18 Introduction to Electrochemistry
18A Characterizing Oxidation/reduction reactions
redox reactions: Ce4+ + Fe2+ → Ce3+ + Fe3+
oxidizing agent, oxidant (electron acceptor): Ce4+ + e- → Ce3+ (reduction of Ce4+ )
reducing agent, reductant (electron donor): Fe2+ → Fe3+ + e- (oxidation of Fe2+ )
MnO4- + 5e- + 8H+ → Mn2+ + 4H2O
5Fe2+ → 5Fe3+ + 5eMnO4- + 5Fe2+ + 8H+ → Mn2+ + 5Fe3+ + 4H2O
18A-1 Comparing Redox Reaction to Acid/Base Reactions
acid1 + base2 → base1 + acid2 (acid/base reaction)
Ared + Box → Aox + Bred
(oxidation/reduction reaction)
Ex. 18-1 2H+
+ Cd(s) → H2(g) + Cd2+
2Ag+ + H2(g) → 2Ag(s) + 2H+
Cd2+ + Zn(s) → Cd(s) + Zn2+
What can we deduce regarding the strengths of H+, Ag+, Cd2+ and Zn2+ as
electron acceptors (or oxidizing agents)?
∵ H+ > Cd2+, Ag+ > H+ , Cd2+ > Zn2+ ∴ Ag+ > H+ > Cd2+ > Zn2+
18A-2 Oxidation/Reduction Reaction in Electrochemical Cells
Ag+
Ag+
+ e- → Ag(s)
Cu(s) → Cu2+ + 2e+ Cu(s) → Ag(s) + Cu2+
(a)
Fig. 18-2 (a) a galvanic cell at open
circuit; (b) a galvanic cell doing work;
(c)an electrolytic cell
(c)
(b)
120
18B ELECTROCHEMICAL CELLS
Consists of two conductors (electrodes), each of which is immersed in an electrolyte
solution.
18B-1 Cathodes and Anodes
cathode : reduction occurs:
NO3- +
Anode : oxidation occurs:
Ag+ + e- → Ag(s)
Fe3+ + e- → Fe2+
10H+ + 8e- → NH4+ +
Cu(s) → Cu2+ + 2e2Cl- → Cl2(g) + 2eFe2+ → Fe3+ + e-
3H2O
18B-2 Types of Electrochemical Cells
galvanic (voltaic) cells : store and supply electrical energy
proceed spontaneous reaction and produce a flow of electron from the anode to the
cathod via an external conductor.
electrolytic cells : required an external source of electrical energy
2Ag(s) + Cu2+ → 2Ag+ + Cu(s)
18B-3 Representing Cells Schematically
Cu|Cu2+(0.0200 M)∥Ag+(0.0200 M)|Ag
anode
cathode
Cu|CuSO4(0.0200 M)∥AgNO3(0.0200 M)|Ag
liquid-junction potential: develop across the interface between two solutions that
differ in their electrolyte composition.
18B-4 Currents in Electrochemical Cells
Electricity is carried by
movement of anions toward
the anode and cations toward
the cathode
Fig. 18-3 Movement of
charge in a galvanic cell
121
18C ELECTRODE POTENTIALS
2Ag+
+
Cu(s) → 2Ag(s) + Cu2+ (Fig. 18-4a)
ΔG = –nFEcell
0
ΔG 0 = −nFEcell
= −RT lnK eq
R : the gas constant
T : the absolute temperature
(a)
(a)
the high-resistance
voltmeter prevents any
significant electron
flow, and the full open
circuit cell potential is
measured. For the
concentrations shown,
this value is + 0.0412V
(b)
(b)
the voltmeter is
replaced with a lowresistance current meter
and the cell discharges
with time until
eventually equilibrium
is reached.
(c)
(c) after equilibruium is
reached, the cell
potential is again
measured with a
voltmeter and found to
be 0.000V. The
concentrations in the
cell are now those at
equilibrium as shown.
Fig. 18-4 Change in cell potential after passage of current until equilibrium is reached.
122
18C-1 Sign Convention for Cell Potential
An electrode potential is, by definition, a reduction potential.
An oxidation potential is the potential for the half-reaction written in the opposite way.
The sign of an oxidation potential is, therefore, opposite that for a reduction potential,
but the magnitude is the same.
The IUPAC sign convention is based on the actual sign of the half-cell of interest when
it is connected to the standard hydrogen electrode.
Cell potential Ecell
Ecell = Eright – Eleft
Fig 18-5
Cell potential in the galvanic cell of Fig. 18-4b
as a function of time. The current, which is
directly related to the cell potential, also
decreases with the same time behavior.
18C-2 The Standard Hydrogen Reference Electrode
Standard Hydrogen Electrode (SHE), Normal
Hydrogen Electrode
2H+ (aq) + 2e- → H2(g)
Pt, H2(p = 1.00 atm)|([H+] = x M)∥
The potential of the standard hydrogen electrode:
0.000 V at all temp. at pH2 = 1.00, aH+ = 1.00,
Fig. 18-6 The hydrogen gas electrode.
18C-3 Defining Electrode Potential and Standard Electrode Potential
Electrode potential: the potential of a cell consisting of the electrode of interest acting as
a cathode and the standard hydrogen electrode acting as an anode.
Ecell = Eright – Eleft = EAg – ESHE = EAg – 0.000 = EAg
Standard electrode potential of a half-reaction Eº:
the activities of the reactants and products are all unity.
half-reaction Ag+ + e- → Ag(s)
Pt, H2(p = 1.00 atm)|H+ ( aH+ = 1.00)∥Ag+(aAg+ = 1.00)|Ag
SHE∥Ag+(aAg+ = 1.00)|Ag
Standard electrode potential (EºAg+/Ag) = + 0.799 V
123
Fig. 18-7 Measurement of the electrode
potential for a Ag electrode. If the silver
ion activity in the right-hand
compartment is 1.00, the cell potential is
the standard electrode potential of the
Ag+/Ag half-reaction.
Fig. 18-8 Measurement of the standard
electrode potential for Cd2+ + 2e- → Cd(s)
Cd2+ + 2e- → Cd(s)
EºCd2+/Cd+ = -0.403 V
Cd(s) + 2H+ → Cd2+ + H2(g)
Half-Reaction
Standard Electrode Potential, V
+
The relative strength of the
+ 0.799
Ag + e → Ag(s)
four ionic species as electron
+ 0.000
2H+ + 2e- → H2(g)
acceptors (oxidizing agents):
Cd2+ + 2e- → Cd(s)
− 0.413
Ag+ > H+ > Cd2+ > Zn2+
2+
Zn + 2e → Zn(s)
− 0.763
18C-5 Effect of Concentration on Electrode Potentials: The Nernst
Equation
aA
+
bB +
....
RT [C]c [ D]d ⋅ ⋅ ⋅
E=E −
ln
nF [A]a [B]b ⋅ ⋅ ⋅
o
ne- →
+
cC +
dD +
...
Eº = standard electrode potential
R = gas constant 8.314 J K-1 mol-1
T = temperature in kelvins
n = no. of moles of e- that appear in the half-reaction
F = faraday = 96485 C (coulombs)
ln = natural logarithm = 2.303 log
[C] c [ D] d ⋅ ⋅ ⋅
0.0592
E=E −
log
n
[A] a [B] b ⋅ ⋅ ⋅
o
Ex. 18-2 (1) Zn2+
(2) Fe3+
+ 2e+
e-
→
Zn(s)
→ Fe2+
124
E = Eo −
0.0592
1
log
2
[Zn 2 + ]
0.0592
[Fe 2 + ]
E=E −
log
1
[Fe3+ ]
o
(3) 2H+ + 2e-
PH 2
0.0592
E=E −
log
2
[H + ]2
o
→ H2(g)
(4) MnO4- + 5e-
+ 8H+
→
Mn2+
+ H2O
0.0592
[Mn 2 + ]
E=E −
log
5
[MnO -4 ][H + ]8
(5) AgCl(s) + e- → Ag(s) + ClE = Eº – (0.0592) log [Cl-]
o
18C-6 The Standard Electrode Potential, Eº (Table 18-1)
1.
2.
3.
4.
Eº → the electron potential when all reactants and products of a halfreaction have unity activity.
Eº relative potential (reference electrode: SHE = 0 V)
a relative reduction potential
the reactants and products are at unit activities
Eº independent of the number of moles reactant and product
Fe3+ +
e- → Fe2+
Eº = + 0.771 V
5Fe3+ + 5e- → 5Fe2+
0.0592
[Fe 2 + ]
E = 0.771 −
log
1
[Fe3+ ]
Eº = + 0.771 V
[Fe 2 + ] 5
0.0592
[Fe 2 + ]5
0.0592
E = 0.771 −
log
log(
)
= 0.771 −
5
5
[Fe3+ ]5
[Fe3+ ]
5 × 0.0592
[Fe 2 + ]
log
= 0.771 −
5
[Fe3+ ]
5. E: + → stronger oxidant than H+, – → weaker oxidant than H+
6. Eº: temperature dependent
Systems Involving Precipitates or Complex Ions
Ag+ + e- → Ag(s)
EºAg+/Ag = + 0.799 V
EºAgCl/Ag = + 0.222 V
AgCl(s) + e- → Ag(s) + Cl32Ag(S2O3)2 + e → Ag(s) + 2S2O3
EºAg(S2O3)23- = + 0.017 V
0.0592
1
[Cl - ]
0.0592
0
0
=
−
E=E +
log
log
E +
−
Ag /Ag
1
Ag /Ag
K sp
1
[Ag+ ]
= E0 +
+ 0.0592 log K sp − 0.0592 log[Cl- ]
Ag /Ag
[Cl ] = 1.00, E= EºAgCl/Ag
E 0AgCl/Ag = E 0
+
Ag /Ag
+ 0.0592 log 1.82 × 10 −10 − 0.0592 log 1.00
= 0.799 + (-0.577) - 0.000 = 0.222 V
125
Fig. 18-9
Measurement of the standard
electrode potential for an Ag/AgCl
electrode.
Ex. 18-3 Calculate the electrode potential of a silver electrode immersed in a 0.0500 M
solution of NaCl using (a) EºAg+ = 0.799 V, and (b) EºAgCl = 0.222 V.
+ 0.799
(a) Ag+ + e- → Ag(s)
1.82 × 10 −10
[Ag ] =
=
= 3.64 × 10 −9 M
0.0500
[Cl- ]
+
Ksp
E = 0.799 − 0.0592 log
1
3.64 × 10
−9
= 0.299 V
-
(b) E = 0.222 – 0.0592 log[Cl ] = 0.222– 0.0592 log 0.0500 = 0.299 V
18C-7 Limitations to the Use of Standard Electrode Potentials
Use of Concentrations Instead of Activities
Effect of Other Equilibria
Formal Potentials
The electrode potential when the ratio of
analytical concentrations of reactants and
products of a half-reaction is exactly 1.00
and the molar concentrations of any other
solutes are specified.
Fig.18-10 Measurement of the formal
potential of the Ag/Ag+ couple in 1M
HClO4.
126