The hydrogen economy myth-part 1

Fuel_Cell_Apparatus-_Pressurized_Liquid_ReactantsBeen having a bit of a blether about the hydrogen economy recently after a supposed breakthrough in hydrogen production at Glasgow University.  This got me thinking about this whole hydrogen economy myth which we covered in our book.  I’m afraid these two posts are going to be complicated so if you don’t like science leave now…

For the purposes of these posts we will mainly be talking about fuel cells and more specifically proton exchange membrane fuel cells (PEMFC’s).  These have a platinum anode and cathode separated by a semi-permeable membrane (usually TEFLON with sulphonate groups added).  The platinum acts as a catalyst.  To make hydrogen we use an electrolyser.  In this the platinum anode and cathode are in the same compartment.

At the anode the reaction (technically an oxidation reaction) is;

2 H2O(l) → O2(g) + 4 H+(aq) + 4e

At the cathode (technically a reduction reaction) is ;

2 H+(aq) + 2e → H2(g)

combining these two and balancing gives;

2 H2O(l) → 2 H2(g) + O2(g)

In the fuel cell all the above are reversed.  At the anode hydrogen is split by the catalyst to protons and electrons.  The electrons flow through the conductor joining the two and some are used by us to do “useful work”.  At the cathode oxygen, electrons from the circuit and protons which have migrated though the semi-permeable membrane combine to make water.  The idea is that the semi-permeable membrane allows the protons to migrate to the cathode compartment where it combines with the hydrogen and oxygen in the reverse reaction but does not allow the oxygen to enter the anode compartment.

All chemical reactions have three criteria that are vital.  The first is kinetics, that is the speed of the reaction.   The second is the thermodynamics of the reaction.  This is whether the reaction is energetically favourable.  Very importantly thermodynamics also has a bearing on the maximum efficiency of any chemical reactions and how much energy can be produced.  The third is where the chemical equilibrium lies i.e. whether the reaction intrinsically favours formation of the products or the reverse.  Both the forward and reverse reactions above require a catalyst.  This speeds up the reaction (kinetics), but has no bearing on the thermodynamics of whether the reaction is favourable.  Most reactions are reversible.  As reactants proceed to products if the products are not removed the reverse reaction will occur more and more until potentially an equilibrium is reached with no overall change.  Equilibrium however can be related to thermodynamics.

Splitting water to hydrogen is not energetically favourable and spontaneous change will not occur, therefore energy in the form of electrons has to be provided.  The reverse reaction is energetically favourable and it produces energy in the form of electrons.

To finish part 1 of this post, thermodynamics has two laws (actually three but the third is irrelevant here) and three components (enthalpy (H), entropy (S) and free energy (G).

The first law says that energy can neither be created or destroyed.  This has a number of  connotations but one is that it does not matter how you get from reactant to product H, S and G are the same by whichever route you go.

The second states that no spontaneous change occurs unless the entropy in a system increases.

The three components are heat (self-explanatory often reactions produce or require heat).  Entropy or disorder is slightly harder to visualise.  The greater the disorder the greater the entropy.  If you have gas in a cylinder its entropy is low, open a valve and its entropy increases to the maximum extent possible as gas molecules exit chaotically  filling the space.  Free energy is the amount of energy available to do “useful work”, in our example above the electrons liberated as hydrogen is split into protons and electrons.

The three are related by the following equation;


Δ means sum of the change over the reaction between the start and finish and where T is the temperature (usually taken as 25ºC) at which a reaction occurs.  For spontaneous change to take place ΔG must be negative.  So in our example above;

H2 + O2 -> H2O

must have a negative value of ΔG.  The more the equilibrium moves towards the products the more negative ΔG will be (up to a theoretical maximum).  (I won’t bore you with the equation linking chemical equilibrium and thermodynamics.)  Equilibrium cannot ultimately trump thermodynamics in making a reaction occur, but can influence it.  Reaction conditions can be tweaked to maximise all the above, using other compounds to “donate” free energy or heat or pressure.  In part 2 of the hydrogen economy myth I will look at maximum and practical efficiencies and why in my view the hydrogen economy doesn’t add up.


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