1. Cells and Batteries

2. 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.

3. 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!!

4. 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!

5. The chemical reactants are supplied from an external source (gas tanks
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!

7. 2H2 (g)+ 4OH-(aq)  4H2O(l) + 4e-
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?

8. O2(g) + 2H2O(l) + 4e-  4OH-(aq)
The electrons generated flow round the circuit to the other electrode.
O2(g) + 2H2O(l) + 4e-  4OH-(aq)
The oxygen is reduced
Why?

9. 2H2(g) + O2(g)  2H2O(l) What is the overall reaction?
Notice that the hydroxide concentration isn’t changed.
This type of fuel cell generates about 1.23V

10. 2H2(g)  4H+(aq) + 4e- O2(g) + 4H+(aq)  2H2O(l)
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

11. 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!

12. 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

13. See what I did there?
Eh – oh = A O = anode oxidation

14. Advantages and Disadvantages of Fuel Cells
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

15. 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!

16. Summary

17. 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

18. The Lead Acid Battery

20. 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!

22. 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 . . .

24. At the positive electrode:
PbO2(s) + 4H+(aq) + SO42-(aq) + 2e-  PbSO4(s) H2O(l)
Lead is reduced – why?
Goes from Pb(IV) to Pb(II)
So this is the
Cathode

25. What’s the overall reaction:
Pb(s) + PbO2(s) + 2H2SO4(aq)  2PbSO4(s) H2O(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!) -

26. 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+ + SO e-
So what’s the overall reaction?

27. 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

28. Advantages:
Easily recharged
Can deliver a large amount of energy for a short time
Disadvantages:
Heavy
Acid can spill

29. NiCad batteries 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

30. At the –ve electrode
Cd(s) + 2OH-(aq)  Cd(OH)2(s) + 2e-
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

31. 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!

32. Advantages:
Light
Easy to transport
Long life
Disadvantages:
Expensive (compared to lead - acid batteries)
Lower voltage than lead – acid batteries
Cd is toxic and must be disposed of carefully
Memory effect

33. 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

34. 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.

35. 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

37. 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?

38. 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!!!

39. Discuss the similarities and differences between fuel cells and rechargeable batteries
Both convert chemical energy into electrical energy
Both make use of spontaneous redox reactions

40. 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!!!!!

41. Did we address all the 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.