Electrolysis and Fuel Cells: Key Concepts and Reactions

Electrolysis

The decomposition (breaking down) of an ionic substance using electricity.

Direct current (d.c.) power supply

A power supply that provides a constant flow of electricity in one direction.

Electrodes

Conductors through which electricity enters and leaves.

Electrolyte

Ionic substance in molten or aqueous form, where ions are free to move.

Cations

Positive ions that move to the negative electrode (cathode) during electrolysis.

Anions

Negative ions that move to the positive electrode (anode) during electrolysis.

OILRIG

Oxidation Is Loss, Reduction Is Gain (of electrons).

Cathode

The electrode where reduction occurs.

Anode

The electrode where oxidation occurs.

Electrolysis of molten lead bromide (PbBr₂)

At cathode (-): Pb²⁺ + 2e⁻ → Pb; At anode (+): 2Br⁻ → Br₂ (gas) + 2e⁻.

Metallic lead and bromine gas

The products formed from the electrolysis of molten lead bromide.

General rule for molten compounds

Metal forms at cathode; Non-metal forms at anode.

Reactivity series

A list that ranks metals by their reactivity, determining extraction methods.

Aluminium oxide (Al₂O₃)

The ore from which aluminium is extracted, found in bauxite.

Cryolite (Na₃AlF₆)

A substance used to dissolve aluminium oxide, lowering its melting point.

Electrolysis cell

A steel tank lined with graphite acting as the cathode, with graphite anodes inserted.

Costs of electrolysis

Expensive due to large electricity requirements and the need to replace anodes.

Hydrogen gas production

Occurs at the cathode if the metal is more reactive than hydrogen.

Aqueous solutions

Contain both the compound's ions and ions from water (H⁺ and OH⁻).

Electrolysis of aqueous solutions

More complex due to the presence of additional ions from water.

Anode (+)

If halide ions are present (Cl⁻, Br⁻, I⁻), the halogen is released. Otherwise, oxygen is released (from OH⁻ ions).

Cathode reaction

2H⁺ + 2e⁻ → H₂ (hydrogen gas).

Anode reaction

2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas).

Left in solution

Sodium hydroxide (NaOH).

Products of electrolysis

Hydrogen gas, chlorine gas, sodium hydroxide solution.

Examiner Tip for electrolysis

Always list all possible ions in aqueous electrolysis before deciding which products form.

Aim of electrolysis practical

Investigate products formed in aqueous electrolysis using graphite electrodes.

Electrolysis method

Set up electrolysis cell with beaker, electrodes, d.c. power supply. Insert inert electrodes (graphite). Pass current through solution. Collect and test gases produced.

Hydrogen test

Lit splint makes a squeaky pop.

Oxygen test

Relights a glowing splint.

Chlorine test

Bleaches damp blue litmus paper.

Example of electrolysis

CuCl₂ → Copper at cathode, chlorine at anode.

Example of NaCl(aq) electrolysis

Hydrogen, chlorine, NaOH left.

Balanced half equations

Always write balanced half equations for each electrode.

Half equations in electrolysis

Half equations show electron transfer at one electrode.

Cathode half equations

Cu²⁺ + 2e⁻ → Cu and 2H⁺ + 2e⁻ → H₂.

Anode half equations

2Cl⁻ → Cl₂ + 2e⁻ and 4OH⁻ → O₂ + 2H₂O + 4e⁻.

Examiner Tip for balancing

Balance atoms first, then add electrons to balance charges. Check total charge on both sides.

Cell definition

A cell produces electricity from a chemical reaction between two electrodes in an electrolyte.

Voltage factors

Voltage depends on the metals used, the difference in reactivity between them, and the electrolyte.

Batteries

Two or more cells joined in series → greater voltage.

Rechargeable batteries

Reaction is reversible → electricity recharges them.

Non-rechargeable batteries

Reaction is irreversible. Once reactants are used up, battery goes flat and must be replaced.

Fuel cell definition

Device that produces electricity from a continuous supply of fuel and oxygen.

Hydrogen fuel cell reaction

2H₂ + O₂ → 2H₂O.

Advantages of fuel cells

Only product = water (non-polluting). Continuous electricity as long as fuel supplied. Lightweight and compact compared to batteries.

Disadvantages of fuel cells

Hydrogen difficult to store safely. Hydrogen often produced from fossil fuels → CO₂ emissions. Expensive technology.

Electrode reactions in hydrogen fuel cells

At anode (-): Hydrogen oxidised. 2H₂ → 4H⁺ + 4e⁻. At cathode (+): Oxygen reduced. O₂ + 4H⁺ + 4e⁻ → 2H₂O.

Overall reaction of hydrogen fuel cells

2H₂ + O₂ → 2H₂O.

Examiner Tip for fuel cells

Learn both half equations AND overall equation. Marks are often lost if students only write the overall equation.

Explain why electrolysis is used to extract aluminium, not reduction with carbon.

Aluminium is more reactive than carbon, so it cannot be displaced or reduced by carbon. Therefore, electrolysis is required to break down aluminium oxide into aluminium and oxygen.

Why is cryolite used in the extraction of aluminium?

Cryolite is used because it lowers the melting point of aluminium oxide. This reduces the energy required and makes the process cheaper, as less electricity is used.

Why must the positive electrodes (anodes) be replaced regularly in aluminium extraction?

The oxygen produced at the anode reacts with the carbon (graphite) electrodes to form carbon dioxide. This causes the anodes to gradually wear away and they must be replaced.

Predict the products of the electrolysis of aqueous sodium chloride solution (brine).

At the cathode, hydrogen gas is produced because sodium is more reactive than hydrogen.At the anode, chlorine gas is produced because chloride ions are present.The solution left behind is sodium hydroxide.

Write the half equations for the electrolysis of sodium chloride solution.

At the cathode: 2H⁺ + 2e⁻ → H₂At the anode: 2Cl⁻ → Cl₂ + 2e⁻

How could you test the gases produced in the electrolysis of aqueous sodium chloride?

Test for hydrogen: insert a lit splint into the gas → produces a squeaky pop.

Test for chlorine: damp blue litmus paper → turns red then bleaches.

Explain why ions can conduct electricity when molten or dissolved, but not when solid.

In solid ionic compounds, ions are fixed in a lattice and cannot move.When molten or in solution, the ions are free to move and carry charge, allowing electricity to pass.

What happens to the lead ions (Pb²⁺) during the electrolysis of molten lead bromide?

Pb²⁺ ions move to the cathode where they gain two electrons (reduction) to form lead atoms.Equation: Pb²⁺ + 2e⁻ → Pb

What is a cell and how is a battery different?

A cell produces electricity from a chemical reaction between two different electrodes in an electrolyte.A battery consists of two or more cells connected in series to produce a higher potential difference.

Why does the potential difference of a cell depend on the metals used?

The voltage depends on the difference in reactivity of the two metals. The greater the difference in reactivity, the greater the potential difference produced.

What is the main difference between a rechargeable battery and a non-rechargeable battery?

In a rechargeable battery, the chemical reactions can be reversed by passing an electric current through the cell.In a non-rechargeable battery, the reactions are irreversible, so once reactants are used up the battery goes flat and cannot be reused.

Give two advantages of hydrogen fuel cells compared to rechargeable batteries.

Only product is water, so no polluting emissions.

Will produce electricity continuously as long as hydrogen and oxygen are supplied.

Hydrogen fuel cells are lightweight and compact.

Give two disadvantages of hydrogen fuel cells.

Hydrogen is difficult to store safely as it is explosive.

Hydrogen is often produced from fossil fuels, releasing carbon dioxide.

Expensive to manufacture fuel cells.

Write the half equations for the hydrogen fuel cell.

At the anode: 2H₂ → 4H⁺ + 4e⁻At the cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂OOverall: 2H₂ + O₂ → 2H₂O

Compare the use of hydrogen fuel cells and rechargeable batteries in vehicles.

Hydrogen fuel cells produce only water, so no pollution, but hydrogen must be stored at high pressure and is difficult to transport.

Rechargeable batteries can be recharged and are convenient, but they have a limited lifetime and need replacing.

Hydrogen fuel cells can run continuously if fuel is supplied, while batteries need regular recharging.

Explain why aluminium is extracted using electrolysis and describe the process. (6 marks)

Aluminium is more reactive than carbon, so it cannot be reduced by carbon.

Aluminium oxide is mixed with cryolite to lower its melting point and reduce energy costs.

Electrolysis splits the aluminium oxide into ions: Al³⁺ and O²⁻.

At the cathode, Al³⁺ ions gain 3 electrons (reduction) to form aluminium metal.

Equation: Al³⁺ + 3e⁻ → Al

At the anode, O²⁻ ions lose 2 electrons (oxidation) to form oxygen gas.

Equation: 2O²⁻ → O₂ + 4e⁻

Oxygen reacts with the carbon anodes, producing CO₂, so anodes must be replaced regularly.

The electrolysis of sodium chloride solution (brine) produces useful products. Explain what these products are and how they are formed. (6 marks)

Brine contains Na⁺, Cl⁻, H⁺ and OH⁻ ions.

At the cathode, hydrogen is produced because hydrogen is less reactive than sodium.

2H⁺ + 2e⁻ → H₂

At the anode, chlorine gas is formed because chloride ions are discharged.

2Cl⁻ → Cl₂ + 2e⁻

The solution left behind contains sodium hydroxide (NaOH).

The products are useful: hydrogen (fuel), chlorine (bleach/cleaning), and sodium hydroxide (soap production).

Explain why ionic compounds can only conduct electricity when molten or in solution. (4 marks)

In the solid state, ions are held in a fixed lattice and cannot move.

When molten or dissolved, the ionic lattice breaks apart.

Ions are free to move and carry an electric current.

Therefore, ionic compounds only conduct when molten or in aqueous solution

Compare the advantages and disadvantages of hydrogen fuel cells and rechargeable batteries for powering cars. (6 marks)

Fuel cells advantages:

Only product is water, so no pollution.

Will run continuously if hydrogen and oxygen are supplied.

Lightweight and compact compared to some batteries.

Fuel cells disadvantages:

Hydrogen is flammable and difficult to store safely.

Hydrogen often produced from fossil fuels, releasing CO₂.

Expensive to manufacture.

Rechargeable batteries advantages:

Can be recharged many times.

Convenient charging infrastructure developing (electric charging stations).

Safer to store than hydrogen.

Rechargeable batteries disadvantages:

Have a limited lifetime, eventually need replacing.

Disposal can cause environmental issues (toxic metals).

Need frequent recharging.

Describe how a hydrogen fuel cell produces electricity. (6 marks)

Hydrogen is supplied to the anode, where hydrogen molecules are split into H⁺ ions and electrons.

2H₂ → 4H⁺ + 4e⁻

Electrons flow through an external circuit, producing an electric current.

Hydrogen ions (H⁺) pass through the electrolyte to the cathode.

At the cathode, oxygen reacts with hydrogen ions and electrons to form water.

O₂ + 4H⁺ + 4e⁻ → 2H₂O

Overall reaction: 2H₂ + O₂ → 2H₂O

This process continues as long as hydrogen and oxygen are supplied.

Explain how the voltage of a chemical cell depends on the metals used. (4 marks)

hVoltage depends on the difference in reactivity between the two electrodes.

A larger reactivity difference gives a higher potential difference (voltage).

Different metals in the electrochemical series will produce different voltages.

The electrolyte also affects the size of the voltage.