Batteries, fuel cells and corrosion

Cells in a box

Open a torch and you find a battery. It is a galvanic cell, or several in a line, packed so that it can be carried and will not spill. Batteries come in two kinds:

• A primary battery can be used only once. The reaction cannot be reversed by sending current back through it. A torch cell is an example.
• A secondary battery can be recharged. A current from outside runs the cell reaction backwards, so the battery is as good as new. A car battery is an example.

The dry cell

The common torch cell is the dry cell, also called the Leclanché cell. It looks like a small tin can with a carbon rod in the middle.

• The anode is the zinc can: $\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$.
• The cathode is the carbon rod, surrounded by a paste of manganese dioxide and carbon. The reaction is $\text{MnO}_2(s) + \text{NH}_4^+(aq) + e^- \rightarrow \text{MnO(OH)}(s) + \text{NH}_3(g)$.
• The electrolyte is a moist paste of ammonium chloride and zinc chloride. It is not a flowing liquid, which is why it is called dry.

The cell gives about $1.5\ \text{V}$. As it works, the zinc can becomes thin, and a very old cell can leak. The voltage falls steadily during use.

The mercury cell

A small, flat button cell in a watch or a hearing aid is often a mercury cell.

• Anode: zinc, mixed with mercury. Cathode: mercury(II) oxide with carbon. Electrolyte: a paste of potassium hydroxide and zinc oxide.
• The overall reaction is $\text{Zn(Hg)} + \text{HgO}(s) \rightarrow \text{ZnO}(s) + \text{Hg}(l)$.

The electrolyte does not change in this reaction. No ions are used up, and so the voltage stays steady at about $1.35\ \text{V}$ for the whole life of the cell. The drawback is that mercury is poisonous, so these cells must be disposed of with care.

The lead storage battery

The battery in a car is a lead storage battery. Each cell has a lead plate as the anode and a plate of lead dioxide as the cathode. Both stand in sulphuric acid, about 38 per cent by mass.

While the battery discharges:

• Anode: $\text{Pb}(s) + \text{SO}_4^{2-}(aq) \rightarrow \text{PbSO}_4(s) + 2e^-$
• Cathode: $\text{PbO}_2(s) + \text{SO}_4^{2-}(aq) + 4\text{H}^+(aq) + 2e^- \rightarrow \text{PbSO}_4(s) + 2\text{H}_2\text{O}(l)$

Add them and you get $\text{Pb} + \text{PbO}_2 + 2\text{H}_2\text{SO}_4 \rightarrow 2\text{PbSO}_4 + 2\text{H}_2\text{O}$.

Look at what happens. Both plates become covered with lead sulphate. The acid is used up and water is made, so the acid becomes weaker and its density falls. A garage measures the density with a hydrometer to see how charged the battery is. A full cell has a density near $1.3\ \text{g/mL}$, and a nearly flat one has less than $1.2\ \text{g/mL}$.

Each cell gives about $2\ \text{V}$. Six of them in series give $12\ \text{V}$.

To recharge, a car's generator drives current the other way. The cell becomes an electrolytic cell and the reactions reverse: lead sulphate on one plate goes back to lead, and on the other it goes back to lead dioxide. The acid is made again.

Other rechargeable cells

• Nickel-cadmium cell. The anode is cadmium and the cathode is nickel(III) oxide hydroxide. The overall reaction is $\text{Cd} + 2\text{NiO(OH)} + 2\text{H}_2\text{O} \rightarrow \text{Cd(OH)}_2 + 2\text{Ni(OH)}_2$. It gives about $1.4\ \text{V}$ and can be recharged many times.
• Lithium-ion cell. Lithium ions move between a graphite anode and a metal oxide cathode, and back again when the cell is charged. It is light and gives about $3.7\ \text{V}$, which is why phones use it.

Fuel cells

A battery holds its fuel inside, and it runs out. A fuel cell is fed from outside, so it keeps running as long as the gases are supplied.

In the hydrogen and oxygen fuel cell, hydrogen goes to one porous electrode and oxygen to the other. The electrolyte is potassium hydroxide solution.

• Anode: $2\text{H}_2(g) + 4\text{OH}^-(aq) \rightarrow 4\text{H}_2\text{O}(l) + 4e^-$
• Cathode: $\text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4e^- \rightarrow 4\text{OH}^-(aq)$

The overall reaction is $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$, with $E^\circ = 1.23\ \text{V}$. The only product is water. Spacecraft use these cells, and the water is drunk by the crew.

Why use a fuel cell? When hydrogen is burned in an engine, much of the energy is lost as heat. A fuel cell turns the chemical energy straight into electricity. The most it could give is the ratio of $\Delta G^\circ$ to $\Delta H^\circ$, which is $\dfrac{237.1\ \text{kJ/mol}}{285.8\ \text{kJ/mol}}$, or about 83 per cent. Real fuel cells reach 40 to 60 per cent, and a heat engine does about 40 per cent at best.

Corrosion

Corrosion is the slow eating away of a metal by its surroundings. The corrosion of iron is called rusting, and it costs a great deal of money each year.

Iron does not rust in dry air or in water without air. It needs both oxygen and water. Rusting is an electrochemical process, a small galvanic cell on the surface:

• Somewhere on the surface, iron is oxidised. This is the anode: $\text{Fe}(s) \rightarrow \text{Fe}^{2+}(aq) + 2e^-$.
• The electrons run through the iron to another spot. This is the cathode, where oxygen is reduced. In the presence of dissolved carbon dioxide, which makes a little acid, the reaction is $\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}$.
• The $\text{Fe}^{2+}$ ions move through the water film, which acts as the salt bridge.

Add the two for the overall reaction: $2\text{Fe} + \text{O}_2 + 4\text{H}^+ \rightarrow 2\text{Fe}^{2+} + 2\text{H}_2\text{O}$. Its $E^\circ = 1.23 - (-0.44) = 1.67\ \text{V}$, which is large and positive, so it can happen.

Later, the air oxidises $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$, which ends as rust, $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$.

Rusting is faster when salt or acid is present, because they make the water conduct better. That is why a car near the sea rusts quickly.

Stopping corrosion

• A barrier. Paint, oil, grease or plastic keep air and water away from the metal.
• Galvanising. A coat of zinc protects iron even when it is scratched. Zinc has a lower potential than iron, so zinc is oxidised first. This is called sacrificial protection.
• Sacrificial anodes. A block of magnesium or zinc is bolted to a ship's hull or buried beside a pipe and joined to it by a wire. The block corrodes in place of the iron and is replaced from time to time.
• Tin plating. Tin has a higher potential than iron. It protects the iron only while the coat is whole. A scratched tin can rusts faster.
• Alloying. Stainless steel holds chromium and nickel. The chromium forms a thin, tight oxide layer that stops further attack.
• Cathodic protection. A buried iron pipe can be joined to the negative terminal of a source, so the pipe is forced to be the cathode.

Change the current while you watch the line. Does a battery that supplies twice the current last twice as long, or half as long?

Practise with new numbers

Try it in Python

The program follows a lead storage battery as it supplies a steady current.

Common mistakes

• Calling a recharged battery a new reaction. Charging runs the same reaction backwards.
• Saying that the acid in a lead battery is unchanged. It is used up on discharge, and its density falls.
• Thinking that a fuel cell burns hydrogen. The reaction takes place at the electrodes. There is no flame.
• Believing that rust forms where the iron is oxidised. The iron is oxidised at the anode, and rust forms later from the $\text{Fe}^{2+}$ ions.
• Trusting a coat once it is scratched. Only a metal with a lower potential than iron, such as zinc, protects the iron when the coat is broken.

Remember these

• Primary battery: used once. Secondary battery: rechargeable.
• Dry cell: zinc anode, $\text{MnO}_2$ and carbon cathode, about $1.5\ \text{V}$. Mercury cell: steady $1.35\ \text{V}$.
• Lead storage battery: Pb anode, $\text{PbO}_2$ cathode, $\text{H}_2\text{SO}_4$ about $2\ \text{V}$ a cell. Discharge makes $\text{PbSO}_4$ and water.
• Fuel cell: $\text{H}_2$ and $\text{O}_2$ fed in, water out, $E^\circ = 1.23\ \text{V}$.
• Rusting needs oxygen and water. Iron is oxidised at the anode. A metal with a lower potential protects iron.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Why does a mercury cell give a steady voltage for its whole life?

2. What happens to the density of the acid in a lead storage battery as it discharges, and why?

3. What is the difference between a battery and a fuel cell?

4. A ship has a magnesium block bolted to its iron hull. Explain how this protects the hull.