Power from Chemicals: Understanding Batteries and Fuel Cells

Hey everyone! Ever wondered how the battery in your remote control works? Or how some new, high-tech cars can run on hydrogen and only produce water? It's all about chemistry! In this chapter, we're going to dive into the amazing world of chemical cells. We'll learn how simple chemical reactions can be harnessed to create electricity. We will explore three main types:

  • The everyday Zinc-Carbon cell (your classic AA battery).
  • Special cells that use non-reactive inert electrodes.
  • The futuristic Hydrogen-Oxygen fuel cell.

This topic is super important because it's the foundation of all portable electronics and a key part of our search for cleaner energy. Don't worry if it sounds complicated – we'll break it down step-by-step!


1. The Zinc-Carbon Cell: The Everyday Battery

You've definitely used one of these. The zinc-carbon cell, also known as a dry cell, is the standard, non-rechargeable battery you find in torches, clocks, and toys. It's a perfect example of a primary cell.

What's a Primary Cell?

Think of it like a single-use disposable camera. You use it until the chemicals inside are used up, and then you have to throw it away (responsibly, of course!).
A primary cell is a non-rechargeable cell. The chemical reaction inside it cannot be easily reversed.

Inside a Zinc-Carbon Cell

Let's look at what's inside. It might seem complex, but it's just a few key parts working together.

  • Negative Electrode (Anode): The outer casing is made of zinc (Zn). This is where oxidation happens.
  • Positive Electrode (Cathode): A carbon (graphite) rod runs down the centre. It's inert, meaning it doesn't react. It just acts as a site for reduction to occur and to conduct electrons.
  • Electrolyte: This isn't a liquid, which is why it's called a "dry cell". It's a moist paste containing ammonium chloride (NH₄Cl) and manganese(IV) oxide (MnO₂).
Memory Aid: OIL RIG and AN OX RED CAT

To remember what happens where, use these mnemonics:
OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
AN OX: Anode is where Oxidation occurs.
RED CAT: Reduction occurs at the Cathode.

How It Works: The Chemistry

The magic happens through a redox reaction.

Step 1: At the Negative Electrode (Anode)

The zinc casing is more reactive, so it loses electrons. This is oxidation. As the battery is used, the zinc casing gets thinner! This is why old batteries can sometimes leak.

Equation:

$$ \mathbf{Zn(s) \rightarrow Zn^{2+}(aq) + 2e^-} $$

Step 2: At the Positive Electrode (Cathode)

The electrons released by the zinc travel through your device (like a lightbulb) to the carbon rod. Here, ions in the electrolyte paste gain these electrons. This is reduction.

Equation:

$$ \mathbf{2NH_4^+(aq) + 2e^- \rightarrow 2NH_3(g) + H_2(g)} $$

The manganese(IV) oxide (MnO₂) is also very important. It's a depolariser that reacts with the hydrogen gas produced, preventing a build-up of gas that would stop the battery from working.

Key Takeaway: Zinc-Carbon Cell

A zinc-carbon cell is a cheap, common, primary (non-rechargeable) battery. It produces electricity (~1.5V) because the zinc casing gets oxidised (loses electrons), and chemicals in the electrolyte paste get reduced (gain electrons) at the carbon rod.


2. Chemical Cells with Inert Electrodes

Sometimes, the chemicals that react to make electricity are all dissolved in the solution as ions. They aren't solid metals that we can use as electrodes. So, what do we connect the wires to? We use inert electrodes!

What are Inert Electrodes?

Inert electrodes are materials that conduct electricity but do not take part in the chemical reaction.
Analogy: Think of a referee in a football match. They are on the field and control the game (conduct electrons), but they don't score goals or defend (they don't react).
The most common inert electrodes are graphite (carbon) and platinum. They are used because they are unreactive and good conductors.

Example: The Iron(II)/Iron(III) and Iodide/Iodine Cell

Let's build a cell to see how this works. We'll connect two different half-cells.

  1. Half-Cell 1: A beaker containing a solution of both iron(II) ions (e.g., from FeCl₂) and iron(III) ions (e.g., from FeCl₃). We dip a platinum electrode into it.
  2. Half-Cell 2: A beaker with a solution of iodide ions (e.g., from KI) and iodine (I₂). We dip another platinum electrode into this one.
  3. The Connection: We connect the two electrodes with a wire and a voltmeter. To complete the circuit, we connect the two solutions with a salt bridge (e.g., a strip of filter paper soaked in potassium nitrate solution).
Quick Review: The Salt Bridge

The salt bridge is essential! It allows ions to move between the beakers to balance the charge. Without it, charge would build up and the flow of electrons would stop almost instantly.

Predicting the Reaction (Using the Electrochemical Series)

So which way do the electrons flow? We need to consult the Electrochemical Series (ECS). The ECS tells us which substance is a stronger oxidising agent (more likely to be reduced).

  • The `Fe³⁺(aq) / Fe²⁺(aq)` system is a stronger oxidising agent than the `I₂(aq) / I⁻(aq)` system.
  • This means `Fe³⁺(aq)` will pull electrons from the other half-cell. In other words, `Fe³⁺(aq)` will be reduced.
  • Therefore, `I⁻(aq)` must be oxidised.

Writing the Equations

Now we can write the half-equations for each electrode.

At the Negative Electrode (Anode): Oxidation

Iodide ions lose electrons to become iodine. This happens at the surface of the platinum electrode.

$$ \mathbf{2I^-(aq) \rightarrow I_2(aq) + 2e^-} $$

Observation: The colourless solution starts to turn brown as iodine is formed.

At the Positive Electrode (Cathode): Reduction

Iron(III) ions gain electrons to become iron(II) ions.

$$ \mathbf{Fe^{3+}(aq) + e^- \rightarrow Fe^{2+}(aq)} $$

Observation: The yellow/brown solution of Fe³⁺ turns into the pale green solution of Fe²⁺.

The Overall Ionic Equation

To get the overall equation, we need to make sure the electrons cancel out. We need to multiply the cathode equation by 2.
`2Fe³⁺(aq) + 2e⁻ → 2Fe²⁺(aq)`
Now, we add the two half-equations together and cancel the electrons (`2e⁻`) from both sides:

$$ \mathbf{2I^-(aq) + 2Fe^{3+}(aq) \rightarrow I_2(aq) + 2Fe^{2+}(aq)} $$
Key Takeaway: Inert Electrode Cells

When the reactants are ions in solution, we use inert electrodes (like platinum or graphite) to conduct electrons. We use the Electrochemical Series to predict which substance will be oxidised (at the anode) and which will be reduced (at the cathode).


3. Fuel Cells: Clean Energy for the Future

Imagine a battery that never runs out, as long as you keep adding fuel to it. That's a fuel cell! It’s an electrochemical cell that converts the chemical energy from a fuel and an oxidising agent directly into electricity.

Analogy: A normal battery is like a sealed lunchbox – once you eat the food, it's empty. A fuel cell is like having a chef who keeps bringing you more food as you eat – you can keep going as long as the chef (fuel supply) is there!

The Hydrogen-Oxygen Fuel Cell

The most famous type is the hydrogen-oxygen fuel cell.

  • Fuel: Hydrogen gas (H₂)
  • Oxidising Agent: Oxygen gas (O₂) from the air
  • Only Product: Water (H₂O)!
Did you know?

Hydrogen-oxygen fuel cells were used on the NASA Space Shuttle and Apollo missions. They provided both electricity for the spacecraft and clean drinking water for the astronauts!

How It Works (in an Alkaline Electrolyte)

The cell contains two electrodes separated by an alkaline electrolyte, like potassium hydroxide (KOH).

At the Negative Electrode (Anode): Oxidation

Hydrogen gas is fed in. It reacts with hydroxide ions from the electrolyte and releases electrons.

$$ \mathbf{H_2(g) + 2OH^-(aq) \rightarrow 2H_2O(l) + 2e^-} $$

At the Positive Electrode (Cathode): Reduction

Oxygen gas is fed in from the other side. It takes the electrons (which have travelled through the external circuit) and reacts with water to make more hydroxide ions.

$$ \mathbf{O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq)} $$

The Amazing Overall Equation

If we balance and combine the half-equations, everything except hydrogen, oxygen, and water cancels out. The overall reaction is simply:

$$ \mathbf{2H_2(g) + O_2(g) \rightarrow 2H_2O(l)} $$

This is the same reaction as burning hydrogen in oxygen, but instead of releasing all the energy as heat, a large portion is converted directly into useful electrical energy.

Pros and Cons of Hydrogen Fuel Cells

Advantages (Pros)
  • Very high efficiency: They convert a higher percentage of chemical energy into useful electrical energy compared to traditional combustion engines.
  • Non-polluting: The only product is pure water. There are no greenhouse gases (like CO₂) or pollutants (like NOx or SO₂) produced by the cell itself.
  • Continuous power: They can run continuously as long as fuel (H₂) and oxygen are supplied.
Disadvantages (Cons)
  • Cost: The electrodes require expensive catalysts, like platinum, to work efficiently.
  • Hydrogen production: Most hydrogen is currently produced from fossil fuels, which creates pollution. Making it cleanly from the electrolysis of water requires a lot of electricity.
  • Hydrogen storage and safety: Hydrogen is a very flammable gas and is difficult to store safely and compactly in a vehicle.
Key Takeaway: Fuel Cells

A fuel cell generates electricity by continuously reacting a fuel (like hydrogen) and an oxidant (like oxygen). The hydrogen-oxygen fuel cell is highly efficient and clean, producing only water. However, its widespread use is limited by the high cost and the challenges of producing and storing hydrogen fuel safely.