Redox and fuel cells

In this lesson, we will learn:

• The redox reaction used to produce energy in a hydrogen-oxygen fuel cell.
• The variation of the redox reaction for acidic and basic conditions.
• The advantages and disadvantages of using the hydrogen-oxygen fuel cell as an energy source.

• When you calculate cell potential for a possible redox reaction (like in Predicting redox reactions using cell potential), getting a positive value tells you that this reaction is feasible and it results in a net release of energy. You are converting chemical potential energy into electrical energy using hydrogen gas – it is an energy source.
  This is a growing area of research where electricity can be generated to power cars and other vehicles, replacing the use of fossil fuels.
  This lesson will look at the specific redox reactions involved, the advantages and the disadvantages of using hydrogen fuel cells.


• The best-known redox reaction as a source of energy uses the oxidation of hydrogen by oxygen to produce water. Looking straight from the standard potentials table1 we can see this is feasible:
• O₂ + 4H⁺ + 4e^-$\, \rightleftharpoons \,$ 2H₂O E⁰ = +1.23 V
• 2H⁺ + 2e^- $\, \rightleftharpoons \,$ H₂ E⁰ = +0.00 V

We are reducing oxygen and oxidising hydrogen, so the reaction at each electrode is:
• Oxidation at anode: (H₂ $\,$→$\,$ 2H⁺ + 2e^- ) x 2
• Reduction at cathode: (O₂ + 4H⁺ + 4e^- $\,$→$\,$ 2H₂O ) x 1

The overall reaction, then, with species cancelled out:

2H₂ + O₂ + 4H⁺ $\,$→$\,$ 2H₂O + 4H⁺
And the cell potential is going to be:

E⁰_cell = +1.23 + 0.00 = 1.23 V
In other words, it is feasible for oxygen to oxidise hydrogen gas and produce water. The reaction will produce a voltage that can power cars and other vehicles. When used as a source of power, these are batteries known as hydrogen-oxygen fuel cells.


• This reaction can be run in acidic or basic conditions.
• The example above is in acidic conditions – note the presence of H⁺ in the redox half-equations.
• Oxidation at anode: H₂ $\,$→$\,$ 2H⁺ + 2e^-
• Reduction at cathode: O₂ + 4H⁺ + 4e^- $\,$→$\,$ 2H₂O

The oxidation of hydrogen gas at the anode produces H⁺ and electrons which will flow in the same direction toward the cathode.
At the cathode, the reduction of oxygen uses electrons to produce the H₂O waste product.
See the diagram below:


• The other option to this is in alkaline conditions. This can be done in the presence of a strong base like KOH or NaOH.
• Oxidation at anode: H₂ + 2OH^- $\,$→$\,$ 2H₂O + 2e^-
• Reduction at cathode: O₂ + 2H₂O + 4e^- $\,$→$\,$ 4OH^-

The oxidation of hydrogen gas here uses hydroxide ions to produce water, the redox product. These hydroxide ions are provided by the electrolyte and the reduction of oxygen at the cathode.
See the diagram below:


• There are several advantages of using hydrogen-oxygen fuel cells:
• This process converts the chemical potential energy of H₂ to electrical energy without large amounts of wasted heat energy.
• The oxygen-hydrogen fuel cell could replace existing battery fuel-cells, which often contain toxic heavy metals.
• The hydrogen-oxygen fuel cell is less polluting as it produces only water, unlike combustion of fossil fuels which creates CO₂, SO₂ and NO_x gases as well.
• These fuel cell batteries do not require charging up – they can run as long as hydrogen fuel and oxygen is still supplied to the cell.


• The biggest drawbacks to the use of hydrogen fuel cells are the following:
• Hydrogen gas has largely non-renewable sources: hydrogen is rarely found in its elemental (H₂) form where it is used in fuel cells.
  It therefore needs to be sourced from other substances like hydrocarbons, which are sourced from fossil fuels. We also need hydrogen to make ammonia which is extremely important for food production. Ultimately, the hydrogen-oxygen fuel cell still has non-renewable sources.
• Large-scale production: There is still no reliable, efficient, large-scale method to produce the hydrogen gas we need for these fuel cells.
  The H₂ gas will likely come from fossil fuels, or the electrolysis of water. Electrolysis of water is the exact opposite of this redox reaction, so it requires the exact voltage that this redox cell releases.
  That required electricity will have to be put in, probably from fossil fuel burning, so we are likely no better off than just using the fossil fuels anyway once we consider (transporting and storing the hydrogen we mass-produced).
• Storage issues: As a gas, the volume of space hydrogen takes up is quite large compared to the same amount of a liquid fuel like petrol/gasoline. Fuel capacity is substantially lower because of this and refueling would have to happen frequently. This is solved (slightly) by pressurizing H₂ gas but that leads to safety concerns.
• Safety concerns: Hydrogen gas is highly flammable and explosive. The simple issue of storing gases economically is to pressurize them. This is a major safety risk for flammable gases, however.



¹Source for data: ATKINS, P. W., & DE PAULA, J. (2006).?Atkins' Physical chemistry. Oxford, Oxford University Press.