Fuel Cells

• Fuel cells convert chemical energy directly into electrical energy, operating in a way that is comparable to batteries. However, unlike batteries, they do not run out or require recharging as long as there is a constant supply of fuel and oxygen.
• They are based on the principles of redox reactions. Typically, in a hydrogen fuel cell, hydrogen gas (H2) acts as the anode, while oxygen gas (O2) acts as the cathode.
• At the anode, hydrogen gas is oxidised to produce protons (H+ ions) and electrons. This reaction can be represented as: H2 -> 2H+ + 2e-
• At the cathode, oxygen gas undergoes reduction by taking up the electrons and protons formed at the anode to produce water. This reaction can be represented as: O2 + 4H+ + 4e- -> 2H2O
• The overall cell purpose is the conjunction of these two half-reactions. The overall redox reaction can be represented as: 2H2 + O2 -> 2H2O
• In hydrogen fuel cells, the electrolyte is often a proton exchange membrane (PEM) that only permits protons to pass through it. Thus, the electrons are forced to travel through an external circuit, producing a flow of electric current.
• Fuel cells are often used in applications where a continuous source of power is needed, such as in spacecraft and hospitals. They are efficient and clean sources of energy as the by-product is simply water, making them an attractive option for future energy supply.
• Various fuels can be used in fuel cells, including hydrogen, methanol, and natural gas. The choice of fuel depends on the application and the specific design of the cell.
• Limitations of fuel cells include high costs, the requirement for pure fuels to avoid catalyst poisoning, and challenges associated with storing and transporting fuels safely, especially hydrogen.
• Understanding the oxidation number changes and half-cell reactions in a fuel cell can further illuminate the principles of redox reactions, equipping you with the knowledge necessary for tackling related calculation and theory-based questions.