Electrochemical cells
= electronic conductor
+ surrounding
electrolyte
If two different
electrolytes are used:
electrode compartment
Galvanic cell: electrochemical cell in which electricity is produced as a result
of a spontaneous reaction (e.g., batteries, fuel cells, electric fish!)
Electrolytic cell: electrochemical cell in which a non-spontaneous reaction is
driven by an external source of current
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
0 +IV -II
E.g.,
CuS(s) + O2(g) → Cu(s) + SO2(g)
oxidized
reduced
Any redox reactions can be expressed as the difference
between two reduction half-reactions in which e- are taken up
Reduction of Cu2+: Cu2+(aq) + 2e- → Cu(s)
Reduction of Zn2+: Zn2+(aq) + 2e- → Zn(s)
Difference: 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
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 + νe- → Red
Any reaction can be expressed in redox half-reactions:
2 H+(aq) + 2e- → H2(g, pf)
2 H+(aq) + 2e- → H2(g, pi)
Expansion of gas: H2(g, pi) → H2(g, pf)
AgCl(s) + e- → Ag(s) + Cl-(aq)
Ag+(aq) + e- → Ag(s)
Dissolution of a sparingly soluble salt: AgCl(s) → Ag+(aq) + Cl-(aq)
−
Reaction quotients: Q = aCl − ≈ [Cl ]
Q=
1
a Ag +
1
≈
[ Ag + ]
Nils Walter: Chem 260
Reactions at electrodes
Galvanic cell:
Red1 →
Ox1 + νe-
Ox2 + νe→ Red2
Half-reactions
Electrolytic cell:
Red1 →
Ox1 + νe-
Ox2 + νe→ Red2
Nils Walter: Chem 260
Types of electrodes I
Gas electrode:
Insoluble-salt electrode:
Gas (e.g., H2)
Porous, insoluble
salt (e.g., AgCl)
solution metal
(e.g., H+) (e.g., Pt)
solution metal
(e.g., Cl-) (e.g., Ag)
Pt(s)|| H2(g)||
Ag(s)|| AgCl(s)|| Cl-(aq) Q = aCl −
2H+ +
2e-
H+(aq)
p ( H 2 ) AgCl(s) + eAg(s) + Cl-(aq)
H2(g) Q =
2
Nils Walter: Chem 260
aH +
Types of electrodes and how to put them
together in a galvanic cell
Redox electrode:
Daniell cell:
Left:
Zn2+(aq) + 2e→ Zn(s)
Right:
Cu2+(aq) + 2e→ Cu(s)
Pt(s)|| Fe2+(aq),Fe3+(aq)
Fe3+ + e-
Fe2+
aRe d a Fe 2+
=
Q=
aOx
a Fe3+
Zn(s)|| ZnSO4(aq) | | CuSO4(aq)|| Cu(s)
Cu2+(aq) + Zn(s)
Cu(s) + Zn2+(aq)
Q=
aZn 2+
Walter: Chem 260
aCuNils
2+
Cell reaction and potential
Cathode (Right): Cu2+(aq) + 2e- → Cu(s)
Anode (Left): Zn2+(aq) + 2e- → Zn(s)
Overall (R-L): Cu2+(aq) + Zn(s)
}
this way the reaction
becomes spontaneous
Cu(s) + Zn2+(aq)
e- disappear
Cell reaction: Difference of electrode half-reactions
(Reduction at Cathode - Oxidation at Anode)
Cell potential E: Potential difference between
the electrodes
Ox + νe- → Red Þ νNA e- transferred per mole of
reaction Faraday constant = eNA
= 96,485 Cmol-1
Þ νNA × (-e) = -ν
νF charge transferred
Maximum electrical work done in
260
a galvanic cell: w’ = - νNils
F ×Walter:
E = ∆Chem
G
r