Unless otherwise stated, all images in this file have been reproduced from:
Blackman, Bottle, Schmid, Mocerino and Wille,
Chemistry, 2007 (John Wiley)
ISBN: 9 78047081 0866
Slide 35-1
Chem 1101
A/Prof Sébastien Perrier
Room: 351
Phone: 9351-3366
Email: [email protected]
Prof Scott Kable
Room: 311
Phone: 9351-2756
Email: [email protected]
A/Prof Adam Bridgeman
Room: 222
Phone: 9351-2731
Email: [email protected] 35-2
Highlights of last lecture
Electrolysis…
CONCEPTS
How does electrolysis work?
Concept of over-potential,
Chlor-alkali process (Cl2, NaOH)
Hall-Herault process (Aluminium)
Selecting which species are electrolysed
CALCULATIONS
Minimum voltage required
How much substance is deposited during electrolysis
(or how much time to produce known amount of
substance)
Slide 35-3
This Lecture
Commercial Applications
Voltaic cells
Batteries
Corrosion (and anti-corrosion)
Electrolytic cells
Aluminium electrolysis
refining of Cu
Industrial production of Cl2 and NaOH
Slide 35-4
Batteries: Commercial Uses of Redox
We will examine the operation, construction and
chemistry of 3 classes of batteries:
Primary batteries. These are non-reversible and once
reactants converted to products the battery is “dead”.
Secondary batteries: In these batteries the cell
reaction can be reversed and the battery recharged
Fuel cells: Fuel (chemicals) pass through the battery,
which converts chemical energy into electrical energy.
Slide 35-5
Primary Batteries
The Dry Cell
Construction:
A carbon rod acts as an inert
or inactive cathode.
The rod is coated with MnO2
(oxidant) and graphite powder
(for electrical conductivity).
A paste of NH4Cl and ZnCl2
acts as an electrolyte (like the
salt bridge), although the NH4+
is also reduced.
The Zn metal can act as the
active anode.
Figure 12.16 Blackman
Slide 35-6
Primary Batteries
Chemistry:
Anode: The anode reaction is the same as we have been using: the
oxidation of zinc.
Zn → Zn2+ + 2e-
E 0 = 0.76V
Cathode: The cathode reaction is complex, and not completely
understood. The MnO2 (O.N. = ___) is reduced to Mn2O3 (O.N. = ___)
through a series of steps also involving some acid-base chemistry as
NH4+ (weak acid) to NH3 (weak base)
2MnO2 + 2NH4+ + 2e- → Mn2O3 + 2NH3 + H2O *
E 0 = 0.7 V
Zn(s) + 2MnO2(s) + 2NH4+(aq) → Mn2O3(s) + Zn2+(aq) + 2NH3(aq) + H2O(l)
E 0 = 1.5 V
* Different texts show different cathode reactions. This is an indication of the
unknown cathode chemistry. The Cl- probably plays a role in reacting with Zn to
form crystalline Zn(NH3)2Cl2. When dry cells get swollen it is due to build up of
gases such as NH3.
Slide 35-7
Primary Batteries
Alkaline battery
The alkaline (e.g. “Duracell”)
battery is an improved dry
cell. It uses an alkaline
environment which prevent the
build up of gases, and
preserves the zinc electrode:
Figure 12.17 Blackman
E 0 = 1.25V
Anode:
Zn + 2OH- → ZnO + H2O + 2e-;
Cathode:
2MnO2 + H2O + 2e- → Mn2O3 + 2OH-; E 0 = 0.12V
Overall:
Zn(s) + 2MnO2(s) → ZnO(s) + Mn2O3(s);
E 0 = 1.4 V
Slide 35-8
Primary Batteries
Mercury or silver
battery
Steel button
(cathode)
+
Insulator
E 0(silver) = 1.6 V
Zinc case
(anode)
-
E 0(mercury) = 1.3 V
Porous
separator
Ag2O or HgO in
KOH/Zn(OH)2
paste
Anode:
Cathode (Silver):
Zn + 2OH- → ZnO + H2O + 2e-;
Ag2O + H2O + 2e- → 2Ag + 2OH-;
E 0=1.25V
E 0=0.34V
Cathode (Mercury):
HgO + H2O + 2e- → Hg + 2OH-;
E 0=0.05V
Slide 35-9
Secondary Batteries
Lead-acid battery
Probably the most familiar of all
secondary batteries – this type of
battery has started cars for 100
years.
Figure 12.14 Blackman
Anode:
Pb + HSO4- → PbSO4 + H+ + 2e-;
Cathode: PbO2 + 3H+ + HSO4- + 2e- → PbSO4 + 2H2O;
E 0=0.30 V
E 0=1.63 V
Pb(s) + PbO2(s) + 2H+(aq) + 2HSO4-(aq) → 2PbSO4(s) + 2H2O(l); E 0=1.93 V
6 cells in series
→ 12 V
Slide 35-10
Secondary Batteries
Some features of the
Pb-acid battery:
Large surface area of
electrodes means that the
battery can deliver very large
currents
The reaction is reversible so
the battery can be recharged
Figure 12.14 Blackman
Slide 35-11
Secondary Batteries
Nickel-cadmium batteries
Very popular replacement for primary batteries. The cell potential is the same so
they can be used interchangeably. They can be recharged hundreds of times.
E 0=0.82 V
E 0=0.6 V
Anode:
Cd + 2OH- → Cd(OH)2 + 2e-;
Cathode: NiO(OH) + H2O + e- → Ni(OH)2 + OH-;
Cd(s) + 2NiO(OH) + 2H2O → Cd(OH)2 + 2Ni(OH)2;
E 0=1.4 V
Lithium batteries
Perhaps the batteries of the future because Li is so light, and the cell
potential is one the highest of common batteries. Lithium has the
highest oxidation potential of all, so little Li is needed.
Anode:
Cathode:
Li → Li+ + e-;
MnO2 + Li+ + e- → LiMnO2;
E 0=3.0 V
E 0=0.0 V
Li + MnO2 + → LiMnO2;
E 0=3.0 V Slide 35-12
Fuel Cells
A fuel cell is a voltaic cell where the reactants are a fuel, e.g. H2, CH4.
The fuel undergoes a normal (overall) combustion reaction, however
the two half-reaction are separated and the electrons harnessed.
Fuel cells are still in the experimental stage, and their most notable
success is probably in space.
Anode:
Cathode:
Anode:
Cathode:
H2 → 2H+ + 2e-;
½O2 + 2H+ + 2e- → H2O;
E 0=0.0 V
E 0=1.23 V
H2(g) + ½O2(g) → H2O(g);
E 0=1.23 V
CH4 + 2H2O → CO2 + 8H+ + 8e-;
4 x {½O2 + 2H+ + 2e- → H2O};
E 0=-0.3 V
E 0=1.23 V
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g); E 0=0.92 V
Slide 35-13
Fuel Cells - H2 / O2 fuel cell
Anode:
Cathode:
H2 → 2H+ + 2e-;
½O2 + 2H+ + 2e- → H2O;
H2(g) + ½O2(g) → H2O(g);
E 0=0.0 V
E 0=1.23 V
E 0=1.23 V
Electrolyte: Teflon ( CF2 )n
with RSO3- groups
H+
Figure 12.20 Blackman
Slide 35-14
Corrosion: Unwanted voltaic cells
We have already examined the process of converting a metal
oxide to a metal (L.30). We saw that it required reducing the
oxide and a lot of heat energy. This means that the reverse,
oxidation of a metal to its oxide will be exothermic, and likely
to be spontaneous.
Reduction + energy
Metal
Metal
oxide
Oxidation, spontaneous
Economically, the most important corrosion process is iron or steel.
Slide 35-15
Corrosion
Corrosion occurs commonly in our environment. It is
considered a product-favoured reaction. That is, Eo for
the reaction is positive (a spontaneous process).
Such reactions are
undesirable because the
corrosion results in loss of
structural strength.
For instance, rust holes in
cars, corrosion of
structural supports in
bridges.
Slide 35-16
Corrosion
Fig. from Silb. p.926
1) Oxidation of Fe at active anode forms a pit and yields e- which travel
through the metal;
2) Electrons at the Fe (inactive) cathode reduce O2 to OH-;
3) The Fe2+ migrates through the drop and reacts OH- and then O2
to form rust.
Slide 35-17
Redox chemistry of corrosion
The rusting of iron involves two (or more) redox reactions:
Anode:
Cathode:
2 x {Fe → Fe2+ + 2e-};
O2 + 4H+ + 4e- → H2O;
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l);
E 0=0.44 V
E 0=1.23 V
E 0=1.67 V
The Fe2+ is further oxidised at the edges of the droplet, where [O2] is
highest:
Anode:
2 x {Fe2+ → Fe3+ + e-}
E 0=-0.77 V
Cathode:
½ x {O2 + 4H+ + 4e- → H2O};
E 0=1.23 V
2Fe2+(aq) + ½O2(g) + 2H+(aq) → 2Fe3+(aq) + H2O(l);
E 0=0.46 V
Iron (III) forms a very insoluble oxide, so rust is deposited at the edge:
2Fe3+(aq) + (3+n) H2O(l) → Fe2O3•n H2O(s) + 6H+(aq)
H+ is a
Overall: 2Fe(s) + 3/2 O2(g) + n H2O(l) → Fe2O3•n H2O(s)
catalyst!
Slide 35-18
Chemistry of corrosion
You should now be able to explain some of the
known features of rusting:
Iron does not rust in dry air;
No water ⇒ no “salt bridge”
Iron does not rust in oxygen-free water, such as ocean
depths;
No oxygen ⇒ no oxidant
Iron rusts more quickly in acidic environments;
H+ is a catalyst
Iron rusts more quickly at the seaside.
More conductivity in the “salt bridge”
Slide 35-19
Protection against corrosion
Iron acts as both an active anode and inactive cathode. Anything
that gives up electrons more readily will instead act as anode and
the iron will not oxidise. These sacrificial anodes can be made of
any metal that is more active than iron. If you remember the
“Activity Series of Metals” lecture experiment, you will recall that
both zinc and magnesium are more active than iron. This is called
“cathodic protection”, and is used frequently in large iron structure
such as ships, pipes, bridges, etc
Galvanising:
Water
Mg
rod
Fe
pipe
Zinc
H2O + ½O2
→ 4OH-
Iron
2e−
Zn → Zn2+
Slide 35-20
Corrosion Protection
An example of anodic inhibition would be to prevent oxygen
from reaching the metal. This could be achieved by
greasing (oiling the surface) or allowing a thin film (metal
oxide film) to form on the surface.
A common method is the treatment of a surface with a
solution of sodium chromate.
2Fe(s) + 2Na2CrO4(aq) + 2H2O(l) → Fe2O3(s) + Cr2O3(s) + 4NaOH(aq)
The surface of the iron is oxidised by the chromate salt to
give Fe(III) and Cr(III) oxides. These form a coating that
cannot be penetrated by oxygen and water, which inhibits
further oxidation.
Slide 35-21
Corrosion Protection
Erin
Brockovich
This was the approach used
by PG&E in the small town of
Hinkley in US. Cr(VI) was
used to prevent corrosion of
the power plant’s cooling
towers.
• Unfortunately, the company
then disposed of the Cr(VI)
containing water into
unlined ponds leading to
environmental disaster.
• The result was numerous
reports of cancer, in excess
of 648 law suits and >$10M
of payouts.
Slide 35-22
Electrolytic cell applications:
Production of NaOH and Cl2
Chemicals ranked #9 and #10 in worldwide production are NaOH and
Cl2. Both are produced in the same electrolysis process:
Reduction:
2H2O + 2e- → H2 + 2OHNa+ + e- → Na
E 0 = -(0.41 + 0.6*) V
E 0 = -2.71
Oxidation:
2Cl- → Cl2 + 2e-
E 0 = -1.36 V
2H2O → O2 + 4H+ + 4e* over-potential
E 0 = -(0.82 + 0.6*) V
Overall:
2Na+(aq) + 2Cl-(aq) + 2H2O(l) → Cl2(g) + H2(g) + 2Na+(aq) + 2OH-(aq)
#9
#10
Slide 35-23
Electrolytic cell applications:
Production of NaOH and Cl2
Slide 35-24
Electrolytic cell applications:
Production of Aluminium
Can Al metal be produced from Al3+(aq)?
Al is the third most abundant element in the Earth’s crust, but it
remained a rare and expensive metal because of the stability of the
oxide. In 1886, a 22 year old American (Hall) and 23 year old Frenchman
(Herault) independently developed a method for electrolytic production
of Al metal, that is still used today.
Problems:
-Al can’t be electrolysed from solution because H2O is preferentially
reduced.
- Al can’t be electrolysed from the molten oxide because it melts at too
high a temperature (2045 ºC).
Solution:
Dissolve Al2O3 in hot cryolite (sodium hexafluoroaluminate,Na3AlF6,
MP = 1000 ºC)
Slide 35-25
Electrolytic cell applications:
Production of Aluminium
Anode and cathode reactions are complex. O2 is given off at the anode,
which reacts with the graphite to produce CO2 (as a result the electrodes
are rapidly used up). Very high currents are used (~250,000 Amp) to
produce enough Al on an industrial scale.
Slide 35-26
Electrolytic cell applications:
Refining of Cu
Electro-refining is the principle method
by which Cu is refined to high purity
Less easily reduced metals remain in
solution.
Noble metals are not oxidised so fall to
the bottom as “mud”
Slide 35-27
Summary
CONCEPTS
Difference between primary and secondary batteries
How common batteries work
Chemical reactions for common batteries
Fuel cells and how they work
Corrosion and corrosion protection
CALCULATIONS
None specifically, except redox calculations involving
batteries.
Slide 35-28