2.3Chemical Bonding, Application Of Chemical Reactions and Organic Chemistry
Metals and Their Extraction — WJEC GCSE Chemistry Revision Notes & Topical Questions
20 min read 34 parts Lesson 3 of 7
Lesson summary
Complete PDF with figures and original page layoutPage 1 WJEC GCSE Chemistry Your notes 2.3 Metals & Their Extraction Contents Metal Ores Native Metals Metal Reactivity Metal Displacement Reactions Oxidation & Reduction Extraction of Iron …
Section 1 of 34Metal ores
- Crust
- the Earth’s crust contains metals and metal compounds such as gold, copper, iron oxide and aluminium oxide
- Abundance
- aluminium is the most abundant metal in the crust and, being reactive, forms compounds
- Ore = a rock containing enough metal to make extraction worthwhile
- Extraction methods
- electrolysis, a blast furnace, or reaction with a more reactive material
- Reduction
- many ores are oxides, so extraction removes oxygen and is a reduction
| Ore | Metal compound | Formula |
|---|---|---|
| Haematite | Iron oxide | Fe2O3 |
| Bauxite | Aluminium oxide | Al2O3 |
| Rock salt | Sodium chloride | NaCl |
| Rutile | Titanium dioxide | TiO2 |
Section 2 of 34Native metals and extraction method
- Native metals
- unreactive metals found as the uncombined element because they do not easily react, such as gold and silver
- Rule
- the extraction method depends on the metal’s position in the reactivity series
| Metals | Method of extraction |
|---|---|
| Potassium, sodium, lithium, calcium, magnesium, aluminium (above carbon) | Electrolysis of the molten chloride or oxide, which uses large amounts of electricity and is expensive |
| Zinc, iron (below carbon) | Heating with a reducing agent such as carbon or carbon monoxide |
| Copper | Often heated with carbon or hydrogen to extract from its ore, or found as the pure element |
| Silver, gold | Found as pure elements |
Section 3 of 34The reactivity series
- Basis
- metals are placed in order of reactivity based on their reactions with other substances
- Meaning
- reactive metals lose electrons easily to form positive ions
- Position
- reactivity decreases going down the series, and a metal higher in the series is more reactive
- Carbon and hydrogen
- these non-metals are included because they reduce metal oxides
Section 4 of 34Displacement reactions with metal oxides
- Rule
- a more reactive metal displaces a less reactive metal from its compounds, and can be done with a metal oxide on heating
- Reducing agent
- zinc reduces copper(II) oxide, Zn + CuO → ZnO + Cu, so zinc is the reducing agent
| Mixture | Result | Equation |
|---|---|---|
| Iron(III) oxide and aluminium | Iron and aluminium oxide | Fe2O3 + 2Al → 2Fe + Al2O3 |
| Sodium oxide and magnesium | No reaction, sodium is more reactive | — |
| Silver oxide and copper | Silver and copper(II) oxide | Ag2O + Cu → 2Ag + CuO |
| Zinc oxide and calcium | Zinc and calcium oxide | ZnO + Ca → Zn + CaO |
| Lead(II) oxide and silver | No reaction, lead is more reactive | — |
- Thermite reaction
- iron(III) oxide with aluminium, Fe2O3 + 2Al → 2Fe + Al2O3
- Aluminium is more reactive, so it displaces the iron.
- It is so exothermic that the iron is molten, so it is used in welding and incendiary devices.
Section 5 of 34Displacement in salt solutions
- Principle
- the reactivity of two metals is compared using the solution of one metal’s salt
- Observation
- the more reactive metal slowly disappears and the less reactive metal is displaced
- Magnesium and copper(II) sulfate
- Mg + CuSO4 → MgSO4 + Cu
- The blue colour fades to colourless magnesium sulfate.
- Copper coats the magnesium and also falls to the bottom as a solid.
| Mixture | Result | Equation |
|---|---|---|
| Magnesium and iron(II) sulfate | Magnesium sulfate and iron | Mg + FeSO4 → MgSO4 + Fe |
| Zinc and sodium chloride | No reaction | — |
| Lead and silver nitrate | Lead(II) nitrate and silver | Pb + 2AgNO3 → Pb(NO3)2 + 2Ag |
| Copper and calcium chloride | No reaction | — |
| Iron and copper(II) sulfate | Iron(II) sulfate and copper | Fe + CuSO4 → FeSO4 + Cu |
- Syllabus note
- the approximate position of common metals, carbon and hydrogen should be known, but the series need not be recalled in full
Section 6 of 34Practical: metal displacement reactions
- Aim
- to determine the order of reactivity of zinc, magnesium, copper and iron
- Apparatus
- 1 cm strips of the four metals, 5 cm3 of 0.1 mol dm-3 zinc, magnesium, copper(II) and iron(II) sulfates, four dropping pipettes and a dimple tile
- Add one to two drops of zinc sulfate to the first row of dimples.
- Add one piece of each metal to different dimples.
- Record any changes in a table, with a tick for reaction and a cross for no reaction.
- Expect no change where the metal matches the compound.
- Use a different pipette for each solution to avoid contamination.
- Repeat for the other sulfate solutions.
- Practical tip
- do not overfill the dimples, or solutions overflow and contaminate other reactions
Section 7 of 34Practical: results and conclusion
| Zinc | Magnesium | Copper | Iron | |
|---|---|---|---|---|
| Zinc sulfate | — | Reaction | No reaction | No reaction |
| Magnesium sulfate | No reaction | — | No reaction | No reaction |
| Copper(II) sulfate | Reaction | Reaction | — | Reaction |
| Iron(II) sulfate | Reaction | Reaction | No reaction | — |
- Magnesium
- it displaces all the other metals from solution
- Zinc and iron
- zinc displaces iron and copper, and iron displaces only copper
- Copper
- it displaces none of the metals
- Order of reactivity
- magnesium (most reactive), zinc, iron, copper (least reactive), because the metal with most reactions is most reactive
- Equations
- Mg + ZnSO4 → Zn + MgSO4, Mg + CuSO4 → Cu + MgSO4, and Mg + FeSO4 → Fe + MgSO4
Section 8 of 34Oxidation and reduction in terms of oxygen
- Redox
- oxidation and reduction occur together in the same reaction
- Oxidation = a reaction in which oxygen is added to an element or compound
- Reduction = a reaction in which oxygen is removed from an element or compound
- Zinc oxide and carbon
- zinc oxide loses oxygen so it is reduced, and carbon gains oxygen so it is oxidised
- Blast furnace
- iron(III) oxide reacts with carbon monoxide in a redox reaction where the iron oxide is reduced and the carbon monoxide is oxidised
Section 9 of 34Blast furnace: overview and raw materials
- Why carbon works
- carbon is above iron in the reactivity series, so it reduces iron(III) oxide, and it is the iron(III) oxide that is reduced, not iron
- Scale
- modern blast furnaces make about 10,000 tonnes of iron per day
- Continuous process
- raw materials are added and products removed all the time because of the time and cost of reaching temperature
- Input
- iron ore, coke and limestone are added at the top, and hot air is blown in at the bottom
| Raw material | Formula | Use |
|---|---|---|
| Iron ore (haematite) | Fe2O3 | Source of iron |
| Coke | C | Provides carbon |
| Limestone | CaCO3 | Removes acidic impurities |
Section 10 of 34Blast furnace: reactions in the three zones
- Zone 1: coke burns in hot air, C(s) + O2(g) → CO2(g), which is exothermic and heats the furnace.
- Zone 2: at high temperature more coke reduces carbon dioxide, CO2(g) + C(s) → 2CO(g).
- Zone 3: carbon monoxide reduces iron(III) oxide, Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g), and molten iron collects at the bottom and is tapped off.
- Redox in the furnace
- iron(III) oxide is reduced and carbon monoxide is oxidised, since it gains oxygen to form carbon dioxide
Section 11 of 34Blast furnace: removing impurities
- Limestone decomposes
- CaCO3(s) → CaO(s) + CO2(g)
- Neutralisation
- calcium oxide reacts with the acidic impurity silicon dioxide, CaO(s) + SiO2(s) → CaSiO3(l)
- Slag
- the calcium silicate melts and floats as molten slag on the iron, and is tapped off separately
Section 12 of 34Electrolysis of molten ionic compounds
- Electrolysis
- an electric current decomposes a molten ionic compound, or an aqueous solution of one
- Electrolyte
- an ionic liquid or solution that conducts electricity, and covalent compounds cannot be electrolysed
- Terms
- Electrodea rod of metal or graphite through which current enters or leaves.
- Anodethe positive electrode, which attracts anions.
- Cathodethe negative electrode, which attracts cations.
- Electrolytic cellthe whole set-up.
- Why molten
- the ions are free to move and carry charge only when the compound is molten or dissolved
- Predicting products
- for a binary molten compound, the cathode product is the metal and the anode product is the non-metal
- Inert electrodes
- graphite or platinum are used so they do not take part in side reactions
Section 13 of 34Electrolysis of molten lead(II) bromide
- Heat lead(II) bromide in a crucible until molten.
- Add two graphite electrodes connected to a power pack.
- Switch on the current.
- Bromide ions move to the anode and lose electrons, giving brown bromine gas.
- Lead ions move to the cathode and gain electrons, giving grey lead metal.
- Anode half-equation (oxidation)
- 2Br- → Br2 + 2e-
- Cathode half-equation (reduction)
- Pb2+ + 2e- → Pb
Section 14 of 34Oxidation and reduction in terms of electrons
- Oxidation
- a substance loses electrons
- Reduction
- a substance gains electrons
- Memory aids
- OIL RIG, and RED CAT / AN OX (reduction at the cathode, oxidation at the anode)
- (oxidation is loss, reduction is gain)
- OIL RIG, and RED CAT / AN OX (reduction at the cathode, oxidation at the anode)
- Iron and copper sulfate
- Fe + CuSO4 → FeSO4 + Cu
- Ionic equation without spectator ionsFe + Cu2+ → Fe2+ + Cu.
- Iron loses electrons and is oxidised, and Cu2+ gains electrons and is reduced.
- In electrolysis
- positive ions are reduced at the cathode (Pb2+ + 2e- → Pb) and negative ions are oxidised at the anode (2Br- → Br2 + 2e-)
- Identifying changes
- in V3+ + Fe3+ → V4+ + Fe2+, V3+ loses an electron so it is oxidised, and Fe3+ gains one so it is reduced
Section 15 of 34Extraction of aluminium: method
- Why electrolysis
- aluminium is above carbon in the reactivity series, so carbon cannot reduce its oxide
- Ore
- bauxite, purified to aluminium oxide, Al2O3
- Cryolite
- aluminium oxide is dissolved in molten cryolite because pure aluminium oxide melts above 2000 °C,
- which would use a lot of energy, and the mixture melts lower without interfering
- Cell
- a steel cell lined with graphite, which is the negative electrode, with large graphite blocks as positive electrodes
- Cost
- a large amount of electricity is needed, which is a major expense
Section 16 of 34Extraction of aluminium: electrode reactions
1Cathode (reduction)3 points
- aluminium ions gain electrons
- Al3+ + 3e- → Al.
- Molten aluminium collects at the bottom and is siphoned off from time to time, and fresh aluminium oxide is added.
2Anode (oxidation)3 points
- oxide ions lose electrons
- 2O2- → O2 + 4e-.
- The graphite anode reacts with the oxygen, C(s) + O2(g) → CO2(g), so it wears away and must be replaced regularly.
Section 17 of 34Properties and uses of metals
- Principle
- different properties decide the use, such as aluminium’s low density, conductivity and strength making it suitable for overhead power cables
| Metal | Properties | Uses |
|---|---|---|
| Steel | Hard and strong | Buildings, reinforced concrete, cars, tools |
| Aluminium | Strong, low density, good conductor of heat and electricity, resistant to corrosion | Aeroplane parts, cans, window frames, foil, overhead cables |
| Copper | Very good conductor, malleable and ductile, attractive colour and lustre | Wiring, pipes, motors |
| Titanium | Hard, strong, low density, corrosion-resistant, high melting point | Surgical instruments, joint replacements, spacecraft, sports equipment |
- Reinforced concrete
- steel rods increase the strength of concrete in buildings
- Exam skill
- explain why a substance suits its use, such as copper conducting electricity in wiring
Section 18 of 34Transition metals
- Position
- found in the centre of the Periodic Table
- Typical metal properties
- high melting and boiling points, malleable, high density, and conduct heat and electricity
- Extra properties
- they form coloured compounds, can form more than one type of ion (such as Fe2+ and Fe3+) and are useful as catalysts
- Catalyst uses
- Iron in the Haber process to make ammonia.
- Platinum in catalytic converters to reduce nitrogen oxides and carbon monoxide in car exhausts.
Section 19 of 34Coloured compounds (Higher tier)
- Reason
- transition metals form more than one ion and their compounds are often coloured, such as blue copper(II) sulfate containing Cu2+
| Ion | Colour of compounds |
|---|---|
| Fe2+ | Pale green |
| Fe3+ | Brown |
| Cu2+ | Blue |
- Precision
- Fe2+ must be described as pale green, not just green
Section 20 of 34Identifying metal ions with sodium hydroxide (Higher tier)
- Method
- add a few drops of aqueous NaOH to the solution of metal cations and note the precipitate colour
| Ion | Equation | Precipitate |
|---|---|---|
| Fe2+ | Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s) | Pale green |
| Fe3+ | Fe3+(aq) + 3OH-(aq) → Fe(OH)3(s) | Orange-brown |
| Cu2+ | Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s) | Blue |
- State symbols
- give them only when asked
Section 21 of 34Alloys
- Alloy = a mixture of two or more metals, or a metal with a non-metal such as carbon
- Properties
- they can be stronger, harder, and resistant to corrosion or extreme temperatures, which makes them more useful than pure metals
- Not a compound
- an alloy is a mixture and its elements are not chemically combined
- Why alloys are harder
- Atoms of different sizes distort the regular lattice.
- It is harder for the layers to slide over each other.
Section 22 of 34Common alloys and uses
- Steel
- iron with carbon, and different steels are made by varying the carbon content and adding other metals
- Stainless steel
- iron with chromium, nickel and carbon, used in cutlery for hardness and corrosion resistance
- Brass
- copper and zinc, much stronger than either metal, used in musical instruments, ornaments and door knobs
- Iron–tungsten alloys
- extremely hard and resistant to high temperatures
- Iron with chromium or nickel
- resistant to corrosion
- Aluminium alloys
- aluminium with copper, manganese and silicon for aircraft bodies, stronger but still low density
Section 23 of 34Electrolysis of water
- Reaction
- 2H2O(l) → 2H2(g) + O2(g), done in a Hoffman voltameter
- Cathode
- H+ ions are attracted and gain electrons (reduced)
- 2H+(aq) + 2e- → H2(g).
- Anode
- OH- ions are attracted and lose electrons (oxidised)
- 2OH-(aq) → O2(g) + 2H+(aq) + 4e-.
- Gas volumes
- twice as much hydrogen as oxygen forms, because the equation shows 2 mol H2 to 1 mol O2
- Gas tests
- hydrogen gives a squeaky pop with a lit splint, and oxygen relights a glowing splint
- Syllabus note
- the cathode equation must be written, and higher tier must balance the anode equation
Section 24 of 34Electrolysis of aqueous solutions: ions present
- Water
- it is always present and partly ionises, H2O(l) ⇌ H+(aq) + OH-(aq)
- Anode
- OH- and other negative ions (Cl-, Br-) are attracted and lose electrons, such as 2Cl- → Cl2 + 2e-
- Cathode
- H+ and other positive ions (Cu2+, Ag+) are attracted and gain electrons
- Metal at the cathode
- only metals below hydrogen in the reactivity series form there, such as Cu2+ + 2e- → Cu
| Solution | Product at anode | Product at cathode |
|---|---|---|
| Copper(II) chloride | Chlorine | Copper |
| Sodium bromide | Bromine | Hydrogen |
- Collecting gases
- use inverted test tubes filled with water, so the gas displaces the water
Section 25 of 34Competing ions (Higher tier)
- Anode rule
- If a halide ion and OH- are both present, the halide is discharged and forms the halogen.
- If no halide is present, OH- is discharged and forms oxygen.
- The other negative ion stays in solution.
- Cathode rule
- Only one of H+ and a metal ion gains electrons.
- If the metal is above hydrogen in the reactivity series, hydrogen forms and bubbles are seen.
- If the metal is below hydrogen, the metal forms.
| Solution | Ions present | Anode | Cathode |
|---|---|---|---|
| Sodium chloride | H+, OH-, Na+, Cl- | 2Cl-(aq) → Cl2(g) + 2e- | 2H+ + 2e- → H2(g) |
| Copper(II) chloride | H+, OH-, Cu2+, Cl- | 2Cl-(aq) → Cl2(g) + 2e- | Cu2+(aq) + 2e- → Cu(s) |
- Explanation
- in copper(II) chloride copper is below hydrogen so Cu2+ is discharged,
- and in sodium chloride sodium is above hydrogen so H+ is discharged, with chlorine formed at the anode in both
- in copper(II) chloride copper is below hydrogen so Cu2+ is discharged,
- Syllabus note
- anode products for nitrate or sulfate solutions need not be known
Section 26 of 34Practical: electrolysis of copper(II) sulfate (graphite electrodes)
- Apparatus
- 250 cm3 beaker, two graphite electrodes, 12 V DC supply, leads and crocodile clips, and 400 cm3 of about 0.5 mol dm-3 copper(II) sulfate
- Pour the copper(II) sulfate solution into a beaker.
- Place two graphite rods in the solution.
- Connect one electrode to the negative terminal and the other to the positive terminal of a DC supply.
- Fill two small test tubes with copper(II) sulfate solution and invert one over each electrode.
- Switch on the power supply and observe each electrode.
- Test any gas with a glowing splint and a burning splint.
- Record observations at each electrode, including the gas tests.
Section 27 of 34Copper sulfate electrolysis with graphite electrodes: results
- Anode
- bubbling occurs and the gas relights a glowing splint, so it is oxygen
- Cathode
- an orange-brown solid, copper metal, forms on the electrode
- Tip
- the anode is the electrode where gas is produced and bubbles are seen
- Cathode equation
- Cu2+(aq) + 2e- → Cu(s), a reduction
- Anode equation
- 2OH-(aq) → O2(g) + 2H+(aq) + 4e-
Section 28 of 34Copper sulfate electrolysis with copper electrodes: method
- Apparatus
- 250 cm3 beaker, two copper electrodes (foil), 12 V DC supply, leads and crocodile clips, and 400 cm3 of about 0.5 mol dm-3 copper(II) sulfate
- Pour copper(II) sulfate solution into a beaker.
- Weigh and record the mass of a piece of copper foil.
- Connect it to the negative terminal and dip it in the solution.
- Weigh a second piece of foil and connect it to the positive terminal.
- Keep the electrodes apart and switch on the supply.
- Adjust to a constant current and leave for 20 minutes.
- Remove one electrode, wash it with distilled water, dip it in propanone, and let it evaporate dry without wiping.
- Weigh and record its mass.
- Repeat with the other electrode, taking care to know which is which.
- Repeat with fresh electrodes and different currents.
- Practical tip
- make sure the electrode is dry, or the mass will be false
Section 29 of 34Copper sulfate electrolysis with copper electrodes: results and conclusion
- Example data
- the anode falls from 18.0 g to 17.6 g and the cathode rises from 18.0 g to 18.4 g, so the cathode gains and the anode loses 0.4 g
- Anode
- copper atoms are oxidised to ions: Cu(s) → Cu2+(aq) + 2e-
- Cathode
- copper ions are reduced to atoms: Cu2+(aq) + 2e- → Cu(s)
- Mass link
- the mass gained equals the mass lost, so the copper deposited is the same copper lost from the anode
- Concentration
- Cu2+ ions leave and enter the solution at the same rate, so the blue colour is maintained
Section 30 of 34Electroplating
- Electroplating
- coating the surface of one metal with a layer of a different metal
- Anode
- made from the pure metal to be plated
- Cathode
- the object to be plated
- Electrolyte
- an aqueous solution of a soluble salt of the plating metal
- Iron plated with tin (electrolyte tin(II) chloride)
- AnodeSn(s) → Sn2+(aq) + 2e-.
- CathodeSn2+(aq) + 2e- → Sn(s), coating the iron.
- Uses
- to resist corrosion or damage, such as chromium and nickel plating, and to improve appearance, such as silver on cutlery and jewellery
Section 31 of 34Purification of copper (Higher tier)
- Principle
- electrolysis separates a metal from its impurities
- Anode
- impure copper, which becomes thinner as copper atoms lose electrons: Cu → Cu2+ + 2e-
- Cathode
- a thin sheet of pure copper, which becomes thicker as ions gain electrons: Cu2+ + 2e- → Cu
- Electrolyte
- an aqueous solution of a soluble copper salt such as CuSO4
- Sludge
- impurities fall to the bottom of the cell as anode sludge,
- which is valuable and further refined because it often contains precious metals such as silver
Section 32 of 34Electrolysis of brine (Higher tier)
- Brine
- concentrated aqueous sodium chloride, containing Na+, H+, Cl- and OH-
- Cathode
- H+ is discharged because it is less reactive than Na+: 2H+ + 2e- → H2
- Anode
- Cl- is discharged: 2Cl- → Cl2 + 2e-
- Solution left
- Na+ and OH- remain, forming sodium hydroxide solution
| Product | Used to make |
|---|---|
| Chlorine | Bleach |
| Hydrogen | Margarine |
| Sodium hydroxide | Soap and detergents |
- Syllabus note
- the electrode equations are needed, but details of membrane cells are not
Section 33 of 34Economic factors and site of an extraction plant
- Viability
- extracting the metal must give more financial gain than it costs
- Costs
- raw materials, energy, labour and the site
- Site requirements
- Good transport links such as rail and road.
- Near a populated area for a workforce, but not so close that noise is a problem.
- Close to power stations because large amounts of energy are needed.
- Near the coast to import raw materials.
- Common error
- plants are sited near the coast mainly to import raw materials, not just to export products
Section 34 of 34Sustainability and recycling
- Sustainability
- resources meet present needs and remain available for future generations
- Aims
- reduce environmental impact, conserve resources, use energy-efficient methods and renewable energy, minimise waste, recycle,
- and cut the carbon footprint
- reduce environmental impact, conserve resources, use energy-efficient methods and renewable energy, minimise waste, recycle,
- Aluminium
- extraction needs huge amounts of electricity and energy to keep the electrolyte molten, so a power station should be close
- Iron
- raw materials such as haematite and limestone are imported because of cost and the environmental impact of quarrying
- Recycling
- It is economically beneficial, especially for costly metals like aluminium.
- Recycling aluminium needs about 5% of the energy used to extract it from bauxite.
- It is a growing industry that provides jobs.
- It reduces the environmental damage of mining and lessens landfill waste and transport energy.
- Example
- Wylfa power station powered Anglesey Aluminium before decommissioning
Last updated:
🌳 Student Search Tree & Frequently Asked Exam Questions: metals and their extraction
Step-by-step problem-solving guide, mark scheme criteria, and experimental error analysis curated by Dr. Khaled Dabour (DABOURHUB).
Q1: How to solve metals and their extraction examination questions step-by-step?
1. Identify all given values and convert non-standard units, 2. State the primary formula or physiological/chemical principle, 3. Calculate to 3 significant figures and clearly state units.
Q2: What are the common mistakes and experimental sources of error for metals and their extraction?
Avoid confusing independent vs dependent variables in practical questions. In calculation questions, double-check percentage change formula: (change / original) × 100.
🔍 Related Student Searches & Problem Trees for metals and their extraction:
🏷️ how to solve metals and their extraction step by step🏷️ step by step calculation formula for metals and their extraction🏷️ common student errors in metals and their extraction exams🏷️ sources of error and investigation variables for metals and their extraction🏷️ interactive revision quiz on metals and their extractionmetals and their extraction past paper questions by topic with mark schemewhere to find metals and their extraction exam questions with answersmetals and their extraction mark scheme explained step by stepwhy is metals and their extraction the answer in this mcqmetals and their extraction 4 mark question with answermetals and their extraction 6 mark question mark scheme wordingmetals and their extraction mcq practice with answers and explanationshow to revise metals and their extraction in one day before the exammetals and their extraction short revision notes with a quizmetals and their extraction common mistakes students make in examsmetals and their extraction sources of error alternative to practicalmetals and their extraction command words for paper 4 answers
Where can I find metals and their extraction past paper questions by topic with the mark scheme?
On DABOURHUB every metals and their extraction question is grouped by syllabus point and shown with the mark-scheme wording, so you can see exactly which sentence earns the mark instead of scrolling through whole year papers.
What is the most common metals and their extraction exam question and why do students lose marks on it?
The repeated pattern on IGCSE papers is a short stem that asks you to apply metals and their extraction and then justify it with a command word (explain, suggest, compare). Marks are lost by stating the fact without linking it back to the wording of the question, so each metals and their extraction item here is written as a full-mark answer.
Is there an alternative to Save My Exams or Physics and Maths Tutor for metals and their extraction?
Yes: the metals and their extraction notes, topical questions and MCQ quizzes on this site follow the same syllabus headings as the official specification, and they open straight from the search result with no download step.
How do I revise metals and their extraction in two weeks before the exam?
Three passes: read the metals and their extraction notes once, then do the topical questions and mark them against the mark scheme, then redo only the ones you missed. The quiz on the page gives instant feedback so the third pass takes minutes rather than an evening.