Simple electrochemical processes

A cell is a device that transforms chemical energy into electrical energy thanks to a chemical reaction taking place inside it. Explain in your own words what this phrase means and name two devices in your home that run on batteries.

At the science fair, a group inserts a zinc plate and a copper plate into a lemon, connects them to an LED, and it lights up: they have built a "lemon battery", as shown in the diagram.

1. a) What acts as the electrolyte (the medium that conducts the current inside the cell) in this setup?
2. b) Why do you think two different metals are needed and two identical plates would not work?

Electrolysis is the reverse process of a cell: an electric current forces a chemical reaction to occur that would not happen on its own. Complete the summary with the missing words: "In a cell, a chemical reaction $\dots$ electric current; in electrolysis, electric current $\dots$ a chemical reaction".

When direct current is passed through water (with a few drops of acid so that it conducts electricity), bubbles of two different gases appear at the electrodes, as shown in the diagram.

1. a) Write the balanced equation for the decomposition of water.
2. b) Name the two gases given off.

To silver-plate a spoon, the spoon is connected to the negative terminal of a power supply and a silver sheet to the positive terminal, both immersed in a silver nitrate solution ($AgNO_3$), as shown in the diagram.

1. a) On which of the two electrodes is the silver deposited?
2. b) The solution would gradually lose silver ions were it not for the fact that the sheet dissolves: explain what role the sheet connected to the positive terminal plays.
3. c) Reason what would happen to the spoon if it were accidentally connected to the positive terminal.

In the Daniell cell, these two half-reactions occur: at the zinc electrode, $Zn \rightarrow Zn^{2+} + 2\ e^-$; at the copper electrode, $Cu^{2+} + 2\ e^- \rightarrow Cu$.

1. a) Add the two half-reactions together and write the overall reaction of the cell.
2. b) Which electrode acts as the negative terminal, that is, the one that donates electrons to the circuit?
3. c) Explain, using the half-reactions, why the zinc electrode loses mass and the copper electrode gains mass while the cell is operating.

Classify these cells and batteries as "single-use" or "rechargeable": the alkaline battery of a remote control, the lithium-ion battery of a mobile phone, the lead battery of a car and the button cell of a watch. Explain in your own words what it means for a battery to be rechargeable.

When you plug your mobile phone into the charger, the chemical reaction in its battery occurs "in reverse". Explain, using the concept of a reverse process, what the charger does with the electric current and how recharging a battery is similar to electrolysis.

Compare a cell and electrolysis by answering these three questions for each process:

1. a) Which process produces electric current and which consumes it?
2. b) In which is the chemical reaction spontaneous and in which is it forced by the current?
3. c) What energy conversion occurs in each?

In addition, give a real-life example of each process.

In the electrolysis of water, $2\ H_2O \rightarrow 2\ H_2 + O_2$, the volume of hydrogen collected is always twice that of oxygen, as shown in the diagram.

Explain, using the coefficients of the balanced equation, why this happens.

In the laboratory, water is electrolysed, collecting the gases in two inverted graduated tubes placed over the electrodes, as shown in the diagram.

When the hydrogen tube contains $20\ \text{mL}$, what volume of oxygen do you expect to find in the other tube, measured under the same conditions? Explain the ratio using the coefficients of the balanced equation.

In an electrolytic cell, water is decomposed and, between the two electrodes, a total of $72\ \text{mL}$ of gas is collected, measured under the same conditions.

1. a) What volume of hydrogen and what volume of oxygen have been collected?
2. b) Under these conditions, $1\ \text{mol}$ of any gas occupies $24\ \text{L}$: calculate the number of moles of each gas.
3. c) Using the molar mass of water, $18\ \text{g/mol}$, calculate what mass of water has decomposed.

Aluminium is extracted by the electrolysis of its molten oxide (alumina), a process that consumes huge amounts of electricity.

1. a) Explain, using the concept of a "reaction forced by current", why producing aluminium is so energy-intensive.
2. b) Recycling a can consumes only $5\ \%$ of the energy required to produce new aluminium. For every $100\ \text{kWh}$ it would cost to produce primary aluminium, calculate how much energy is saved if that aluminium is obtained by recycling.
3. c) Using this data, argue why recycling cans is an important measure from both a chemistry and an environmental perspective.

The steel hull of ships has zinc blocks welded to it called "sacrificial anodes": it is the zinc that corrodes instead of the iron, as shown in the diagram.

1. a) Taking into account what happens in the Daniell cell, where zinc wears away by losing electrons more easily than copper, explain why a zinc block protects the iron of the hull.
2. b) Reason what would happen if the blocks were made of copper instead of zinc.
3. c) Why must these blocks be inspected and replaced periodically?