IGCSE Chemistry Paper 4 (Theory / Extended)

Topic 4: Electrochemistry -- Mock Exam 1

Duration1 hour 15 minutes

Total Marks80

Questions7

Time Remaining75:00

Syllabus0620

Instructions

Answer all questions in the spaces provided.
Show all working for calculations.
Use appropriate scientific terminology.
Write balanced half-equations where required.
Your answers will be automatically graded when you submit.

Question Navigation

Question 1 -- Electrolysis of Molten Compounds

Total: 12 marks

A teacher in Mumbai sets up an experiment to electrolyse molten lead(II) bromide (PbBr₂) using carbon electrodes and a d.c. power supply.

(a) [3]

(i) Define electrolysis. [2]

(ii) State why lead(II) bromide must be molten for electrolysis to occur. [1]

Model Answer -- 1(a)

Electrolysis is the decomposition / breakdown of an ionic compound [1]

using an electric current / when molten or in aqueous solution [1]

When molten, the ions are free to move and carry the electric current; when solid, ions are held in fixed positions and cannot move [1]

⚠ If you missed marks here: A definition like 'splitting a compound using electricity' loses marks because it leaves out the key words: the compound is IONIC, and it is decomposed by an electric current when molten or in aqueous solution. For (ii), 'so electrons can flow through it' is wrong – in the melt the current is carried by moving IONS, and solid PbBr₂ contains ions too, so the mark needs the idea that they are only FREE TO MOVE once it is molten.

1 mark -- electrolysis decomposition ionic compound 1 mark -- using electric current molten aqueous solution 1 mark -- molten ions free move carry current solid ions fixed cannot move

(b) [3]

(i) State the name of the electrode connected to the positive terminal of the power supply. [1]

(ii) Describe what you would observe at the cathode during electrolysis. [1]

(iii) Describe what you would observe at the anode during electrolysis. [1]

Model Answer -- 1(b)

The positive electrode is the anode [1]

At the cathode: a silvery / grey metallic bead of lead is deposited [1]

At the anode: brown / orange-brown fumes of bromine gas are produced [1]

⚠ If you missed marks here: 'Lead is formed' and 'bromine is formed' name the products but are not observations – (ii) and (iii) need what you would SEE: a silvery-grey bead of lead at the cathode and brown/orange-brown fumes of bromine at the anode. For (i), the electrode joined to the positive terminal is the anode; the cathode is the negative electrode, which attracts the Pb²⁺ ions.

1 mark -- positive electrode anode 1 mark -- cathode silvery grey metallic lead deposited 1 mark -- anode brown orange fumes bromine gas produced

(i) Write the half-equation for the reaction at the cathode. [2]

(ii) Write the half-equation for the reaction at the anode. [2]

Model Answer -- 1(c)

Cathode: Pb²⁺ + 2e⁻ → Pb [1 for species, 1 for balanced]

Pb²+ + 2e− → Pb

Anode: 2Br⁻ → Br₂ + 2e⁻ [1 for species, 1 for balanced]

2Br− → Br₂ + 2e−

⚠ If you missed marks here: The usual losses are Br⁻ → Br + e⁻ (bromine is diatomic, so it must be 2Br⁻ → Br₂ + 2e⁻) and electrons on the wrong side of the arrow. Electrons are GAINED at the cathode, so they go on the left (Pb²⁺ + 2e⁻ → Pb), and LOST at the anode, so they go on the right; check the charge is the same on both sides of each equation.

1 mark -- cathode lead ion Pb2+ gains electrons 1 mark -- cathode equation balanced 2 electrons 1 mark -- anode bromide ions Br- lose electrons 1 mark -- anode equation balanced 2Br- gives Br2 plus 2e-

(d) [2]

Using the half-equations from part (c):

(i) State which electrode shows reduction, and explain why. [1]

(ii) State which electrode shows oxidation, and explain why. [1]

Model Answer -- 1(d)

Reduction occurs at the cathode because the Pb²⁺ ions gain electrons / there is a gain of electrons [1]

Oxidation occurs at the anode because the Br⁻ ions lose electrons / there is a loss of electrons [1]

⚠ If you missed marks here: Answers about oxygen ('lead loses oxygen') score nothing here, because there is no oxygen in lead(II) bromide – use electrons: Pb²⁺ ions GAIN electrons at the cathode (reduction) and Br⁻ ions LOSE electrons at the anode (oxidation). Swapping the two loses both marks; remember OIL RIG – Oxidation Is Loss, Reduction Is Gain.

1 mark -- reduction cathode gain electrons 1 mark -- oxidation anode lose electrons

Question 2 -- Electrolysis of Aqueous Solutions

Total: 12 marks

A chemistry laboratory in Birmingham investigates the electrolysis of concentrated sodium chloride solution (brine) using inert carbon electrodes.

(a) [2]

State the four types of ions present in concentrated sodium chloride solution. Give the source of each ion.

Model Answer -- 2(a)

Na⁺ and Cl⁻ ions from sodium chloride [1]

H⁺ and OH⁻ ions from the water [1]

⚠ If you missed marks here: Listing only Na⁺ and Cl⁻ loses the second mark – in a solution, WATER also supplies H⁺ and OH⁻ ions, and the question asks for the source of each. Ions written without charges (Na, Cl), or 'H₂O' given as an ion, do not score.

1 mark -- sodium Na+ chloride Cl- ions from sodium chloride 1 mark -- hydrogen H+ hydroxide OH- ions from water

(b) [4]

(i) At the cathode, hydrogen gas is produced rather than sodium metal. Explain why hydrogen is discharged in preference to sodium, using the reactivity series. [2]

(ii) At the anode, chlorine gas is produced rather than oxygen. Explain why chlorine is discharged in preference to oxygen from the concentrated solution. [2]

Model Answer -- 2(b)

Sodium is more reactive than hydrogen / sodium is higher in the reactivity series [1]

The less reactive cation is preferentially discharged at the cathode / hydrogen ions gain electrons more readily than sodium ions [1]

When the solution is concentrated, the halide ion (Cl⁻) is discharged in preference to OH⁻ [1]

Concentration overrides the normal rule: the high concentration of Cl⁻ ions means they are preferentially discharged [1]

⚠ If you missed marks here: A common reversal is 'hydrogen is more reactive than sodium' – sodium is HIGHER in the reactivity series, so its ions stay in solution and the less reactive H⁺ ions are discharged. For (ii), 'chlorine is more reactive than oxygen' scores nothing: chlorine forms because the Cl⁻ ions are at HIGH CONCENTRATION, so they are discharged in preference to OH⁻.

1 mark -- sodium more reactive than hydrogen higher reactivity series 1 mark -- less reactive cation discharged cathode hydrogen ions gain electrons 1 mark -- concentrated halide chloride ion discharged preference hydroxide 1 mark -- concentration overrides high concentration chloride preferentially discharged

(i) Write the half-equation for the reaction at the cathode. [1]

(ii) Write the half-equation for the reaction at the anode. [1]

(iii) Name the solution remaining after electrolysis. [1]

(iv) State one industrial use of the solution named in part (iii). [1]

Model Answer -- 2(c)

Cathode: 2H⁺ + 2e⁻ → H₂ [1]

2H+ + 2e− → H₂

Anode: 2Cl⁻ → Cl₂ + 2e⁻ [1]

2Cl− → Cl₂ + 2e−

The remaining solution is sodium hydroxide (NaOH) [1]

Used in making soap / paper / bleach / ceramics / purifying bauxite [1]

⚠ If you missed marks here: Na⁺ + e⁻ → Na at the cathode contradicts your answer to (b) – hydrogen is discharged, and as H₂ molecules: 2H⁺ + 2e⁻ → H₂, not H⁺ + e⁻ → H. The solution left is sodium hydroxide (the Na⁺ and OH⁻ ions that were NOT discharged), not 'sodium chloride' or 'water'.

1 mark -- cathode 2H+ plus 2e- gives H2 1 mark -- anode 2Cl- gives Cl2 plus 2e- 1 mark -- remaining solution sodium hydroxide NaOH 1 mark -- industrial use soap paper bleach ceramics bauxite

(d) [2]

If a dilute solution of sodium chloride were used instead, state how the products at the anode would differ. Explain your answer.

Model Answer -- 2(d)

Oxygen would be produced at the anode instead of chlorine [1]

In dilute solution, the concentration of Cl⁻ ions is low, so OH⁻ ions are preferentially discharged instead [1]

⚠ If you missed marks here: 'Less chlorine is made' is not enough – with a dilute solution the anode product CHANGES to oxygen, because the Cl⁻ ions are now at low concentration and OH⁻ ions are discharged instead. Do not move hydrogen to the anode: hydrogen still forms at the cathode, and only the anode product changes.

1 mark -- oxygen produced anode instead chlorine 1 mark -- dilute solution low concentration chloride hydroxide ions preferentially discharged

Question 3 -- Extraction of Aluminium

Total: 12 marks

Aluminium is extracted from its ore bauxite at a smelting plant in Jharsuguda, Odisha. The purified ore, aluminium oxide (Al₂O₃), is dissolved in molten cryolite and electrolysed.

(a) [2]

(i) Explain why aluminium cannot be extracted by reduction with carbon. [1]

(ii) State the purpose of cryolite in the extraction process. [1]

Model Answer -- 3(a)

Aluminium is too reactive / too high in the reactivity series to be reduced by carbon [1]

Cryolite lowers the melting point of aluminium oxide from about 2072 °C to about 950 °C, reducing energy costs [1]

⚠ If you missed marks here: For (i), 'carbon is not hot enough' misses the point – aluminium is ABOVE carbon in the reactivity series, so carbon cannot take the oxygen away from aluminium oxide. For (ii), cryolite is NOT a catalyst and it does not lower the melting point of aluminium metal: the aluminium oxide dissolves in molten cryolite, so the electrolyte is molten at about 950 °C instead of 2072 °C, which saves energy.

1 mark -- aluminium too reactive high reactivity series reduced carbon 1 mark -- cryolite lowers melting point reducing energy costs

(b) [3]

(i) State the material used for the electrodes. [1]

(ii) State where the aluminium collects and explain why. [1]

(iii) Identify the gas produced at the anode. [1]

Model Answer -- 3(b)

The electrodes are made of carbon / graphite [1]

Aluminium collects at the bottom of the cell as molten aluminium because it is denser than the electrolyte and sinks [1]

Oxygen gas is produced at the anode [1]

⚠ If you missed marks here: Naming carbon dioxide as the anode gas loses (iii) – the gas discharged at the anode is oxygen (2O²⁻ → O₂ + 4e⁻); CO₂ only forms afterwards, when that oxygen reacts with the hot carbon anode. For (ii), 'at the cathode' is not an explanation: say the molten aluminium sinks to the bottom because it is denser than the electrolyte.

1 mark -- electrodes carbon graphite 1 mark -- aluminium collects bottom cell denser sinks 1 mark -- oxygen gas produced anode

(i) Write the half-equation for the reaction at the cathode. [2]

(ii) Write the half-equation for the reaction at the anode. [2]

Model Answer -- 3(c)

Cathode: Al³⁺ + 3e⁻ → Al [1 for species, 1 for balanced]

Al³+ + 3e− → Al

Anode: 2O²⁻ → O₂ + 4e⁻ [1 for species, 1 for balanced]

2O²− → O₂ + 4e−

⚠ If you missed marks here: Aluminium is in Group III, so its ion is Al³⁺ and the cathode equation needs 3e⁻ – writing Al²⁺ or 2e⁻ loses the mark. At the anode, O²⁻ → O₂ + 2e⁻ is unbalanced: oxygen is diatomic, so it takes TWO oxide ions, and they release 4e⁻ (2O²⁻ → O₂ + 4e⁻).

1 mark -- cathode aluminium ion Al3+ gains electrons 1 mark -- cathode balanced 3 electrons 1 mark -- anode oxide ions O2- lose electrons oxygen formed 1 mark -- anode balanced 2O2- gives O2 plus 4e-

(d) [3]

The carbon anodes must be replaced regularly.

(i) Explain why the anodes need to be replaced. [2]

(ii) Write the equation for the reaction that causes this problem. [1]

Model Answer -- 3(d)

The oxygen produced at the anode reacts with the carbon electrodes at high temperature [1]

This burns away / oxidises the carbon anode, making it smaller over time [1]

C + O₂ → CO₂ [1]

⚠ If you missed marks here: 'The anodes wear out' or 'they dissolve in the cryolite' scores nothing – the marks need the chemistry: the OXYGEN made at the anode reacts with the hot carbon and burns it away as carbon dioxide, C + O₂ → CO₂. That is why it is the anodes, where the oxygen is released, that must be replaced, not the cathode.

1 mark -- oxygen reacts carbon electrodes high temperature 1 mark -- burns away oxidises carbon anode smaller 1 mark -- equation C plus O2 gives CO2

Question 4 -- Purification of Copper by Electrolysis

Total: 12 marks

A copper refinery in Tuticorin, Tamil Nadu uses electrolysis to purify copper for electrical wiring. The electrolyte is copper(II) sulfate solution.

(a) [3]

(i) State what the anode is made of. [1]

(ii) State what the cathode is made of. [1]

(iii) State the electrolyte used. [1]

Model Answer -- 4(a)

The anode is made of impure copper [1]

The cathode is made of pure copper / a thin sheet of pure copper [1]

The electrolyte is copper(II) sulfate solution [1]

⚠ If you missed marks here: Swapping the electrodes is the classic slip: the IMPURE copper is the anode (it dissolves away) and the PURE copper sheet is the cathode (copper builds up on it). Writing just 'copper' for either electrode loses the mark – the words impure and pure are the whole point.

1 mark -- anode impure copper 1 mark -- cathode pure copper thin sheet 1 mark -- electrolyte copper sulfate solution

(b) [4]

(i) Describe what happens to the anode during electrolysis. [1]

(ii) Describe what happens to the cathode during electrolysis. [1]

(iii) Write the half-equation for the reaction at the anode. [1]

(iv) Write the half-equation for the reaction at the cathode. [1]

Model Answer -- 4(b)

The anode dissolves / decreases in mass / gets smaller as copper atoms lose electrons and enter the solution as Cu²⁺ ions [1]

The cathode increases in mass / gets bigger as Cu²⁺ ions from solution gain electrons and are deposited as pure copper [1]

Anode: Cu → Cu²⁺ + 2e⁻ [1]

Cu → Cu²+ + 2e−

Cathode: Cu²⁺ + 2e⁻ → Cu [1]

Cu²+ + 2e− → Cu

⚠ If you missed marks here: If you wrote 4OH⁻ → 2H₂O + O₂ + 4e⁻ or described bubbles at the anode, you used the INERT-electrode rule – with a copper anode the copper atoms themselves lose electrons (Cu → Cu²⁺ + 2e⁻), so the anode dissolves and no oxygen forms. The cathode equation is the exact reverse, Cu²⁺ + 2e⁻ → Cu, so the cathode grows as the anode shrinks.

1 mark -- anode dissolves decreases mass copper atoms lose electrons 1 mark -- cathode increases mass copper ions gain electrons deposited pure 1 mark -- anode equation Cu gives Cu2+ plus 2e- 1 mark -- cathode equation Cu2+ plus 2e- gives Cu

The impure copper contains small amounts of gold and silver.

(i) Explain what happens to the gold and silver during electrolysis. [1]

(ii) Explain why these metals do not dissolve at the anode. [1]

Model Answer -- 4(c)

Gold and silver fall to the bottom of the cell as an anode sludge / anode mud [1]

They are less reactive than copper / lower in the reactivity series, so they do not form ions and dissolve [1]

⚠ If you missed marks here: 'They are heavier, so they sink' gives the wrong reason – gold and silver stay as metal because they are LESS reactive than copper, so they do not lose electrons to form ions; they drop off as the anode dissolves and collect as anode sludge. They are never deposited on the cathode – if they were, the copper there would not be pure.

1 mark -- gold silver fall bottom anode sludge mud 1 mark -- less reactive than copper lower reactivity series do not form ions

(d) [3]

(i) Explain why copper must be purified for use in electrical wiring. [1]

(ii) State one property of copper, other than electrical conductivity, that makes it suitable for wiring. [1]

(iii) Explain why the concentration of the copper(II) sulfate solution remains constant throughout the process. [1]

Model Answer -- 4(d)

Impurities reduce the electrical conductivity of copper / pure copper conducts electricity more efficiently [1]

Copper is ductile / can be drawn into thin wires / is malleable / flexible [1]

For every Cu²⁺ ion removed from solution at the cathode, another Cu²⁺ ion enters the solution from the anode, so the concentration stays the same [1]

⚠ If you missed marks here: In (ii), 'it conducts electricity' is ruled out by the question, and 'it conducts heat' or 'it is shiny' is not why copper makes good wire – the answer is that it is ductile (it can be drawn into thin wires). In (iii) the mark needs the balance: each Cu²⁺ ion removed at the cathode is replaced by one entering the solution from the anode, so the concentration stays the same.

1 mark -- impurities reduce conductivity pure copper conducts efficiently 1 mark -- ductile drawn wires malleable flexible 1 mark -- Cu2+ removed cathode another Cu2+ enters from anode concentration stays same

Question 5 -- Electroplating

Total: 10 marks

A jewellery manufacturer in Jaipur uses electroplating to coat iron bracelets with a thin layer of silver. The electrolyte is silver nitrate solution.

(a) [3]

(i) State two reasons why the manufacturer electroplates the bracelets with silver. [2]

(ii) State what the cathode should be in this electroplating cell. [1]

Model Answer -- 5(a)

To improve appearance / make the bracelet look more attractive [1]

To prevent corrosion / rusting of the iron [1]

The cathode is the iron bracelet / the object to be plated [1]

⚠ If you missed marks here: 'To make it shiny' and 'to make it look attractive' are the SAME reason and score only once – the second mark needs a different one, protecting the iron from rusting (corrosion). The bracelet must be the CATHODE, because the positive Ag⁺ ions are attracted to the negative electrode and are deposited there; as the anode, the iron would dissolve instead.

1 mark -- improve appearance attractive 1 mark -- prevent corrosion rusting iron 1 mark -- cathode iron bracelet object plated

(b) [2]

(i) State what the anode should be made of. [1]

(ii) Explain why this material is chosen for the anode. [1]

Model Answer -- 5(b)

The anode should be made of pure silver [1]

The silver anode dissolves to replace the Ag⁺ ions being deposited from the solution, keeping the concentration of the electrolyte constant [1]

⚠ If you missed marks here: An inert carbon or platinum anode loses the mark – the anode must be pure silver so that it dissolves (Ag → Ag⁺ + e⁻) and replaces the Ag⁺ ions that are plated onto the bracelet. 'Because the coating is silver' is not an explanation for (ii): the point is that the concentration of the electrolyte stays constant.

1 mark -- anode pure silver 1 mark -- silver dissolves replace ions concentration constant

Write the half-equation for the reaction at:

(i) the cathode [1]

(ii) the anode [1]

Model Answer -- 5(c)

Cathode: Ag⁺ + e⁻ → Ag [1]

Ag+ + e− → Ag

Anode: Ag → Ag⁺ + e⁻ [1]

Ag → Ag+ + e−

⚠ If you missed marks here: The silver ion is Ag⁺ (a single positive charge), so each half-equation has ONE electron – Ag²⁺ + 2e⁻ → Ag loses the mark. The two equations are exact reverses: electrons are gained at the cathode (Ag⁺ + e⁻ → Ag) and lost at the anode (Ag → Ag⁺ + e⁻).

1 mark -- cathode Ag+ plus e- gives Ag 1 mark -- anode Ag gives Ag+ plus e-

(d) [3]

(i) State one factor that affects the thickness of the silver coating. [1]

(ii) Explain why the electrolyte must contain ions of the plating metal. [1]

(iii) Suggest why the manufacturer uses a low current during electroplating rather than a high current. [1]

Model Answer -- 5(d)

Time / duration of electrolysis / current size / concentration of electrolyte [1]

The metal ions in the electrolyte are the source of the metal atoms that are deposited on the cathode / the ions gain electrons to form the metal coating [1]

A low current produces a smoother, more even coating; a high current can produce a rough, uneven or powdery deposit [1]

⚠ If you missed marks here: In (ii), 'so the solution conducts' is not enough – any ionic solution conducts; ions of the PLATING metal are needed because they are what gain electrons at the cathode and become the coating. In (iii), 'a low current is safer' or 'cheaper' misses the mark: a high current deposits the silver too fast and gives a rough, uneven or powdery coating.

1 mark -- time duration current concentration affects thickness 1 mark -- metal ions source deposited cathode gain electrons form coating 1 mark -- low current smoother even coating high current rough uneven powdery

Question 6 -- Electrochemical Cells

Total: 12 marks

A student in London sets up a simple electrochemical cell using a strip of zinc and a strip of copper dipping into dilute sulfuric acid. A voltmeter is connected between the two metals.

(a) [3]

(i) State what an electrochemical cell converts. [1]

(ii) Explain why a voltage is produced when two different metals are used. [2]

Model Answer -- 6(a)

An electrochemical cell converts chemical energy into electrical energy [1]

The two metals have different reactivities / different tendencies to lose electrons [1]

This difference in reactivity creates a potential difference / voltage between the electrodes, causing electrons to flow through the external circuit [1]

1 mark -- converts chemical energy electrical energy 1 mark -- different metals different reactivities tendencies lose electrons 1 mark -- difference reactivity creates potential difference voltage electrons flow external circuit

(b) [4]

(i) State the direction of electron flow in the external circuit. [1]

(ii) Explain which metal acts as the negative electrode. [1]

(iii) Write the half-equation for the reaction at the zinc electrode. [1]

(iv) Write the half-equation for the reaction at the copper electrode. [1]

Model Answer -- 6(b)

Electrons flow from zinc to copper through the external circuit [1]

Zinc is the negative electrode because it is more reactive / higher in the reactivity series, so it loses electrons more readily [1]

Zinc: Zn → Zn²⁺ + 2e⁻ [1]

Zn → Zn²+ + 2e−

Copper: 2H⁺ + 2e⁻ → H₂ [1]

2H+ + 2e− → H₂

1 mark -- electrons flow zinc copper external circuit 1 mark -- zinc negative electrode more reactive higher reactivity series 1 mark -- zinc equation Zn gives Zn2+ plus 2e- 1 mark -- copper equation 2H+ plus 2e- gives H2

(i) Predict how the voltage would change if the zinc were replaced by magnesium. Explain your answer. [2]

(ii) Predict the effect on the voltage if both electrodes were made of copper. [1]

Model Answer -- 6(c)

The voltage would increase [1]

Magnesium is more reactive than zinc / further apart in reactivity series from copper, so there is a greater difference in reactivity, producing a higher voltage [1]

No voltage / zero reading, because both electrodes are the same metal so there is no difference in reactivity [1]

1 mark -- voltage increase 1 mark -- magnesium more reactive greater difference reactivity higher voltage 1 mark -- no voltage zero same metal no difference reactivity

(d) [2]

State what you would observe happening to:

(i) the zinc strip over time [1]

(ii) the surface of the copper strip [1]

Model Answer -- 6(d)

The zinc strip dissolves / gets smaller / decreases in size as zinc atoms form Zn²⁺ ions [1]

Bubbles of gas (hydrogen) are produced on the surface of the copper strip [1]

1 mark -- zinc dissolves gets smaller decreases size 1 mark -- bubbles gas hydrogen produced surface copper

Question 7 -- Hydrogen Fuel Cells

Total: 10 marks

The Delhi Metro Rail Corporation is evaluating hydrogen fuel cell buses for use on feeder routes. A fuel cell uses hydrogen and oxygen to generate electricity.

(a) [3]

(i) State what energy conversion takes place in a hydrogen fuel cell. [1]

(ii) Write the overall equation for the reaction in a hydrogen fuel cell. [1]

(iii) State the only product formed. [1]

Model Answer -- 7(a)

Chemical energy is converted into electrical energy [1]

2H₂ + O₂ → 2H₂O [1]

The only product is water [1]

⚠ If you missed marks here: 'Electrical to chemical' is the electrolysis direction – a fuel cell converts CHEMICAL energy into ELECTRICAL energy. H₂ + O₂ → H₂O is unbalanced (2 O atoms on the left, 1 on the right), so it must be 2H₂ + O₂ → 2H₂O, and water is the ONLY product – there is no carbon in the fuel, so no CO₂.

1 mark -- chemical energy converted electrical energy 1 mark -- equation 2H2 plus O2 gives 2H2O 1 mark -- only product water

(b) [3]

State three advantages of using hydrogen fuel cells for buses compared to using petrol or diesel engines.

Model Answer -- 7(b)

No carbon dioxide is produced / does not contribute to global warming / climate change [1]

No sulfur dioxide or nitrogen oxides produced / does not cause acid rain or air pollution [1]

The only waste product is water, which is not a pollutant / no particulates produced [1]

⚠ If you missed marks here: 'It is clean' or 'it is environmentally friendly' scores nothing on its own – each mark needs a named pollutant that the bus avoids. 'No CO₂', 'no greenhouse gases' and 'no global warming' are ONE point, so make your three different: no CO₂ (climate change), no sulfur dioxide or nitrogen oxides (acid rain), and only water made (no particulates).

1 mark -- no carbon dioxide global warming climate change 1 mark -- no sulfur dioxide nitrogen oxides acid rain pollution 1 mark -- only waste water not pollutant no particulates

State two disadvantages or challenges of using hydrogen fuel cells for transport.

Model Answer -- 7(c)

Hydrogen is difficult to store / must be stored under high pressure or as a liquid at very low temperatures [1]

Hydrogen is highly flammable / explosive, posing safety risks / lack of hydrogen refuelling infrastructure / hydrogen is expensive to produce [1]

⚠ If you missed marks here: Two storage points ('hard to store' and 'needs high-pressure tanks') are the SAME challenge and score once, and 'the water it makes is a pollutant' is simply wrong. Your second point must be a different one: hydrogen is highly flammable or explosive, there are few refuelling stations, or it is expensive to produce.

1 mark -- hydrogen difficult store high pressure low temperature 1 mark -- flammable explosive safety infrastructure expensive produce

(d) [2]

State two differences between a hydrogen fuel cell and a conventional rechargeable battery.

Model Answer -- 7(d)

A fuel cell uses an external supply of fuel (hydrogen) which is continuously fed in, whereas a battery stores its chemicals internally [1]

A fuel cell produces electricity as long as fuel is supplied, whereas a battery runs out / needs recharging [1]

⚠ If you missed marks here: 'A fuel cell is cleaner than a battery' does not answer the question – the marks are for how they get their chemicals. A fuel cell is fed hydrogen continuously from outside and works for as long as fuel is supplied, whereas a battery holds a fixed store of chemicals inside and runs flat until it is recharged.

1 mark -- fuel cell external supply fuel continuously fed battery stores chemicals internally 1 mark -- fuel cell produces electricity long fuel supplied battery runs out recharging

Self-Assessment

Tick marks earned, then click Calculate Grade.

Your Score

Out Of

Percentage