Electrochemistry, featuring electrolysis and fuel cells.

3.3 Electrochemistry

We got most of the way through this subtopic and I asked you to go through the remainder for homework.  I thought the electrolysis section might be worthy of a post, as we didn’t look into it in great detail. So here goes…….

Key Ideas from the Subject Outline:

Electrochemical cells are conveniently divided into galvanic cells, which produce electrical energy from spontaneous redox reactions, and electrolytic cells, which use electrical energy from an external source to cause a non-spontaneous chemical reaction.

Galvanic and electrolytic cells involve oxidation at the anode and reduction at the cathode, with electrons being transferred from one electrode to the other through an external circuit.

Electrolytic cells are used in the production of active metals.

In electrolytic cells an external source of energy is used to bring about a non-spontaneous redox reaction.

To summarise what happens:

• The power source pumps electrons through the external circuit from the anode (+) to the cathode (-).
• In the electrolyte, the cations (+) are attracted to the negatively charged cathode and the anions (-) are attracted to the positively charged anode.  This movement of ions within the electrolyte constitutes an electric current within the internal circuit.
• If the electrolyte is in aqueous solution, water molecules will be in contact with the surface of the electrodes.
• An oxidation half-reaction occurs at the anode.  For a molten electrolyte, the anions are oxidised.  For an aqueous electrolyte, either the anions or the water molecules are oxidised, depending on their relative ease of oxidation.
• A reduction half-reaction occurs at the cathode.   For a molten electrolyte, the cations are reduced.  For an aqueous electrolyte, either the cations or the water molecules are reduced, depending on their relative ease of reduction.

So in an aqueous electrolyte, how can you tell whether the anions/cations are oxidised/reduced in preference to water?

At the cathode.

Reduction.

If the cation present is zinc or below on the reactivity series, then it is reduced in preference to water.

If the cation present is aluminium or above on the reactivity series, then water is reduced in preference to the cations.

At the anode.

Oxidation.

If the anion present is chloride, bromide or iodide, then it is oxidised in preference to water.

If the SO₄^2- or NO₃^– anion is present, then water is oxidised.

I’m sure that you will find this animation useful.

But why would electrolysis be used in real life?

Highly reactive metals (group 1 and 2) can be produced by the electrolysis of molten chlorides of the metals using inert electrodes (graphite or steel).

The following animation shows the extraction of aluminium by electrolysis.

So, the similarities between galvanic cells and electrolytic cells:

• the ions in the electrolyte conduct electricity within the cell.
• the anions move towards the anode and the cations move towards the cathode.
• oxidation occurs at the anode.  Reduction occurs at the cathode.
• in the external circuit, electrons travel through the wire from the anode to the cathode.

But the differences?

• galvanic cells produce a spontaneous redox reaction.  Electrolytic cells require an external energy source for the non-spontaneous redox reaction to occur.
• in a galvanic cell the anode is negative and the cathode is positive.  In an electrolytic cell it is the other way around.
• in a galvanic cell the electrons flow from the negative electrode to the positive electrode (in the external circuit).  In electrolytic cells the external power source acts as an ‘electron pump’ driving electrons to the cathode, and withdrawing electrons from the anode.
• galvanic cells convert chemical energy to electrical energy.  Electrolytic cells convert electrical energy to chemical energy.
• galvanic cells usually use a salt bridge to connect 2 separate half cells.  In electrolytic cells the electrodes are immersed in a common electrolyte with no salt bridge.

And now it’s quiz time.

Side note:

We haven’t covered fuel cells, but will do so shortly.  Fuel cells are a special type of galvanic cell.  Gaseous fuels are fed continuously to the anode (-) and an oxidant such as oxygen is fed continuously to the cathode (+). Oxidation occurs at the anode and reduction occurs at the cathode to produce an electric current.

So what makes them so special?

• the electrodes are porous which increases their surface area. They are also impregnated with a catalyst.  This will increase the reaction rate.
• fuel (usually hydrogen gas) is continuously supplied to the anode.  At the same time an oxidant (usually oxygen gas) is continuously supplied to the cathode.
• the anode is connected externally to the cathode by a metal wire.
• the operating temperatures are quite high, between 50 – 1000 degrees celsius, dependent on the electrolyte used.

There are a number of advantages of using fuel cells to produce energy (compared to using traditional galvanic cells or steam turbines):

• they continuously provide an electric current as long as a fuel and oxidant is continuously supplied.
• they are efficient.
• they offer a better mass:power output ratio than conventional galvanic cells.
• they use readily available fuels and oxidants.
• they don’t produce pollutant gases.
• the electrodes and electrolyte are not consumed.
• electrode reaction products are removed as they are formed.
• they require minimal maintenance.
• they are silent during operation (no moving parts).

But the disadvantages?

• impurities in the fuel or oxidant can ‘poison’ the electrode.
• high purity fuels and oxidants are costly.
• medium to high temperatures are required for the cells to function effectively.
• metal electrode catalysts (palladium and platinum) are costly.
• some of the electrolytes are very corrosive.

An animation for the operation of a fuel cell can be located here.

And just in case you need some revision for galvanic cells, you can view this animation.

Happy revising, I can feel a test coming up…..