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 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!

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?

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?

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

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

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 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

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) 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 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- 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!

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

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 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 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.