQ )
E = E0 when log quotient is unity
E0 is relative to SHE
E0 is measure of driving force for half-cell reduction
3. Electrical double layer*
electrons transferred at electrode surface by redox reactions, occurring at
solution/electrode interface
Electrical double layer formed
1. tightly bound inner layer
2. loosely bound outer layer
Fig. 22-3 (p.632) Electrical double layer formed at electrode surface as a
result of an applied potentials