© www.chemsheets.co.uk
A2 046
20-Jul-12
Zn2+(aq) + 2 e–  Zn(s)
- electrode
anode
oxidation
At this electrode the
metal loses electrons
and so is oxidised to
metal ions.
+ electrode
cathode
reduction
electron flow
At this electrode the metal
ions gain electrons and so
is reduced to metal atoms.
Zn
These electrons make
the electrode negative.
Zn  Zn2+ + 2 eoxidation
Cu
As electrons are used up,
this makes the electrode
positive.
Cu2+ + 2 e-  Cu
reduction
Standard Conditions
Concentration
1.0 mol dm-3 (ions involved in ½ equation)
Temperature
298 K
Pressure
100 kPa (if gases involved in ½ equation)
Current
Zero (use high resistance voltmeter)
S tandard
H ydrogen
E lectrode
Emf = E = Eright Eleft
V
high resistance
voltmeter
H2 at 100 kPa
salt bridge
o
o
o
o
temperature
= 298 K
Cu
o
o
o
o
o
o
o
Pt
o
1.0 M H+(aq)
1.0 M Cu2+(aq)
E =
Eright
V
high resistance
voltmeter
H2 at 100 kPa
salt bridge
o
o
o
o
temperature
= 298 K
o
o
o
o
o
Cu
o
o
Pt
o
1.0 M H+(aq)
1.0 M Cu2+(aq)
Pt(s) | H2(g) | H+(aq) || Cu2+(aq) | Cu(s)
ROOR
Ni(s) | Ni2+(aq) || Sn4+(aq), Sn2+(aq) | Pt(s)
K(s) | K+(aq) || Mg2+(aq) | Mg(s)
ELECTRODE POTENTIALS – Q1
Emf = Eright - Eleft
- 2.71 = Eright - 0
Eright = - 2.71 V
ELECTRODE POTENTIALS – Q2
Emf = Eright - Eleft
Emf = - 0.44 - 0.22
Emf = - 0.66 V
ELECTRODE POTENTIALS – Q3
Emf = Eright - Eleft
Emf = - 0.13 - (-0.76)
Emf = + 0.63 V
ELECTRODE POTENTIALS – Q4
Emf = Eright - Eleft
+1.02 = +1.36 - Eleft
Eleft = + 1.36 - 1.02 = +0.34
V
ELECTRODE POTENTIALS – Q5
Emf = Eright - Eleft
a) Emf = + 0.15 - (-0.25) = +0.40 V
b) Emf = + 0.80 - 0.54
= +0.26 V
c) Emf = + 1.07 - 1.36
= - 0.29 V
ELECTRODE POTENTIALS – Q6
Emf = Eright - Eleft
a) Eright = +2.00 - 2.38 = - 0.38 V
Ti3+(aq) + e-  Ti2+(aq)
b) Eleft = -2.38 - 0.54 = - 2.92 V
K+(aq) + e-  K(aq)
c) Eright = - 3.19 + 0.27 = - 2.92 V
Ti3+(aq) + e-  Ti2+(aq)
ELECTRODE POTENTIALS – Q7
a) Cr(s) | Cr2+(aq) || Zn2+(aq) | Zn(s)
Emf = -0.76 - (-0.91) = +0.15 V
b) Cu(s) |Cu2+(aq)|| Fe3+(aq),Fe2+(aq)| Pt(s)
Emf = +0.77 - 0.34 = +0.43 V
c)
Pt(s) | Cl-(aq)| Cl2(g) || MnO4-(aq),H+(aq),Mn2+(aq)| Pt(s)
Emf = +1.51 – 1.36 = +0.15 V
Standard electrode potentials
Increasing
oxidising
power
E/V
F2(g) + 2 e-

2 F-(aq)
+ 2.87
MnO42-(aq) + 4 H+(aq) + 2 e-

MnO2(s) + 2 H2O(l)
+ 1.55
MnO4-(aq)
Cr2O72-(aq)
+
-
Cl2(g) + 2 e-
+ 8 H (aq) + 5 e
+

2+
Mn (aq) + 4 H2O(l)
+ 1.51

2 Cl-(aq)
+ 1.36

+ 14 H (aq) + 6 e
-
3+
Br2(g) + 2 e
-
+ 1.33
-

2 Br (aq)
+ 1.09
Ag+(aq) + e-

Ag(s)
+ 0.80
2 Cr (aq) + 7 H2O(l)
-

Fe (aq)
+ 0.77
MnO4-(aq) + e-

MnO42-(aq)
+ 0.56
-
3+
Fe (aq) + e
2+

-
2 I (aq)
+ 0.54

Cu(s)
+ 0.34
Hg2Cl2(aq) + 2 e-

2 Hg(l) + 2 CI-(aq)
-

Ag(s) + Cl (aq)
+
-

H2(g)
0.00
2+
-

Pb(s)
- 0.13
2+
-

Sn(s)
- 0.14
V3+(aq) + e-

V2+(aq)
- 0.26
2+
-

Ni(s)
- 0.25
2+
-

Fe(s)
- 0.44
2+
-

Zn(s)
- 0.76
3+
-

Al(s)
- 1.66
Mg2+(aq) + 2 e-
I2(g) + 2 e
2+
Cu (aq) + 2 e
AgCl(s) + e
2 H (aq) + 2 e
Pb (aq) + 2 e
Sn (aq) + 2 e
Ni (aq) + 2 e
Fe (aq) + 2 e
Zn (aq) + 2 e
Al (aq) + 3 e
-
+ 0.27
+ 0.22

Mg(s)
- 2.36
Na (aq) + e
-

Na(s)
- 2.71
2+
-

Ca(s)
- 2.87
K+(aq) + e-

K(s)
- 2.93
+
Ca (aq) + 2 e
Increasing
reducing
power
The more +ve electrode
gains electrons
(+ charge attracts electrons)
–
0
–ve electrode
+
+ve electrode
e–
+ 1.10 V
+ 0.34 V
Cu2+ + 2 e-  Cu
– 0.76 V
Zn2+ + 2 e-  Zn
Cu2+ + Zn → Cu + Zn2+
PREDICTING REDOX REACTIONS – Q1
–
0
–ve electrode
+ve electrode
e–
+ 0.51 V
– 0.25 V
Ni2+ + 2 e-  Ni
– 0.76 V
Zn2+ + 2 e-  Zn
Ni2+ + Zn → Ni + Zn2+
+
PREDICTING REDOX REACTIONS – Q2
0
+
–ve electrode
+ve electrode
e–
+ 0.46 V
+ 0.80 V
Ag+ + e-  Ag
+ 0.34 V
Cu2+ + 2 e-  Cu
2 Ag+ + Cu → 2 Ag + Cu2+
PREDICTING REDOX REACTIONS – Q3
+
0
+ 1.36 V
Cl2 + 2 e-  2 Cl-
+ 0.77 V
+ 1.51 V
Fe3+ + e-  Fe2+
NO
MnO4- + 8 H+ + 5 e-  Mn2+ + 4 H2O
+ 1.33 V
Cr2O72- + 14 H+ + 6 e-  2 Cr3+ + 7 H2O
NO
YES
PREDICTING REDOX REACTIONS – Q4a
0
+
–ve electrode
2.19 = 0.34 - Eleft
+ve electrode
Eleft = 0.34 – 2.19 = – 1.85 V
e–
+ 2.19 V
+ 0.34 V
Cu2+ + 2 e-  Cu
?V
Be2+ + 2 e-  Be
Be + Cu2+ → Be2+ + Cu
PREDICTING REDOX REACTIONS – Q4b
–
0
–ve electrode
When using SHE
+ve electrode
E = cell emf = – 1.90 V
e–
1.90 V
+ 0.00 V
2 H+ + 2 e-  H2
?V
Th4+ + 4 e-  Th
4 H+ + Th → 2 H2 + Th4+
PREDICTING REDOX REACTIONS – Q5a
Pt(s)|H2(g)|H+(aq)||Br2(aq),Br-(aq)|Pt(s)
+
0
–ve electrode
+ve electrode
e–
+ 1.09 V
+ 1.09 V
Br2 + 2 e-  2 Br-
0.00 V
2 H+ + 2 e-  H2
H2 + Br2 → 2 H+ + 2 Br-
PREDICTING REDOX REACTIONS – Q5b
Cu(s)|Cu2+(aq)||Fe3+(aq),Fe2+(aq)|Pt(s)
0
–ve electrode
+ve electrode
e–
+ 0.43 V
+ 0.77 V
Fe3+ + e-  Fe2+
+ 0.34 V
Cu2+ + 2 e-  Cu
2 Fe3+ + Cu → 2 Fe2+ + Cu2+
+
PREDICTING REDOX REACTIONS – Q6a
Mg(s)|Mg2+(aq)||V3+(aq),V2+(aq)|Pt(s)
–
0
–ve electrode
+ve electrode
e–
+ 2.10 V
– 0.26 V
V3+ + e-  V2+
– 2.36 V
Mg2+ + 2 e-  Mg
YES: Mg reduces V3+ to V2+
PREDICTING REDOX REACTIONS – Q6b
+
0
–ve electrode
+ve electrode
e–
+ 0.59 V
+ 1.36 V
Cl2 + 2 e-  2 Cl-
+ 0.77 V
Fe3+ + e-  Fe2+
NO: Cl- won’t reduce Fe3+ to Fe2+
PREDICTING REDOX REACTIONS – Q6c
0
Pt(s)|Br-(aq),Br2(aq)||Cl2(g)|Cl-(aq)|Pt(s)
–ve electrode
+
+ve electrode
e–
+ 0.27 V
+ 1.36 V
Cl2 + 2 e-  2 Cl-
+ 1.09 V
Br2 + 2 e-  2 Br-
YES: Cl2 oxidises Br- to Br2
PREDICTING REDOX REACTIONS – Q6d
Sn(s)|Sn2+(aq)||Fe3+(aq),Fe2+(aq)|Pt(s)
–
+
0
–ve electrode
+ve electrode
e–
+ 0.91 V
+ 0.77 V
Fe3+ + e-  Fe2+
– 0.14 V
Sn2+ + 2 e-  Sn
YES: Sn reduces Fe3+ to Fe2+
PREDICTING REDOX REACTIONS – Q6e
+
0
–ve electrode
+ve electrode
e–
+ 0.03 V
+ 1.36 V
Cl2 + 2 e-  2 Cl-
+ 1.33 V
Cr2O72- + 14 H+ + 6 e-  2 Cr3+ + 7 H2O
NO: H+/Cr2O72- won’t oxidise Cl- to Cl2
PREDICTING REDOX REACTIONS – Q6f
Pt(s)|Cl-(aq)|Cl2(g)||MnO4- (aq),H+(aq),Mn2+(aq)|Pt(s)
+
0
–ve electrode
+ve electrode
e–
+ 0.03 V
+ 1.51 V
+ 1.36 V
MnO4- + 8 H+ + 5 e-  Mn2+ + 4 H2O
Cl2 + 2 e-  2 Cl-
YES: H+/MnO4- oxidises Cl- to Cl2
PREDICTING REDOX REACTIONS – Q6g
Fe(s)|Fe2+(aq)||H+(aq)|H2(g)|Pt(s)
–
0
–ve electrode
+ve electrode
e–
+ 0.44 V
0.00 V
2 H+ + 2 e-  H2
– 0.44 V
Fe2+ + 2 e-  Fe
YES: H+ oxidises Fe to Fe2+
PREDICTING REDOX REACTIONS – Q6h
0
+
–ve electrode
+ve electrode
e–
+ 0.34 V
+ 0.34 V
Cu2+ + 2 e-  Cu
0.00 V
2 H+ + 2 e-  H2
NO: H+ won’t oxidise Cu to Cu2+
Electrochemical cells
i
appreciate that electrochemical cells can be used as a
commercial source of electrical energy
j
appreciate that cells can be non-rechargeable
(irreversible), rechargeable and fuel cells
k
be able to use given electrode data to deduce the
reactions occurring in non-rechargeable and rechargeable
cells and to deduce the e.m.f. of a cell
l
understand the electrode reactions of a hydrogen-oxygen
fuel cell and appreciate that a fuel cell does not need to
be electrically recharged
m appreciate the benefits and risks to society associated
with the use of these cells
Non-rechargeable (primary) cells – Zinc-carbon
-0.80 V Zn(NH3)22+ + 2 e-  Zn + 2 NH3
+0.70 V 2 MnO2 + 2 H+ + 2 e-  Mn2O3 + H2O
Determine: a) cell emf
b) overall reaction during discharge
• Standard cell
• Short life
Non-rechargeable (primary) cells – alkaline
-0.76 V Zn2+ + 2 e-  Zn
+0.84 V MnO2 + H2O + e-  MnO(OH) + OHDetermine: a) cell emf
b) overall reaction during discharge
• Longer life
Non-rechargeable (primary) cells – lithium
• Very long life
Determine: a) cell emf
b) overall reaction during discharge
• High voltage
Rechargeable (secondary) cells
• In non-rechargeable (primary) cells, the
chemicals are used up so the voltage drops
• In rechargeable (secondary) cells the reactions
are reversible – they are reversed by applying
an external current.
• It is important that the products from the
forward reaction stick to the electrodes and are
not dispersed into the electrolyte.
Rechargeable (secondary) cells – Li ion
+0.60 V Li+ + CoO2 + e-  LiCoO2
-3.00 V Li+ + e-  Li
Determine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge
• Rechargeable
• Most common
rechargeable cell
Rechargeable (secondary) cells – lead-acid
+1.68 V PbO2 + 3 H+ + HSO4- + 2 e-  PbSO4 + 2 H2O
-0.36 V PbSO4 + H+ + 2 e-  Pb + HSO4• Used in sealed car
Determine: a) cell emf
batteries (6 cells
b) overall reaction during discharge
giving about 12 V
c) overall reaction during re-charge
overall)
Rechargeable (secondary) cells – nickel-cadmium
+0.52 V NiO(OH) + 2 H2O + 2 e-  Ni(OH)2 + 2 OH-0.88 V Cd(OH)2 + 2 e-  Cd + 2 OHDetermine: a) cell emf
b) overall reaction during discharge
c) overall reaction during re-charge
FUEL CELLS
+0.40 V O2 + 2 H2O + 4 e-  4 OH-0.83 V 2 H2O + 2 e-  H2 + 2 OHDetermine: a) cell emf
b) overall reaction
• High efficiency (more
efficient than burning
hydrogen)
• How is H2 made?
• Input of H2/O2 to
replenish so no need
to recharge
From wikipedia (public domain)
Pros & cons of cells
+
portable source of electricity
Pros & cons of non-rechargeable cells
+
cheap, small
–
waste issues
Pros & cons of rechargeable cells
+
less waste, cheaper in long run
–
still some waste issues
Pros & cons of fuel cells
+
water is only product
–
most H2 is made using fossil fuels, fuels cells expensive