CHEMICAL REACTIONS FOR A HYDROGEN E CONOMY .
At present, hydrogen production is dominated by the manipulation of hydrocarbons. These same
organic compounds already form the basis of our current, flawed industry, which requires energy
from a diminishing supply of non-renewable resources. If we are to extricate ourselves from a
dependence on fossil fuels, such feedstocks are exactly what a hydrogen economy should avoid.
The hydrogen economy is the conception of using hydrogen as our main fuel source. Since hydrogen
does not naturally occur in its diatomic form H2, in order to realise this vision a chemical reaction
must be found that allows hydrogen production to take place quickly, cheaply and safely. Thus far,
four methods have been proposed that could fulfil the criteria; steam reforming, electrolysis,
photolysis and thermochemical splitting. Yet none of these equations are complete. Not one can
produce hydrogen in a way that could compete with fossil fuels nor provide sufficient fuel for the
planet - so our hydrogen economy remains just a vision, for now.
By looking more closely at the changes occurring in each reaction, the reasons for their shortcomings
can be identified. In each process, a reactant, condition or waste product does not fit requirements
and by highlighting these faults, it becomes clear in which direction current research is headed – and
how close we are to a perfect reaction.
The failings of the most common reaction – hydrocarbon steam reforming – are substantial.
Nonetheless, the steam reforming feed materials are composed of light hydrocarbons, such as
methane, from natural gas and LPG, which are abundant and low-cost. These gases undergo two
reactions to produce hydrogen. The primary reaction is the conversion of the hydrocarbon into
hydrogen and carbon monoxide through a reversible reaction with water. Water must be present as
steam for this reaction to occur with desirable conversion occurring at 700-1000°C. Because carbon
monoxide is a poisonous gas, the standard replacement is a straightforward conversion into carbon
dioxide. The water gas shift reaction produces the CO2 along with more hydrogen in another
reversible reaction. Lower temperatures move the position of equilibrium to the right because the
reaction is moderately exothermic, however in order for a homogenous reaction to take place, both
reactants must be at least 200°C. With that the case, they react as gases to produce synthesis gas.1
CH4 + H2O → CO + 3 H2
CO + H2O → CO2 + H2
reforming reaction
water gas shift reaction
Hydrogen can be extracted by pressure swing adsorption. As the gas passes over an adsorbent, gas
molecules will form Van der Waal’s forces with the solid, leading to an accumulation of gas particles
on the solid surface. Any gas that does not form intermolecular forces will pass through the unit,
although increasing the pressure can cause weak adsorbates to be taken up more readily. In this
case, hydrogen will not adsorb due to its high volatility and lack of polarity. Within the pressure
swing adsorption unit, a solid surface with an affinity for CO2 will remove carbon dioxide and other
impurities such as CO from the syngas, leaving pure hydrogen. The carbon dioxide is then desorbed.2
This inevitably releases carbon dioxide into the atmosphere, where it accumulates, contributing to
global warming. While it lowers emissions compared to combustion of these fuels, it is not a longterm solution for our energy crisis; the limited supply of reactants are being consumed, the
conditions require energy to reach the desired temperatures and the by-products consist of
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greenhouse gases. Inevitably, research into carbon neutral methods of hydrogen production have
been undertaken.
The aim to eliminate the release of compounds that are detrimental to the environment led to the
reconsideration of a method proposed in 1966. The concept of containing reactions in a cycle was
conceived to prevent any unwanted products from leaving the reaction chamber, instead forming
part of a subsequent reaction. These are the thermochemical cycles, all of which receive their
hydrogen source in the form of water – an abundant, harmless molecule.
Over 200 different thermochemical cycles have been tested using different reagents and reactions,
and the exploration of these continues. Water splitting is conducted at high temperatures by
converting water into hydrogen and oxygen via a series of reactions with other species. To date, the
most efficient cycle is the sulfur-iodine cycle, in which iodine and sulfur act as catalysts to promote
water splitting. Other than the oxygen and hydrogen, the chemicals are continuously recycled, which
reduces waste.3
1.) 4H2O + 2SO2 + 2I2 → 2H2SO4 + 4HI
2.) 2H2SO4 → 2H2O + 2SO2 + O2
3.) 4HI → 2I2 + 2H2
FIGURE 1: The sulfur-iodine cycle
In reaction 1, the reactants are mixed to form sulfuric acid and hydrogen iodide. This is an
exothermic reaction, which occurs spontaneously to form products with lower bond enthalpies than
the reactants. The HI is distilled and separated from the sulfuric acid, and each product is
transferred to a different chamber. The ensuing reactions are endothermic; they use more energy
forming the bonds within the new compounds than could be provided by breaking the bonds of the
original compounds. As a result, large amounts of heat are required to provide sufficient energy. At
850-900°C, the sulfuric acid thermally decomposes into water, sulfur dioxide and oxygen vapours
(reaction 2). Water and sulfur dioxide can be condensed and oxygen gas will be released.
Meanwhile, reaction 3 takes place: the hydrogen iodide is heated to 300°C where it will decompose
to form iodine and hydrogen. Because the iodine will condense to form a liquid at a higher
temperature than hydrogen, upon cooling, H2 will remain a gas that can be tapped off from the top
of the cooling chamber.4
Despite the heat consumption, this reaction can be carbon zero. By using nuclear power or solar
concentrators that focus sunlight, temperatures of up to 2000°C can be created. The drawback is
infact the reactants used in reactions 2 and 3; H2SO4 is a strong acid and highly corrosive. Equally,
hydrogen iodide is one of the most corrosive acid gases (formed from the covalent bonding of
hydrogen and a halide). Through a process known as hydrogen embrittlement, it causes loss of
ductility and tensile strength of metals. Hydrogen atoms can arise spontaneously from the HI and
react to form H2. Due to the high temperature, however, the hydrogen atoms are likely to dissolve
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into the matrix of the metal reaction container. So, the reduction of hydrogen atoms to H2 takes
place within the metal, rather than on the surface, forming pockets of gas. As pressure builds up, the
metal becomes weaker and can crack.5 The development of a material that can withstand both the
high temperatures and the corrosive effect has not yet been achieved.
An alternative way to carry out water splitting, and at the same time remove all unnecessary
products and intermediates, is electrolysis. Water splitting is not thermodynamically favourable at
RTP therefore energy is required to drive the reaction. This energy can be supplied by creating a
potential, and in electrolysis this is obtained using two electrodes of opposite charge (figure 2). The
negative dipole in water causes an attraction between the positive electrode (anode) and the water;
four electrons are transferred from the water molecules to the electrode. This decomposes water
into an oxygen molecule and four protons. Since each electrode is attached to a power terminal, the
donation of electrons completes the circuit and so an equal number of electrons are displaced to the
cathode where they attract the protons to form hydrogen.6
Overall redox reaction: 2H2O → O2 + 2H2
FIGURE 2: Electrolysis requires electrodes, over
which a potential difference is placed.
Its role in electron transferral means the metal coating on the electrodes must have the correct
metal-oxygen binding strength. At the anode, oxygen will bond to the metal atom and release an
electron and proton, creating a hydroxide atom. The loss of another electron and proton leaves only
oxygen bound to the metal. Two adjacent, bonded oxygen molecules must then form covalent bonds
to become O2. For this to occur, the bond between oxygen and electrode cannot be too strong or
too weak.7 Platinum is the traditional metal catalyst used, since it binds appropriately with the
oxygen in water in order to encourage the reaction, but use of platinum is not feasible for large-scale
use due to its high cost and scarcity.
A recent analysis of materials at MIT has revealed that the binding affinities of the catalyst are due
to the arrangement of valance electrons and the way in which they bond with oxygen. This can now
be assessed in order to estimate performance before the more specific trials are done, saving time.
It involves mapping how the electrons are distributed within orbitals, with an ideal configuration as a
guide. This also allows the possibility to tailor chemical compositions with more favourable
properties (such as durability) into effective catalysts by changing their electron configuration, giving
opportunity for more abundant metal oxide catalysts to be introduced in the near future. 8
However, this alone will not make the redox equation employable on a large scale because of the
potential needed to carry out such a reaction. The major flaw in this equation is the reaction
conditions, and this includes energy usage as well as catalysts. Theoretically, 1.23V is all that is
needed for this particular reaction, however some energy will inevitably be lost to heat or
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resistance. Any voltage exceeding the ideal (1.23V) is known as overpotential. Even with platinum,
the overpotential is too great to be cost-effective. I believe that bio-inspired catalyst is needed.
It is promising that nature has already verified this way of evolving oxygen by oxidation.
Photosynthesis takes place in plants at a continually high rate, without access to mains electricity,
and therefore it should be possible to achieve the same synthetically. By investigating water-splitting
in plants, a new process - photolysis –is being developed, using the same equation under a different
procedure.
2H2O → O2 + 4H+ + 4e- oxidation
4H+ + 4e- → 2H2
reduction
Photolysis requires light and a catalyst to achieve water splitting. Photons striking an adapted
photocell can be used to provide energy for oxygen evolution and a further step unites electrons
with hydrogen.
Unlike electrolysis, the light splits water directly without needing to convert into electricity, which
increases production and minimises energy wastage. It also avoids hydrocarbon fuel, which is used
in the production of electricity. In this way, it makes a suitable alternative to electrolysis, and
scientists developing the process hope to reduce the overpotential that has plagued this redox
reaction by exploring the catalysts used by plants. The reduction step can already be achieved with
negligible overpotential; it is the initial oxidation that requires the correction.
Plants use the protein complex known as Photosystem II (PSII) to extract protons and electrons for
ATP synthesis, releasing oxygen as a waste product. Research into the mechanism of this reaction
within PSII could be the key to producing our own catalyst. As of this year, the oxygen-evolving
region has been located to a cofactor composed of manganese, calcium and oxygen, with a
bicarbonate ion bound to this cluster that has been identified as the active site for the reaction. The
structure has been mapped at the stage just before oxygen production to show four manganese
atoms held by oxide bridges to a central calcium ion. As shown in figure 3, there are five changes
that occur to the inorganic cofactor, and a greater understanding of these changes should show
where our own catalysts could be modified for efficiency.9
FIGURE 3: A mechanism of PSII. The lightning bolts represent a photon. B The structure of
the Mn4CaO5 cluster.
The identification of manganese as the active metal element is already a major step as manganese is
abundant and easy to access. Plants receive the metal from naturally occurring manganese minerals
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but until recently, it was not known how plants engineer such effective catalysts from this mineral.
At neutral pH, manganese is inactive, and yet this is the form in which water is present. Because of
this, the synthetic catalyst MnO2 demands high energies to induce an interaction between the
manganese and water. On the other hand, PSII will remove electrons from water easily, using a
recently discovered method known as proton-coupled electron transfer (PCET).
MnO2 is an ionic compound formed from manganese4+ and two oxygen2-. Its interaction with water
causes the manganese to accept an electron creating Mn3+. Due to instability, it disproportionates
into Mn2+ and Mn4+. In this state, water is bound as a ligand (Mn2+ OH2) and it is here that PCET is
employed. As the oxidation reaction begins, an electron is emitted causing the transition from Mn2+
back to Mn3+. In order to create this charge imbalance, a large amount of energy must be added to
the system. A subsequent emission of a proton, which has the opposite charge to an electron,
counters this imbalance. This results in Mn2+ OH, and so the first electron and proton of this four
electron/four proton reaction has occurred.10 Yet by using PCET, plants limit the amount of energy
needed. It is achieved by synchronising the proton and electron transfer, so both are released at
once. This maintains a neutral charge, unlike when the proton and electron are transferred
sequentially.11
PCET lowers activation energy and increases rate of reaction; initial trials have shown a ‘significant
improvement in the catalytic activity of manganese oxides at neutral pH.’12 The Mn2+ OH2 complex
must lose both electrons and protons in a single concerted step in order to regenerate the
manganese with minimal overpotential and produce the electrons and protons of the oxidation
reaction.
The opinion is that PSII uses PCET to avoid charge build-up in the oxidation reaction. Manganese
oxides on their own however, are unable to manage this mechanism and so additional compounds
need to be attached that serve as functional analogues for the component of PSII that regulates
PCET. With these compounds currently being tested, photolysis may prove to be practicable in the
near future. This would mean an equation with abundant reactants, efficient conditions and safe
products.
This is not to say that photolysis is the answer to our fuel crisis; there are many other factors that
must be considered if we are to rebuild our industries on hydrogen. In fact, all the methods of
hydrogen production that have been mentioned are still evolving. There may be changes to
reactions which cannot be predicted or there may be new reactions altogether. Some methods have
further to go than others and there is no guarantee what the hydrogen generators of the future will
look like once logistics and engineering costs have been taken into account. Perhaps developments
in engineering will make thermochemical water splitting possible with no changes. Indeed the
perfect reaction for a hydrogen economy may not be just one equation – the solution may be a
combination of many.
Having looked more closely at the existing reactions that produce hydrogen, the depth and detail of
chemical research has been revealed; chemists are researching mechanisms ever more precisely,
overcoming challenges and testing innovative ideas. Yet if we truly wish to create a sustainable
society, we must continue to advocate such reactions that fully preserve our environment, by being
mindful of both reactants and by-products. This is why I believe that an understanding of nature’s
fuel systems will advance the development of our own fuel system that can work alongside nature.
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Bibliography
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Implementing Agreement [Online] p.3 Available from:
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2.) The linde engineering group (2012) hydrogen recovery by pressure swing adsorption
3.) http://energy.gov/eere/fuelcells/hydrogen-production-thermochemical-water-splitting
4.) BANERJEE A. et al. (October 2007) Studies on sulfur-iodine thermochemical cycle for
hydrogen production BARC newsletter [Online] p.68 Available from:
http://www.barc.gov.in/publications/nl/2007/200710-8.pdf
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7.) http://www1.lsbu.ac.uk/water/electrolysis.html
8.) http://phys.org/news/2011-06-catalyst-materials-fuel-cells-metal.html
9.) http://www.mpg.de/8373743/photosynthetic-water-splitting
10.) YAMAGUCHI A. et al. (June 2014) Regulating proton-coupled electron transfer for
efficient water splitting by manganese oxides at neutral pH Nature Communications [Online]
Available from:
http://www.nature.com/ncomms/2014/140630/ncomms5256/full/ncomms5256.html
11.) MAYER J. (June 2004) Proton-coupled electron transfer: A Reaction Chemist's View Annual
review [Online] Available from:
http://www.annualreviews.org/doi/abs/10.1146/annurev.physchem.55.091602.094446
12.) http://www.riken.jp/en/pr/press/2014/20140630_1/
FIGURE 1: pubs.rsc.org
FIGURE 2: nmsea.org
FIGURE 3: journal.frontiersin.org
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