Cells and Batteries
Syllabus Statements:
• C.5.1 Describe how a hydrogen–oxygen fuel cell
works. (Include the relevant half-equations in
both acidic and alkaline electrolytes.)
• C.5.2 Describe the workings of rechargeable
batteries. ( Include the relevant half-equations.
Examples should include the lead–acid storage
battery, the nickel–cadmium (NiCad) battery and
the lithium-ion battery.
• C.5.3 Discuss the similarities and differences
between fuel cells and rechargeable batteries.
Fuel Cells
• A fuel cell is a device that converts chemical
energy directly into electrical energy.
• The most widely used fuel cell is a
hydrogen/oxygen fuel cell.
• This is used in the space programme!!
• The electrodes are made of porous carbon
impregnated with a catalyst
• (either Pd or Pt for the negative electrode; Pt
for the positive electrode)
• Notice that I’ve been very careful not to talk
about anode and cathode.
• I’ll explain why later!
• The chemical reactants are supplied from an
external source (gas tanks!) directly into the
electrodes.
• The electrodes are surrounded by an
electrolyte of hydroxide solution (KOH or
NaOH)
• Let’s see a diagram!
• At the –ve electrode:
• Hydrogen reacts to give water and surplus
electrons
• 2H2 (g)+ 4OH-(aq)  4H2O(l) + 4e• Everything has been doubled to make it easier
to balance the overall equation!
• The hydrogen has been oxidised
• Why?
• The electrons generated flow round the circuit
to the other electrode.
• O2(g) + 2H2O(l) + 4e-  4OH-(aq)
• The oxygen is reduced
• Why?
• What is the overall reaction?
• 2H2(g) + O2(g)  2H2O(l)
• Notice that the hydroxide concentration isn’t
changed.
• This type of fuel cell generates about 1.23V
• The hydroxide electrolyte can be replaced
with acidic conditions
• In this case the half equations are:
• At the –ve electrode:
• 2H2(g)  4H+(aq) + 4e• At the +ve electrode:
• O2(g) + 4H+(aq)  2H2O(l)
• Make sure you can give either set of half
equations
• A small aside!
• Why was I so sneaky about not
mentioning anode and cathode?
• Well when you were told that the
anode was the +ve electrode and the
cathode was the –ve electrode,
• We lied to you!
• The anode is the electrode where oxidation
occurs.
• In electrolysis this is the +ve electrode
• But in cells, this is the –ve electrode!!
• The cathode is the electrode where reduction
occurs
• In electrolysis this is the –ve electrode
• But in cells this is the +ve electrode
• You will study this more in the oxidation and
reduction module
• See what I did there?
• Eh – oh = A O = anode oxidation
Advantages and Disadvantages of Fuel
Cells
• Advantages:
• Fuels cells are very efficient compared to
conventional cells (70 – 80% efficient)
• They produce no greenhouse gases
• No “thermal pollution”
• The water produced can be drunk (useful in
the space program)
• They are light weight
• Disadvantages:
• Gases ( H2 ; O2 )are hard to store and handle
• The H2 is often produced from the electrolysis
of water, which requires fossil fuels to be
burned.
• They often experience technical problems
(leaks, corrosion, failure of the catalyst)
• They are ****** expensive!
Summary
Rechargeable Batteries
• Primary cells produce electricity through
chemical reactions BUT CANNOT BE
RECHARGED.
• These are discussed in detail in the oxidation
and reduction topic.
• Secondary cells CAN be recharged
• You need to be aware of the details of
– The Lead Acid Battery
– Nickel Cadmium Batteries (NiCad)
– Lithium Ion Batteries
The Lead Acid Battery
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Used in cars
Anode in made of lead plates
Cathode is made of lead(IV) oxide
The electrolyte is sulphuric acid
5.2 mol / dm3 since you asked!
• The half equations are:
• Pb(s) + SO42-(aq)  PbSO4(s) + 2e• Is this the negative electrode or positive
electrode?
• Negative – because it generates electrons
• Is it the anode or the cathode?
• It is the anode because the lead is oxidised (it
loses electrons)
• And . . .
• At the positive electrode:
• PbO2(s) + 4H+(aq) + SO42-(aq) + 2e-  PbSO4(s) +
2H2O(l)
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•
•
•
Lead is reduced – why?
Goes from Pb(IV) to Pb(II)
So this is the
Cathode
• What’s the overall reaction:
• Pb(s) + PbO2(s) + 2H2SO4(aq)  2PbSO4(s) +
2H2O(l)
• Uses up sulphuric acid
• The electrolyte gets less concentrated
• The condition of the battery can be checked
by measuring the strength of the acid
• We don’t bother titrating it (phew!) -
• The whole point of a rechargeable battery is
that it can be recharged (DUH!)
• This is done by passing electricity through it
• PbSO4 + 2e-  Pb + SO42• Pb(II)  Pb(0) Lead is reduced
• PbSO4 + 2H2O  PbO2 + 4H+ + SO42- + 2e• So what’s the overall reaction?
• 2PbSO4(s) + 2H2O(l)  Pb(s) + PbO2(s) + 2H2SO4(aq)
• Notice that this uses up water and regenerates
sulfuric acid
• Lead and lead oxide are also regenerated
• In practice the battery is usually charged by the
alternator (a small generator) whilst it is still in
the car.
• Each cell produces 2V; car batteries are 12V, so 6
cells are connected in series to make a car battery
• Advantages:
• Easily recharged
• Can deliver a large amount of energy for a
short time
• Disadvantages:
• Heavy
• Acid can spill
NiCad batteries
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•
•
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Known as “dry cells”
Produce about 1.4V
The electrolyte is KOH
They are called “dry cells” because the
electrolyte is either soaked onto a paper
separator or made into a paste
• At the –ve electrode
• Cd(s) + 2OH-(aq)  Cd(OH)2(s) + 2e•
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•
•
Cd is oxidised (loses electrons)
Hence this is the anode
At the +ve electrode
NiO(OH)(s) + H2O + e-  Ni(OH)2 + OH-
• Ni is reduced
• Ni(III)  Ni(II)
• Therefore this is the cathode
So what’s the overall reaction?
• Cd + 2NiO(OH) + 2H2O  Cd(OH)2 + 2Ni(OH)2
• (All solids - except water!)
• The reducing and oxidising agents are
regenerated by recharging
• Cd(OH)2 + 2Ni(OH)2  Cd + 2NiO(OH) + 2H2O
• Remember - only the half equations show the
electrons. We cancel them out in the overall
reaction!
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Advantages:
Light
Easy to transport
Long life
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Disadvantages:
Expensive (compared to lead - acid batteries)
Lower voltage than lead – acid batteries
Cd is toxic and must be disposed of carefully
Memory effect
• Memory effect:
• If a NiCad cell is recharged without being
completely discharged, then an unreactive
surface can form on the electrodes, which can
stop the recharging
Lithium Ion Batteries
• These are now used in laptops, mobile phones
etc.
• They are complicated; high tech; prone to
bursting into flames!
• Li is a reactive metal and should be able to
generate lots of electrical energy
• BUT Li quickly gets covered in an oxide layer,
which stops it making contact with an
electrolyte, so the cell won’t work.
• To overcome this problem, lithium ion cells
don’t contain lithium metal
• They contain mobile lithium ions
• The –ve electrode is made of graphite
• Lithium ions can enter the carbon lattice to
form LiC6
• The +ve electrode is a metal compound e.g
MnO2 or CoO2 or NiO2
• At the –ve electrode
• LiC6  Li+ + 6C + e• (oxidation, hence anode)
• The electrolyte is usually an organic solvent
which can carry the lithium ions to the +ve
electrode
• Li+ + e- + MnO2  LiMnO2
• This is reduction
• What’s the overall reaction?
• LiC6 + MnO2  6C + LiMnO2
• The cell is powered by the overall movement of
lithium ions
• The process can be reversed by passing a current
in the opposite direction
• The big advantage is that there is no memory
effect!
• They generate about 3.6V
• This technology is as new to me as it is to you!
• If you want to know more you’ll have to research
it yourself!!!
Discuss the similarities and differences
between fuel cells and rechargeable
batteries
• Similarities:
• Both convert chemical energy into electrical
energy
• Both make use of spontaneous redox
reactions
• Differences:
• Batteries are energy storage devices; fuel cells
are energy conversion devices
• Fuel cells require a constant supply of
reactants. Batteries are a closed system
• Batteries can be recharged – but cannot
generate electricity whilst being recharged.
Fuels cells can operate continually and so have
a longer operating life.
• Fuel cells have inert electrodes.
• Fuel cells are more expensive!!!!!
Did we address all the syllabus
statements?
• C.5.1 Describe how a hydrogen–oxygen fuel
cell works. (Include the relevant halfequations in both acidic and alkaline
electrolytes.)
• C.5.2 Describe the workings of rechargeable
batteries. ( Include the relevant halfequations. Examples should include the lead–
acid storage battery, the nickel–cadmium
(NiCad) battery and the lithium-ion battery.
• C.5.3 Discuss the similarities and differences
between fuel cells and rechargeable batteries.