This document discusses electrochemical cells and half-reactions. It defines electrochemical cells as consisting of two electrodes separated by an electrolyte. In a galvanic cell, spontaneous reactions occur that produce electricity. In an electrolytic cell, an external current is used to drive nonspontaneous reactions. Redox reactions can be written as the sum of oxidation and reduction half-reactions. The document also describes different types of electrodes, including gas electrodes using hydrogen and redox electrodes like the Daniell cell. It explains how the cell potential is the difference between electrode potentials and relates this to the maximum work done by the cell.

1. Nils Walter: Chem 260
Electrochemical cells
= electronic conductor
+ surrounding
electrolyte
electrode compartment
If two different
electrolytes are used:
Electrolytic cell: electrochemical cell in which a non-spontaneous reaction is
driven by an external source of current
Galvanic cell: electrochemical cell in which electricity is produced as a result
of a spontaneous reaction (e.g., batteries, fuel cells, electric fish!)

2. Nils Walter: Chem 260
Reactions at electrodes: Half-reactions
Redox reactions: Reactions in which electrons are transferred
from one species to another
+II -II 0 -II
0 +IV
reduced oxidized
Any redox reactions can be expressed as the difference
between two reduction half-reactions in which e- are taken up
CuS(s) + O2(g) →
→
→
→ Cu(s) + SO2(g)
E.g.,
Reduction of Cu2+: Cu2+(aq) + 2e- →
→
→
→ Cu(s)
Reduction of Zn2+: Zn2+(aq) + 2e- →
→
→
→ Zn(s)
Cu2+(aq) + Zn(s) →
→
→
→ Cu(s) + Zn2+(aq)
More complex: MnO4
-(aq) + 8H+ + 5e- →
→
→
→ Mn2+(aq) + 4H2O(l)
Half-reactions are only a formal way of writing a redox reaction
Difference:

3. Nils Walter: Chem 260
Carrying the concept further
Reduction of Cu2+: Cu2+(aq) + 2e- →
→
→
→ Cu(s)
In general: redox couple Ox/Red, half-reaction Ox + ν
ν
ν
νe- →
→
→
→ Red
Any reaction can be expressed in redox half-reactions:
Expansion of gas: H2(g, pi) →
→
→
→ H2(g, pf)
2 H+(aq) + 2e- →
→
→
→ H2(g, pi)
2 H+(aq) + 2e- →
→
→
→ H2(g, pf)
Dissolution of a sparingly soluble salt: AgCl(s) →
→
→
→ Ag+(aq) + Cl-(aq)
Ag+(aq) + e- →
→
→
→ Ag(s)
AgCl(s) + e- →
→
→
→ Ag(s) + Cl-(aq)
Reaction quotients: ]
[ −
≈
= − Cl
a
Q Cl
]
[
1
1
+
≈
=
+ Ag
a
Q
Ag

4. Nils Walter: Chem 260
Reactions at electrodes
Galvanic cell: Electrolytic cell:
Red1 →
→
→
→
Ox1 + ν
ν
ν
νe-
Ox2 + ν
ν
ν
νe-
→
→
→
→ Red2
Half-reactions
Red1 →
→
→
→
Ox1 + ν
ν
ν
νe-
Ox2 + ν
ν
ν
νe-
→
→
→
→ Red2

5. Nils Walter: Chem 260
Insoluble-salt electrode:
metal
(e.g., Pt)
solution
(e.g., H+)
2H+ + 2e- H2(g) 2
2 )
(
+
=
H
a
H
p
Q
Types of electrodes I
Pt(s)|
||
| H2(g)|
||
| H+(aq)
Gas electrode:
Gas (e.g., H2)
metal
(e.g., Ag)
solution
(e.g., Cl-)
AgCl(s) + e- Ag(s) + Cl-(aq)
−
= Cl
a
Q
Ag(s)|
||
| AgCl(s)|
||
| Cl-(aq)
Porous, insoluble
salt (e.g., AgCl)

6. Nils Walter: Chem 260
Fe3+ + e- Fe2+
+
+
=
=
3
2
Re
Fe
Fe
Ox
d
a
a
a
a
Q
Types of electrodes and how to put them
together in a galvanic cell
Pt(s)|
||
| Fe2+(aq),Fe3+(aq)
+
+
=
2
2
Cu
Zn
a
a
Q
Zn(s)|
||
| ZnSO4(aq) |
||
| |
||
| CuSO4(aq)|
||
| Cu(s)
Redox electrode: Daniell cell:
Right:
Cu2+(aq) + 2e-
→
→
→
→ Cu(s)
Left:
Zn2+(aq) + 2e-
→
→
→
→ Zn(s)
Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq)

7. Nils Walter: Chem 260
Cell reaction and potential
Cell reaction: Difference of electrode half-reactions
(Reduction at Cathode - Oxidation at Anode)
Cathode (Right): Cu2+(aq) + 2e- →
→
→
→ Cu(s)
Anode (Left): Zn2+(aq) + 2e- →
→
→
→ Zn(s)
Overall (R-L): Cu2+(aq) + Zn(s) Cu(s) + Zn2+(aq) e- disappear
this way the reaction
becomes spontaneous
}
Cell potential E: Potential difference between
the electrodes
Maximum electrical work done in
a galvanic cell: w’ = - ν
ν
ν
νF ×
×
×
× E = ∆
∆
∆
∆rG
Ox + ν
ν
ν
νe- →
→
→
→ Red Þ
Þ
Þ
Þ ν
ν
ν
νNA e- transferred per mole of
reaction
Þ
Þ
Þ
Þ ν
ν
ν
νNA ×
×
×
× (-e) = -ν
ν
ν
νF charge transferred
Faraday constant = eNA
= 96,485 Cmol-1