2. Extraction of Iron and Steelmaking

 Doc Brown's Chemistry GCSE/IGCSE/O Level Revision Notes - Mining of Minerals, Methods of Extracting of Metals from Ores These revision notes on the extraction of copper and the electrolytic refining of copper, useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.

Summary for this page

How do we extract iron from its mineral ores like the iron oxide ores haematite and magnetite? How do we convert it into steel? The raw materials needed i.e. limestone, coke, air and iron ore and the chemistry of the blast furnace are fully described. Why convert iron into steel? How do you make steel? The manufacture of steel alloys is further described. Scroll down for revision notes on extraction procedures and theory which should prove useful for school/college assignments/projects on ways of extracting metals from their ores.

Equation note:

The equations are sometimes written three times: (i) word equation, (ii) balanced symbol equation without state symbols, and, (iii) with the state symbols (g), (l), (s) or (aq) to give the complete balanced symbol equation.

2a. Extraction of Iron in a blast furnace

Sadly, now shut down!

Raw Materials

• Iron Ore e.g. haematite ore (iron(III) oxide) the source of iron.

  • Fe₂O₃

• or magnetite ore

  • Fe₃O₄

• coke (carbon, C), both fuel and reducing agent.

• hot air (for the oxygen in it) to burn the coke

  • O₂

• limestone (calcium carbonate) to remove certain impurities like silica.

  • CaCO₃

• Iron ore is used to make iron and steel eg haematite and magnetite. Iron is produced (extracted) in a blast furnace by reducing iron oxides with carbon and it is the carbon that removes the oxygen from the iron oxides – the carbon is known as the reducing agent. Coke is a cheap and readily made reducing agent.
• Iron is not a very reactive metal, so, because its position in the reactivity series of metals, iron can be extracted using carbon in a blast furnace because iron is below carbon (iron is less reactive than carbon). Therefore, iron can be displaced from its eg its oxides, by heating with the theoretically 'more reactive' carbon in a sort of displacement reaction.
• Iron oxide ore is mined in many parts of the world.
  • Examples of rich or high quality ores are
  • haematite, mainly iron(III) oxide, formula Fe₂O₃
  • and magnetite, triiron tetroxide, formula Fe₃O₄.
• A solid mixture of magnetite/haematite ore, coke and limestone is continuously fed into the top of the blast furnace.
• The double role and function of coke (carbon)
• 1st Coke function (i) As a fuel
  • The coke is ignited at the base and hot air blown in to burn the coke (carbon) to form carbon dioxide in an oxidation reaction (C gains O).
  • The heat energy is needed from this very exothermic reaction to raise the temperature of the blast furnace to over 1000^oC to effect the ore reduction. The furnace contents must be he
  • carbon + oxygen ==> carbon dioxide
    • C + O₂ ==> CO₂
      • C(s) + O₂(g) ==> CO₂(g)  (equation with state symbols)
      • The carbon is oxidised to carbon dioxide and in doing so releases a lot of heat energy.
• 2nd Coke function (ii) As a reducing agent
  • At high temperatures the carbon dioxide formed, reacts with more coke (carbon) to form carbon monoxide
  • carbon dioxide + carbon ==> carbon monoxide
    • CO₂ + C ==> 2CO
      • CO₂(g) + C(s) ==> 2CO(g)  (equation with state symbols)
  • Note that in this reaction carbon dioxide, CO₂, is reduced by oxygen loss to the carbon monoxide, and the carbon is oxidised by oxygen, O gain to carbon monoxide.
  • A rather curious case where both an oxidation and a reduction give the same product!
• The carbon monoxide is the molecule that actually removes the oxygen from the iron oxide ore – this is the action of a reducing agent in the context of producing iron from iron oxide ores.
  • REDOX definition reminders – reduction is a process of oxygen loss (or electron gain) and oxidation is a process of oxygen gain (or electron loss).
  • This a reduction reaction, which can be described in two ways.
  • The Fe₂O₃ loses its oxygen O (reduction)
    • or for more advanced students
      • the ion Fe³⁺ gains three electrons to form Fe,
        • noting that electron gain is reduction.
  • The carbon monoxide, CO, is known as the reducing agent because it is the oxygen (O) remover and gets oxidised to carbon dioxide in the process (CO gains oxygen).
• This frees the iron, which is molten at the high blast furnace temperature, and trickles down to the base of the blast furnace and run off.
  • An example of the main reduction reaction is ...
  • iron(III) oxide + carbon monoxide ==> iron + carbon dioxide
    • Fe₂O₃ + 3CO ==> 2Fe + 3CO₂
      • Fe₂O₃(s) + 3CO(g) ==> 2Fe(l/s) + 3CO₂(g)  (equation with state symbols)
      • Note on state symbols (l/s), iron is initially formed as a liquid and then obviously solidifies on cooling.
      • The iron oxide is reduced due to oxygen loss to give free iron.
      • The carbon monoxide is oxidised, gains oxygen to form carbon dioxide.
  • note, as in the two reactions above, oxidation and reduction always go together!
    • Other possible iron ore reduction reactions are direct reduction of the iron oxide by carbon itself ...
    • iron(III) oxide + carbon ==> iron + carbon monoxide
    • Fe₂O₃ + 3C ==> 2Fe + 3CO
      • Fe₂O₃(s) + 3C(g) ==> 2Fe(l/s) + 3CO(g)  (equation with state symbols)
    • or
    • iron(III) oxide + carbon ==> iron + carbon dioxide
    • 2Fe₂O₃ + 3C ==> 4Fe + 3CO₂
      • 2Fe₂O₃(s) + 3C(g) ==> 4Fe(l/s) + 3CO₂(g)  (equation with state symbols)
    • I'm afraid there are quite a few possible equations and most do occur to some extent somewhere in the blast furnace.
    • In each case the iron oxide is reduced to iron by oxygen loss.
    • The carbon is oxidised to carbon dioxide by oxygen gain.
    • Note, that although it doesn't seem obvious, in all these reductions of the iron ores with carbon or carbon monoxide, it is the iron(III) ion that is actually reduced, so the principal chemical change of importance is ...
      • Fe³⁺  +  3e^–  ===>  Fe
  • The iron is initially formed in its liquid state because of the high temperatures of the blast furnace (well over 1000^oC) but when cooled it is cast into solid ingots, or the liquid iron can be transported directly in special insulated 'torpedo' wagons to a steel making plant on the same industrial site complex..
  • At the highest temperatures in a blast furnace the reactions can be written as a direct reduction of the oxide with carbon, and, carbon monoxide (CO) can be formed as well as carbon dioxide (CO₂) e.g.
    • for haematite
      • iron(III) oxide + carbon ==> iron + carbon monoxide
      • Fe₂O₃ + 3C ==> 2Fe + 3CO
        • Fe₂O₃(s) + 3C(s) ==> 2Fe(l/s) + 3CO(g)  (equations with state symbols)
    • or 
      • iron(III) oxide + carbon ==> iron + carbon dioxide
      • 2Fe₂O₃ + 3C ==> 4Fe + 3CO₂
        • 2Fe₂O₃(s) + 3C(s) ==> 4Fe(l/s) + 3CO₂(g)
    • for magnetite
      • 'iron tetroxide' + carbon ==> iron + carbon monoxide
      • Fe₃O₄ + 4C ==> 3Fe + 4CO
        • Fe₃O₄(s) + 4C(s) ==> 3Fe(l/s) + 4CO(g)
    • or 
      • 'iron tetroxide' + carbon ==> iron + carbon dioxide
      • Fe₃O₄ + 2C ==> 3Fe + 2CO₂
        • Fe₃O₄(s) + 2C(s) ==> 3Fe(l/s) + 2CO₂(g)
    • Name note: The correct name for iron tetroxide is diiron(III)iron(II) oxide, but I wouldn't worry about it!
    • Again, in each case the iron oxide is reduced to iron by oxygen loss, but in terms of electrons the reduction is ...
      • Fe³⁺  +  3e^–  ===>  Fe
    • At the same time the carbon is oxidised to carbon dioxide by oxygen gain.
  • I'm afraid the chemistry of the blast furnace can get very complicated indeed!
• The role of limestone in the extraction of iron
  • The original ore contains acidic mineral impurities such as silica (SiO₂, silicon dioxide).
  • These react with the calcium carbonate (limestone) to form a molten slag, the main ingredient being calcium silicate.
  • There are two ways to show the formation of the waste slag, which is mainly calcium silicate.
  • (i) calcium carbonate + silica ==> calcium silicate + carbon dioxide
    • (i) CaCO₃ + SiO₂ ==> CaSiO₃ + CO₂
    • Reaction (i) is a sort of displacement reaction i.e. the less volatile high melting/boiling silicon dioxide (silica) displaces the more volatile gaseous carbon dioxide.
  • However, this is sometimes shown in two stages, i.e. reactions (ii) and (iii):
    • (ii) CaCO₃ ==> CaO + CO₂
    • (iii) CaO + SiO₂ ==> CaSiO₃
    • (ii) is the thermal decomposition of calcium carbonate into calcium oxide and carbon dioxide, and the reaction needs a high temperature of over 900^oC, but that's no problem in the blast furnace!
    • (iii) is the combination of the basic calcium oxide and the acidic silicon dioxide to form calcium silicate.
  • However, whichever way you represent the reaction, its all the same in the end.
    • If you 'add up' chemical reactions (ii) and (iii) you get (i), check for yourself.
• The molten slag forms a layer above the more dense molten iron and they can be both separately, and regularly, drained away. The iron is cooled and cast into pig iron ingots OR transferred directly to a steel producing furnace.
• The waste gases and dust from the blast furnace must be appropriately treated to avoid polluting the environment.
  • The highly toxic carbon monoxide can actually be burnt to provide a source of heat energy, and in the exothermic reaction it is converted into relatively harmless carbon dioxide.
    • carbon monoxide + oxygen ==> carbon dioxide
    • 2CO + O₂ ==> 2CO₂
      • 2CO(g) + O₂(g) ==> 2CO₂(g)
      • This the oxidation of harmful carbon monoxide to harmless carbon dioxide.
  • Acidic gases like sulfur dioxide from sulfide ores, can be removed by bubbling through an alkali solution such as calcium hydroxide solution ('limewater') where it is neutralised and oxidised to harmless calcium sulphate. Cleaning a gas in this way is called 'gas scrubbing'.
  • Any contaminated water must be purged of harmful chemicals before being released into a river or recycled via water treatment plant.
  • The waste slag is used for road construction or filling in quarries which can then be landscaped.
• Iron from a blast furnace is ok for very hard cast iron objects BUT is too brittle for many applications due to too high a carbon content from the coke.

2b. STEEL MAKING

• Why do we need to convert most iron into steel alloys?

  • Most metals in everyday use are alloys.

  • Iron is a good conductor of heat and can be bent or hammered into shape (malleable), quite strong physically – made stronger when alloyed with other materials.

  • This makes iron useful as structural material and for making things that must allow heat to pass through easily and useful construction materials.

  • Steel is an alloy because it is a mixture of a metal (iron) with other elements (carbon and perhaps other metals too).

  • Iron from the blast furnace contains about ~96% iron with ~4% of impurities including carbon, silica and phosphorus.

  • In this state the cast iron is too hard and too brittle for most purposes.

  • Cast iron is hard and can be used directly for some purposes eg manhole covers, ornamental railings because of its strength in compression and is very hard wearing.

  • However, if all the impurities are removed, the resulting very pure iron is too soft for any useful purpose.

  • Therefore, strong useful steel is made by controlling the amount of carbon and selected metals to produce an alloy mixture with the right physical properties fit for a particular application e.g. steel for car bodies, chrome stainless steel, extremely hard and tough tungsten–iron steel alloys etc.

  • The real importance of alloys is that they can be designed to have properties for specific uses in terms of eg compression/tensile strength or corrosion reduction ie less susceptible to rusting.

  • Low–carbon steels (0.1% carbon) are easily shaped for car bodies

  • High–carbon steels (1.5% carbon, often with other metals too) are hard wearing and inflexible and can be used for cutting tools, bridge construction.

  • Stainless steels have chromium (and maybe nickel) added and are much resistant to corrosion (from oxygen/water) than iron or plain steel which readily rust. Objects made of iron or plain steel, particularly those exposed to the weather, regularly have to be painted or coated with some other protective layer from the effects of water and oxygen.

• The properties of iron can be altered by adding small quantities of other metals or carbon to make steel, one of the useful metal alloys in widespread use today.

• Steels are alloys since they are mixtures of iron with other metals or with non–metals like carbon or silicon.

• Making Steel:

  • (1) Molten iron from the blast furnace is mixed with recycled scrap iron

  • (2) Then pure oxygen is passed into the mixture and the non–metal impurities such as silicon or phosphorus are then converted into acidic oxides (the BOS oxidation process) ..

    • e.g. Si + O₂ ==> SiO₂, or 4P + 5O₂ ==> P₄O₁₀

  • (3) Calcium carbonate (a base) is then added to remove the acidic oxide impurities (in an acid–base reaction). The salts produced by this reaction form a slag which can be tapped off separately.

    • e.g. CaCO₃ + SiO₂ ==> CaSiO₃ + CO₂ (calcium silicate slag)

  • Steps (1)–(3) produce pure iron.

  • Calculated quantities of carbon and/or other metallic elements such as titanium, manganese or chromium are then added to make a wide range of steels with particular properties.

  • Because of the high temperatures the mixture is stirred by bubbling in unreactive argon gas!

  • Economics of recycling scrap steel or ion: Most steel consists of >25% recycled iron/steel and you do have the 'scrap' collection costs and problems with varying steel composition.

    • BUT you save enormously because there is no mining cost or overseas transport costs AND less junk lying around

    • Note that some companies send their own scrap to be mixed with the next batch of 'specialised' steel they order, this saves both companies money!). More on this below.

2c. NOTE on RECYCLING Iron and Steel

• About 42% of iron/steel in goods/components manufactured in iron/steel is recycled iron/steel, whether it be steel pans, car bodies, bridge girders, stainless steel cutlery etc.

• This makes good economics because recycling saves on several costs AND allows a mineral resource like iron's haematite or magnetite ore resources to last a lot longer – slower depletion of the Earth's mineral ore resources will make it last longer.

  • Transport costs may be less (ie within UK now), but much more importantly

    • mining costs are omitted – energy/machinery involved in digging out the ore, crushing it, transporting the ore,

    • and the cost of actually extracting the metal from its finite ore resource – chemicals needed (coke, limestone), constructing and running a blast furnace

• So, scrap metal merchants are doing a roaring trade at the moment.

• The savings are partly reduced by the cost off collecting waste/scrap metal.

• There are particular problems due to the varying composition of alloys, but if the composition is known, or obtained from chemical analysis, the different compositions can be blended together to a desired alloy composition.

• The social, economic and environmental impacts of exploiting metal ores including RECYCLING are discussed further on a separate page.

WHERE NEXT?

Other associated KS4 Science GCSE/IGCSE chemistry web pages on this site

• Word-fill quiz "Extraction of Iron"

• GCSE/IGCSE Foundation–tier (easier) Multiple choice QUIZ on METAL EXTRACTION

• or GCSE/IGCSE Higher–tier (harder) Multiple choice QUIZ on METAL EXTRACTION

• and gap–word–fill exercise on METAL EXTRACTION

• Reactivity of metals : all the reactions of metals with oxygen (air), water and acids, displacement reactions

• Oxidation–reduction explained (redox theory)

• Group 1 Alkali Metals: chemical and physical properties

• Extra Industrial chemistry : properties and uses of metals including titanium and aluminium

• Transition metals physical and chemical properties and uses

• Metallic Bonding – structure and properties of metals

• Extra Electrochemistry–electrolysis theory, experiment designs and links to all electrolysis sections

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