Introduction to Chemical Cells and Fuel Cells
Welcome! In this chapter, we are exploring how we can take chemical energy and turn it into electrical energy. This is the science behind everything from the tiny battery in a TV remote to the advanced fuel cells powering space shuttles. If you have ever wondered why some batteries can be recharged while others can't, or how we can produce electricity while only emitting water, you are in the right place!
1. Chemical Cells and Batteries
A chemical cell is a simple system that produces a voltage (a potential difference) until the chemicals inside are used up.
How a Simple Cell is Made
To make a basic cell, you only need three things:
1. Two different metals (called electrodes).
2. An electrolyte (a liquid or gel that contains ions and can conduct electricity).
3. A wire to connect the metals.
The Science Secret: The two different metals have different reactivities. One metal "wants" to give up electrons more than the other. These electrons flow through the wire from the more reactive metal to the less reactive metal, creating an electric current!
Predicting the Voltage
The voltage produced by a cell depends on:
• The types of electrodes: The bigger the difference in reactivity between the two metals, the higher the voltage. For example, a cell made of Magnesium and Copper will have a higher voltage than a cell made of Zinc and Copper.
• The electrolyte: Different ions in the solution will affect how well the electrons flow.
Quick Review Box:
• Cell: A single unit using chemical reactions to make electricity.
• Battery: Two or more cells connected together in series. This provides a greater total voltage (you just add the voltages of the individual cells together!).
Rechargeable vs. Non-Rechargeable Cells
Non-rechargeable cells: (Like alkaline batteries). The chemical reactions are irreversible. Once one of the reactants is used up, the reaction stops, and the battery "dies."
Rechargeable cells: The chemical reactions can be reversed when an external electric current is supplied. This resets the chemicals so they can react again.
Key Takeaway: Voltage is created by the difference in metal reactivity. Batteries are just multiple cells joined together to boost the power.
2. Fuel Cells
Fuel cells are different from normal batteries. While a battery stores a set amount of energy inside it, a fuel cell is supplied by an external source of fuel (like hydrogen) and oxygen.
How they work
In a fuel cell, the fuel is oxidised electrochemically. This means the fuel reacts with oxygen at a low temperature to produce a voltage.
• The most common type is the hydrogen fuel cell.
• The only waste product of a hydrogen fuel cell is water (\(H_2O\)).
The Overall Equation:
\(2H_2 + O_2 \rightarrow 2H_2O\)
Analogy: Think of a normal battery like a lunchbox—once you've eaten the food inside, it's empty. A fuel cell is like a restaurant—as long as food (fuel) keeps being delivered to the kitchen, it can keep serving customers (producing electricity) forever!
Did you know? Hydrogen fuel cells are used in spacecraft because they provide electricity and the "waste" product (pure water) can be used by the astronauts for drinking!
Hydrogen Fuel Cells vs. Rechargeable Batteries
Exam questions often ask you to evaluate which is better. Here is a quick comparison:
• Fuel Cells: They don't need recharging (as long as you have fuel), they don't get less efficient over time, and they are non-toxic. However, hydrogen is a gas, which is hard to store and can be explosive.
• Rechargeable Batteries: They are easy to store and use, but they eventually wear out, take a long time to recharge, and can contain toxic chemicals that harm the environment when thrown away.
Key Takeaway: Fuel cells need a constant supply of fuel but are very clean and efficient compared to traditional batteries.
3. Higher Tier (HT Only): Half Equations
If you are taking the Higher Tier paper, you need to know what happens at the individual electrodes inside a hydrogen fuel cell. Don't worry if this seems tricky; just follow the electrons!
Memory Aid: OIL RIG
Oxidation Is Loss (of electrons).
Reduction Is Gain (of electrons).
At the Negative Electrode (Anode)
Hydrogen is oxidised (it loses electrons):
\(2H_2 \rightarrow 4H^+ + 4e^-\)
At the Positive Electrode (Cathode)
Oxygen is reduced (it gains electrons and reacts with the \(H^+\) ions):
\(O_2 + 4H^+ + 4e^- \rightarrow 2H_2O\)
Common Mistake to Avoid: In electrolysis, the anode is positive and the cathode is negative. However, in a chemical cell, it is the opposite! The negative electrode is where the electrons are released (oxidation).
Key Takeaway: Hydrogen loses electrons at the negative electrode, and oxygen gains them at the positive electrode to form water.
Summary Checklist
Can you:
• Explain how to make a simple chemical cell?
• State why the reactivity of metals affects the voltage?
• Explain the difference between rechargeable and non-rechargeable batteries?
• Write the overall equation for a hydrogen fuel cell?
• Compare the advantages of fuel cells versus batteries?
• (HT Only) Write the two half equations for a hydrogen fuel cell?