How do hydrogen fuel cells work?

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Like batteries, hydrogen fuel cells can transform chemical energy into electrical energy. The difference between batteries and fuel cells is that batteries store the chemical reactants inside their cases whereas in a fuel cell the chemical reactants come from outside the cell, as shown in Figure 1a and 1b, respectively. This important characteristic makes fuel cells, in theory, able to operate indefinitely if they have a continuous supply of hydrogen and oxygen.  Internal combustion engines (ICE) or heat engines also convert chemical energy into electricity throughout several steps. The fuel releases its energy as a thermal energy in a combustion chamber, followed by conversion of the thermal energy into mechanical energy and finally the mechanical energy into electricity - see Figure 1c. In comparison to heat engines, fuel cells do not involve energy conversion steps before they produce electricity, fuel cells have no moving parts like electric generators or combustion chambers that work at high temperatures and are not limited by the Carnot cycle¹. Instead, they convert chemical energy directly into electricity at a relatively low temperature.

 

 

 

The are many types of fuels that can be used in a fuel cell such as methane, propane, or alcohols but hydrogen is considered the best option because, unlike the others, it does not generate carbon dioxide greenhouse gas but only water, heat, and electricity.  Although hydrogen is practically everywhere and exists in vast quantities, it is usually attached to other molecules and its extraction is an energy intense and costly process. Despite this, hydrogen fuel, to be used in a fuel cell, is our best alternative to generate clean electrical energy and reduce the effects of greenhouse gases which lead to climate change.  To add to this, the advantage of hydrogen as an energy carrier, is that it can be generated by renewable energy sources such as photovoltaics or wind power by water electrolysis.

 

 

Principles Behind Fuel Cell Operation

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When hydrogen molecules (H₂) enter the fuel cell, as shown in Figure 1b, they permeate through an anode electrode consisting of a microporous carbon network that supports a gas diffusion layer and a catalyst layer on the opposite side. The carbon network and the gas diffusion layer facilitate the even distribution of hydrogen towards the catalyst layer where the hydrogen oxidation reaction (HOR) to protons (H⁺) occurs, releasing electrical charges in the form of electrons - see Figure 2. Each mol of molecular hydrogen provides two mol electrons or, in engineering terms, two Faraday charges. The HOR is relatively straightforward if the appropriate catalyst such as platinum (Pt) or platinum-ruthenium (PtRu), are used. There are other lower cost catalysts such as nickel (Ni) and nickel oxide (NiO), iron (Fe) and iron oxide (Fe₂O₃), cobalt (Co), and ceria (CeO₂), tungsten carbide (WC), molybdenum carbide (Mo₂C), strontium titanate (SrTiO₃), lanthanum strontium manganite (La₀.₇Sr₀.₃MnO₃) and many others. The typical loading catalyst at the anode varies but it can be between 0.04 to 10 mg cm^-2, depending on the manufacturing method and application².

 

 

The electrons created by the HOR flows, via an electronic conductor material outside the cell, towards the cathode and provides the power for a motor, light or any suitable load. The cathode electrode receiving the flow of electrons has a similar construction to the anode electrode, i.e., gas diffusion and catalyst layers and it is where the oxygen reduction reaction (ORR) occurs - see Figure 3. Typical catalysts for the ORR are platinum (Pt), platinum-iron (Pt-Fe), platinum-cobalt (Pt-Co), iron-nitrogen-carbon (Fe-N-C), manganese oxides (MnO_x), cobalt-nitrogen-carbon (Co-N-C), and many others. Perhaps one of the most important differences between HOR and ORR is that the oxygen reduction is thermodynamically more difficult than the hydrogen oxidation reaction and is substantially slower in comparison. The ORR is one of the major limitations in hydrogen-oxygen fuel cells and, in a way, these energy losses and the sluggishness of the reaction have limited wider commercialisation of fuel cells. Major efforts have been dedicated to developing a catalyst able to facilitate and increase the rate reaction of oxygen reduction.

 

 

 

Despite their similarity in geometry and construction, the anode and cathode electrodes operate in different chemical environments and are under mechanical, thermal, and chemical electrochemical stress and degradation during the fuel cell operation. Another fundamental part of the hydrogen fuel cell is the electrolyte that separates the anode and cathode electrodes. This is typically an ionic solution of sodium hydroxide or a phosphoric acid, for example, which conducts ions but not electrons. However, it is more suitable to use a thin solid polymer electrolyte membrane because this minimises the distance between the two electrodes, the movement of hydrogen ions (H+) towards the cathode is shorter and allows scaling up of the fuel cell. Typical proton exchange membranes (PEMs) measure 50 micrometres thickness and only allow the passage of ions, not electrons. Figure 4 shows a unit cell composed of one anode and one cathode divided by a solid polymer electrolyte membrane. This unit cell is the membrane electrode assembly (MEA) and it is at the heart of hydrogen fuel cells. Typical MEA thickness is 500 micrometres but can change depending on the system. An additional component of the fuel cells are the graphite bipolar plates. These plates hold the MEA together and electrically connect adjacent MEAs in a fuel cell stack, provide heat transfer between the cells and serve as cooling separators. The bipolar plates have flow fields machined on their surface that permit the movement of reactants (hydrogen and oxygen) and products (water) to and from the electrodes. Figure 5 shows an MEA between two bipolar plates on each side. This unit repeats many times to form a fuel cell stack of up to 200 cells or more.

 

 

 

The cell voltage across the anode and cathode E_cell, follows the equation below:

Where E⁰_c - E⁰_a are the electrode potentials of the ORR and HOR respectively, at the theoretically thermodynamic equilibrium, i.e. when not current has been generated, while ∑|η|-∑iR are the sum of the overpotentials and the sum of the potential drop caused by the ionic and electrical resistances of all the cell components, respectively. The terms can be expanded as follows:

η_{c, act} and η_{a, act} are the activation overpotentials, i.e., the potential required to initiate the cathodic and anodic reactions above the equilibrium potential of each reaction, i.e., E⁰_c and E⁰_a. The following terms η_{c, conc} and η_{a, conc}, are the mass transport overpotentials, i.e., the limitations when the hydrogen or oxygen supply is not sufficiently fast to satisfy the current demand and they do not reach the electrode surface to react fast enough, causing a drop in the cell potential. The activation and mass transport overpotentials control the reaction at low and high currents, respectively. On the other hand, the equation that represents the potential drop due to the ionic and electrical resistances ∑|iR| shows the importance of selecting components with low ionic and electrical resistance R, i.e., the membrane electrolyte, current collectors, bipolar plates, wires connectors in the electrical circuit, and all the components of the hydrogen fuel cell.

 

 

Fuel Cell Performance

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The fuel cell generates water, electricity, and heat from the oxidation of hydrogen and the reduction of oxygen in a system like the one described in Figures 2 and 3 and performs the electrochemical reactions written below:

In an ideal hydrogen fuel cell where the overpotentials and the resistance are negligible, i.e., the term
∑|η|-∑iR = 0, the cell potential should be 1.23 volts which is the electrode potential of the cathodic reaction minus the electrode potential of the anodic reaction at the equilibrium. However, as soon as the electrodes are connected the voltage drops and when the current flows the overpotentials start to influence the cell potential which falls below the thermodynamical predicted value of 1.23 volt. The cell potential falls with the current as shown in the curve of Figure 6, known as the polarisation curve.

Figure 6 shows various scenarios of the cell voltage versus the current density. The ideal cell voltage for a single cell would remind 1.23 V as the current density increases if the terms ∑|η|-∑iR were zero as shown at the top of the area highlighted in green. However, this is an ideal behaviour, and in real fuel cells, one needs to consider the limitations of the gas permeation, short circuit, and ORR overpotential, i.e., the green area in Figure 6 that shows how the cell voltage drops as the current density increases. The overpotential due to the HOR, represented by the yellow area, is in general quite low as the current increases and its effect on the cell voltage drop is small. The area in red represents the influence of the resistance whereas the area in blue represents the overpotential due to the mass transport limitations, with the current density. It is clear from the figure that the gas permeation and short circuiting the two electrodes plus the ORR contributes substantially to the volage drop although at very high current densities; the mass transport limitations also cause a high cell voltage to drop.

In general, the cell voltage decreases as the current density demands are larger. To avoid excessive cell voltage drops most fuel cells operate at moderate current densities in what is known as the ohmic region between 200 and 800 mA cm^-2, in the example shown in Figure 6.

 

 

Advantages and challenges

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Hydrogen fuel cells have properties that make them an attractive option for a range of applications. The advantages are many; these include the potential for high efficiency, being highly scalable to larger sizes to provide high power, use of other fuels than hydrogen although with the caveat of carbon dioxide emissions, otherwise zero or near zero carbon dioxide emissions, no moving parts, and the potential to run continuously, unlike batteries. However, there are still many challenges, the main one perhaps is that hydrogen is not widely available and can be costly. In an ideal case hydrogen should come from renewable sources of energy, reduce the cost of the MEA of which the main expenses are the catalysts on the electrodes and the membrane. As the fuel cell performance gradually decreases with time core research is required in catalyst degradation and membrane electrolytes. At the stack level there are other challenges that include optimisation of the air supply management (most fuel cells use air rather than pure oxygen), and water management system (humidified hydrogen and oxygen), fuel supply system (blowers and ejectors), and coolant pumps and fans.

 

 

Fuel Cell Applications

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Fuel cells are needed to replace conventional fossil fuel energy generation. There are many markets where fuel cells have found applications, for example in portable electronic equipment due to their high energy density; in the transport industry car manufacturers have already produced fuel cell cars prototypes; in the aerospace industry hydrogen fuel cells are interesting due to being lightweight, having no moving parts and having reduced carbon dioxide levels although having hydrogen stored in aircrafts comes with many challenges for its potential fire hazard. 

 

 

 

 

Conclusion

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Fuel cells operate with two gas diffusion electrodes divided by a proton exchange membrane. One of the most important characteristics of fuel cells is the relationship between the cell current and the cell voltage shown in the polarisation curve. This relationship graphically represents the influences of the kinetic (rate of HOR and ORR), the ohmic resistances and the mass transport of fuel and oxidant on the performance. In fuel cell stacks the material and energy balances as well as the water and thermal balances are key factors to increase their efficiency. Despite its disadvantages, fuel cells are one of the best routes to achieving a clean energy supply.

 

 

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References

1. For details about the Carnot Cycle see: the Khan Academy, Khanacademy
2. Yan-Jie Wang, David P. Wilkinson, and Jiujun Zhang. Noncarbon Support Materials for Polymer Electrolyte Membrane Fuel Cell Electrocatalysts.
3. Unknown Author is licensed under CC BY-SA
4. Journal of Power Sources Volume 258, 15 July 2014, Pages 189-194. Sciencedirect
5. First London fuel cell bus
6. German Aerospace Center is licensed under CC BY 2.0. https://www.dlr.de/en
7. https://www.horizonfuelcell.com/#!minipak/c156u
8. Fuel Cell Store