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Brown's Chemistry Advanced A Level Notes - Theoretical–Physical
Advanced Level
Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 7
7.5
Electrochemical cells (batteries) and fuel cell systems
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Redox reactions take place in electrochemical cells
('batteries') where electrons are transferred from the reducing agent
(e.g. 'fuel') to the oxidising agent (e.g. O2, MnO2)
indirectly via an external circuit.
A potential difference is created
that can drive an electric current to do work.
Electrochemical cells
have very important commercial applications as a portable supply of
electricity to power electronic devices such as mobile phones, tablets
and laptops. On a larger scale, they can provide energy to power a
vehicle.
So, how do electrochemical cells like simple
batteries work? How does a zinc–carbon battery work? How does a NiCad or
alkaline battery work? How does a lead–acid battery work? How does a
fuel cell work? How does a lithium ion battery work? The advantages and
disadvantages of electrochemical cells - primary cells, secondary cells,
fuel cells, rechargeable cells and non-rechargeable cells etc.
7.5 Electrochemical cells
('batteries') and fuel cell systems
Cells ('batteries') can be:
(i) non-rechargeable
(irreversible) where the chemicals are used up
(ii) rechargeable, where the
chemistry generating the electrical current can be reversed, so
charging up the battery again, AND the chemicals are not used up,
(iii) or fuel cells, which
must be continuously fed chemicals e.g. fuel plus oxidant to
maintain the chemistry producing the electrical current. Fuel cells
used to generate an electric current do not need to be electrically
recharged.
PRIMARY CELLS
are not
rechargeable and are discarded (hopefully by safe recycling systems!)
after they run down when all the chemicals are used up ie no more
chemical potential energy available
SECONDARY CELLS
can be
recharged after they have run down ie the discharge reactions producing
the electricity are reversed to built up the store of chemical potential
energy.
FUEL CELLS
produce electricity
directly from gaseous of liquid fuels such as hydrogen or hydrocarbons
with only safe waste products of water or carbon dioxide.
Important notation
convention:
Note that the +ve and –ve electrode charges in battery cells/fuel
cells are reversed compared to the electrodes in the process of
electrolysis in
electrolytic cells. This is because battery cells and electrolysis cells
operate in 'opposite directions' to each other in terms of oxidation &
reduction and electron flow. Therefore, we have, a –ve cathode in electrolysis
and a +ve cathode in battery/fuel cells AND a
+ve anode in electrolysis and –ve anode in battery/fuel cells. But, as in
electrolysis, you still get reduction at the cathode and oxidation at
the anode, so watch out for the (–) and (+) electrode signs!
-
Primary Cells
-
The first
primary cells were galvanic cells in which the reactants are sealed
in when manufactured and ready for immediate use i.e. the chemicals
are capable of spontaneously reacting and the redox changes released
energy as an electron flow (rather than heat energy).
They cannot be
recharged, and when they run down, that is the chemical reactants
are completely depleted, they stop working and are discarded!
-
The common ones
such as the zinc–carbon batteries are used in
torches, radios, cameras, flashlights, cameras etc.
-
Hopefully
recycling of the materials will be increasingly possible as well as
being worthwhile from the point of view of conserving valuable
resources and minimising environmental pollution from poisonous
metals or their compounds.
-
The Leclanché cell is a primary
cell or battery invented and patented by the French scientist
Georges Leclanché. The battery contains a conducting solution
(electrolyte paste) of ammonium chloride and water, a cathode
(positive terminal) of carbon, a depolarizer of manganese dioxide
(oxidizer), and an anode (negative terminal) of zinc (reductant).
How it works is described and explained below.
-
Dry cell
zinc–carbon battery, 1.5V falling to 0.8V as reaction
products build up.
-
In the zinc–carbon cell
a rod of carbon
cathode (+ convention) is set into a paste of zinc and ammonium chloride (weakly acid
electrolyte) and fine particles of manganese(IV) oxide and carbon
contained in a zinc anode (– convention) 'compartment'. Although called a 'dry'
cell, the paste must contain water, which is thickened with e.g.
starch.
-
Zn(s)|ZnCl2(aq),NH4Cl(aq)||MnO2(s)|MnO(OH)(s)|Cgraphite
-
(–) anode
discharging reaction (i) Zn(s) + 4NH3(aq)
==> [Zn(NH3)4]2+(aq) + 2e–
-
(+) cathode
discharging reaction (ii) MnO2(s) + NH4+(aq)
+ e– ==> MnO(OH)(s) + NH3(aq)
-
overall working
cell reaction (iii) Zn(s) + 4NH3(aq)
+ 2MnO2(s) + 2NH4+(aq)
-
oxidation state
changes: (i) oxidation Zn(0) ==> Zn(+2), (ii) reduction 2Mn(IV) ==>
2Mn(III) to balance
-
Advantages: Low
cost and non–toxic materials.
-
Disadvantages:
Cannot be recycled, can leak (weak acid electrolyte reacts with
zinc), short shelf–life, unstable voltage and current (as battery
'runs down') and low power.
-
The dry cell
alkaline battery, 1.5–1.9V
depending on constituents.
-
In the alkaline dry cell
the electrolyte
is the strong base sodium/potassium hydroxide contained in
'typically' zinc anode (–) compartment and a cathode of
manganese(IV) oxide. Metals like cadmium or aluminium can be used as
the anode, and copper, iron, lead, mercury, nickel and silver oxide
can be used as cathode materials.
-
Zn(s)|ZnO(s)|OH–(aq)||MnO2(s)|Mn(OH)2(s)|Cgraphite
-
(–) anode
discharging reaction (i) Zn(s) + 2OH–(aq)
==> ZnO(s) + H2O(l) + 2e–
-
(+) cathode
discharging reaction (ii)
MnO2(s) + 2H2O(l)
+ 2e– ==> Mn(OH)2(s) + 2OH–(aq)
-
overall cell
reaction (iii) Zn(s) + MnO2(s) + H2O(l)
==> ZnO(s) + Mn(OH)2(s)
-
oxidation state
changes: (i) oxidation Zn(0) ==> Zn(+2), (ii) reduction Mn(IV) ==> Mn(II)
-
Advantages:
-
Disadvantages:
-
Fuels cells
are a development of primary cells but with one
significant difference from their predecessors, the chemical
potential energy source or 'fuel' can be continually fed in to give
the cell a long active life.
-
The hydrogen–oxygen
fuel cell
-
 |
equation |
It uses costly platinum electrodes and
an acid electrolyte such as phosphoric acid, H3PO4 |
|
1. oxidation |
2H2(g)
==> 4H+(aq) + 4e– (at
negative anode
electrode*) |
|
2. reduction |
O2(g)
+ 4H+(aq) + 4e– ==> 2H2O(l) (at
positive cathode electrode*) |
|
3
= 1 + 2 redox |
2H2(g)
+ O2(g) ==> 2H2O(l) |
|
*
Note the +ve and –ve electrode charges are reversed compared to electrolysis, because the system is operating in the
opposite direction. But, as in electrolysis, you still get reduction at
the cathode and oxidation at the anode - it can be confusing
! |
-
The
hydrogen–oxygen cell with an
alkaline electrolyte is known as the 'alkali fuel cell' and
is used in NASA's space shuttle craft.
-
(–) anode reaction
(i)
2H2(g) + 4OH–(aq) ==> 4H2O(l)
+ 4e–
-
(+) cathode reaction
(ii) O2(g) + 4H+(aq) + 4e–
==> 2H2O(l)
-
overall fuel cell
reaction: 2H2(g)
+ O2(g) ==> 2H2O(l)
-
The electrolyte
is the alkali potassium hydroxide solution, KOH(aq).
-
In both acid or
alkaline hydrogen–oxygen fuel cells the oxidation state changes are
-
-
Advantages: Can run on
conventional fuels without the need of expensive metals except for
the catalyst
-
Disadvantages: Quite
costly at the moment eg the platinum catalyst
-
Organic fuel cells are described in
Advanced Redox Chemistry Part III (Organic reactions)
-
Secondary Cells
(electrical 'accumulators')
-
Secondary
cells are galvanic cells that must be charged before they can be
used and rechargeable many times. In the charging process, the
spontaneous–feasible cell reaction that produces electrical energy
is reversed, so building up the chemical potential of the cell
system.
-
Lead–acid
storage battery, 2 V. (usually 6 in series to give 12V
supply).
-
The electrodes
are initially hard lead–antimony alloy plates coated in a paste of
lead(II) sulphate encased in dilute sulphuric acid. During the first
charging some of the lead(II) sulphate is reduced lead(0) on one of
the electrodes (this will acts as the (–) anode in discharging).
Simultaneously in charging, lead(II) sulphate is oxidised to lead(IV)
oxide on the other electrode which acts as the cathode (+) in
discharging.
-
Pb(s)|PbSO4(s)|H+(aq),HSO4–(aq)||PbO2(s)|PbSO4(s)|Pb(s)
-
(–) anode
discharging reaction (i)
Pb(s) + HSO4–(aq)
==> PbSO4(s) + H+(aq) + 2e–
-
(+) cathode
discharging reaction (ii)
PbO2(s) + 3H+(aq)
+ HSO4–(aq) + 2e–
==> PbSO4(s) + 2H2O(l)
-
working cell
reaction (iii) PbO2(s) + 2H+(aq)
+ 2HSO4–(aq) + Pb(s)
==> 2PbSO4(s) + 2H2O(l)
-
oxidation state
changes: (i) oxidation Pb(0) ==> Pb(II) : (ii) reduction Pb(IV) ==> Pb(II)
-
The charging
reactions will be the opposite of (i) and (ii)
-
Advantages:
Inexpensive, high power density (can car starter motor as well as
lights), long shelf life, readily recharges, so has a long working
life of many years.
-
Disadvantages:
Lead needs to be recycled to avoid environmental contamination,
sometimes generates hydrogen gas at the cathode when charging
(explosive in air + spark) – though batteries seem to be made of a high
standard these days in completely sealed units that last many years.
-
Uses: Car
batteries.
-
The NiCad
Cell, 1.25 V.
-
diagram?
-
Cd(s)|Cd(OH)2(s)|KOH(aq)||Ni(OH)3(s)|Ni(OH)2(s)|Ni(s)
-
(–) anode
discharging reaction (i) Cd(s) + 2OH–(aq)
==> Cd(OH)2(s) + 2e–
-
(+) cathode
discharging reaction (ii)
2Ni(OH)3(s) + 2e–
==> 2Ni(OH)2(s) + 2OH–(aq)
-
overall cell
reaction (iii) Cd(s) + 2Ni(OH)3(s) ==>
Cd(OH)2(s) + 2Ni(OH)2(s)
-
oxidation state
changes: (i) oxidation Cd(0) ==> Cd(II), (ii) reduction Ni(III) ==> Ni(II)
-
The charging
reactions will be the opposite of (i) and (ii)
-
Advantages: ?
-
Disadvantages:
Cadmium is a toxic metal.
-
Uses: Portable
computers
-
The lithium cell (the
rechargeable 'Lion battery')
-
A popularly used commercial
battery that powers many a computer!
-
The chemicals involved in one
type are lithium metal, graphite and a lithium-cobalt oxide.
-
The discharging simplified
electrode reactions in a lithium cell to generate electricity are as
follows:
-
At the positive electrode
(reduction): Li+ + CoO2 + e–
==> Li+[CoO2]–
-
Negative electrode via graphite
(oxidation): Li ==> Li+ + e–
-
When the battery is being
recharged, these reactions are reversed.
-
The voltage and power
available from a battery or cell
-
The voltage
depends primarily on the materials used in the chemical process
generating the electrical energy.
-
Since the
voltage is small from an individual cell, (typically 0.4 to 2V),
several cells can be assembled in parallel to increase the voltage.
-
The power
primarily depends on the amount of material and how fast the
chemicals can react.
-
The
advantages and disadvantages of using different types of
electrochemical cells
-
Batteries - commercial
electrochemical cells, have provided us with an extremely useful
source of power, both in the home and industry.
-
They are convenient to carry
around in phones, computers, torches and fitted into a variety of
vehicles.
-
You can take power to remote
places without the need for expensive cabling, and rechargeable
batteries via solar cells can be invaluable in this context e.g.
powering a telephone box in some off the beaten track location in
some way out location in the countryside.
-
The orbiting space station and
spacecraft use fuels cells to provide power and collect the water
produced for domestic use! You can't get a more remote location than
outer space or the moon!
-
BUT, there are drawbacks to all
this electrochemical cell technology and some of the 'pros and cons'
are discussed below.
-
Non-rechargeable cells
-
Non-rechargeable batteries have a
long history and cheap to produce. They can be made in all sorts of
sizes, with a range of voltages and applied to a wide range of
devices from fire alarms to torches.
-
BUT, all the chemicals are 'used
up' and the voltage-power declines as they come towards the end of
their working life.
-
They can't be recharged and are
often just thrown away without any attempt at recycling which is
waste of both energy and materials.
-
Rechargeable cells
-
Using rechargeable cells
obviously reduces the total number of cells used and so far less
will be added to the 'scrapheap' of used 'batteries'.
-
Batteries, such as those in cars
are constantly being recharged when the vehicles is moving and
provide power for lights, starter motor, electric windows and
ignition and any other automated device.
-
Rechargeable cells can be used in
conjunction with solar panels, these work in daylight, but storing
some of the electrical energy enables it to be used at night when
the solar panels are inactive.
-
However, the cadmium in
nickel-cadmium rechargeable cells is toxic and must be carefully
disposed of and the cadmium recovered, hopefully to be recycled.
-
The lead and lead(II) sulfate in
lead-acid car batteries can be recovered and recycled in new
batteries and they are a better design these days and last longer
now in 2017+ than they did back in the 1960s as I recall !
-
The design of lithium-ion cells
has improved over the years and are now widely used in phones,
computers and even larger machines like disability scooters. These
Li-ion batteries are quite costly, but their useful lifetime is
increasing as is the amount of electrical energy they can store
efficiently e.g. several hours use on your computer/phone from a
relatively compact source!
-
Fuel cells
-
The main big advantage of
fuel cells of the hydrogen-oxygen variety is that the only waste
product is harmless water.
-
Theoretically, fuel cells can
replace the power generation from conventional petrol and diesel
engines, thus reducing emission of greenhouse gases and harmful
nitrogen/sulfur oxides and hydrocarbon/carbon particulate
pollution.
-
However, most industrial
hydrogen (e.g. for the Haber synthesis of ammonia) is actually
made from fossil fuel hydrocarbon molecules!
-
Hydrogen is also highly
flammable-explosive with issues over safe storage and
distribution.
-
One possibility is
large-scale solar powered electrolysis of water and collect the
gases, but such an enterprise seems a long way off and the
health and safety issues remain!
-
See
Fuel Cells e.g. the hydrogen - oxygen fuel cell
WHAT NEXT?
Advanced Equilibrium Chemistry Notes Part 1. Equilibrium,
Le Chatelier's Principle–rules
* Part 2. Kc and Kp equilibrium expressions and
calculations * Part 3.
Equilibria and industrial processes * Part 4
Partition between two
phases, solubility product Ksp, common ion effect,
ion–exchange systems *
Part 5. pH, weak–strong acid–base theory and
calculations * Part 6. Salt hydrolysis,
acid–base titrations–indicators, pH curves and buffers * Part 7.
Redox equilibria, half–cell electrode potentials,
electrolysis and electrochemical series
*
Part 8.
Phase equilibria–vapour
pressure, boiling point and intermolecular forces watch out for sub-indexes
to multiple sections or pages
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