GCSE-IGCSE-KS4 Science-CHEMISTRY Revision-Information Notes on

 The Mining of Minerals and Methods of Extracting of Metals   Summary notes on extraction procedures

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Multiple choice Quiz: Foundation or Higher and word-fill

Other associated web pages on this site:

• Reactivity of metals : all the reactions with oxygen (air), water and acids, displacement reactions and oxidation-reduction explained (redox theory)
.

• Group 1 Alkali Metals: chemical and physical properties

• Extra Industrial chemistry : properties and uses of metals

• Transition metals : properties and uses

• Extra Electrochemistry-electrolysis : electrolysis theory, experiment designs and links to all electrolysis section on this site.

• The Earth's crust contains many different rocks. Rocks are a mixture of minerals and from some we can make useful substances.

• A mineral can be a solid metallic or non-metallic element or a compound found naturally in the Earth's crust.

• A metal ore is a mineral or mixture of minerals from which economically viable amounts of metal can be extracted, i.e. its got to have enough of the metal, or one of its compounds, in it to be worth digging out! Ores are often oxides, carbonates or sulphides. They are all finite resources so we should use them wisely!

• In order to extract a metal, the ore or compound of the metal must undergo a process called reduction to free the metal (i.e. the positive metal ion gains negative electrons to form the neutral metal atom, or the oxide loses oxygen, to form the free metallic atoms).

  • The chemical that removes the oxygen from an oxide is called the reducing agent i.e. carbon, carbon monoxide or sometimes hydrogen.

  • Detailed REDOX notes on the metal reactivity page, and on this page where appropriate. 

• Generally speaking the method of extraction depends on the metals position in the reactivity series

• The reactivity series of metals can be presented to include two non-metals, carbon and hydrogen, to help predict which method could be used to extract the metal.

  • lower Pt Au Ag Cu (H) Pb Sn Fe Zn (C) Al Mg Ca Na K higher in series

  • RULE: Any element higher in the series can displace any other lower element

• Metals above zinc and carbon in the reactivity series cannot usually be extracted with carbon or carbon monoxide. They are usually extracted by electrolysis of the purified molten ore or other suitable compound

  • e.g. aluminium from molten aluminium oxide or sodium from molten sodium chloride.

  • The ore or compound must be molten or dissolved in a solution in an electrolysis cell to allow free movement of ions (electrical current). Theory given in the appropriate sections.

• Metals below carbon can be extracted by heating the oxide with carbon or carbon monoxide. The non-metallic elements carbon will displace the less reactive metals in a smelter or  blast furnace e.g. iron or zinc and metals lower in the series.

  • Metals below hydrogen will not displace hydrogen from acids. Their oxides are easily reduced to the metal by heating in a stream of hydrogen, though this is an extraction method rarely used in industry. In fact most metal oxides below carbon can be reduced when heated in hydrogen, even if the metal reacts with acid.

• Some metals are so unreactive that they do not readily combine with oxygen in the air or any other element present in the Earth's crust, and so can be found as the metal itself. For example gold (and sometimes copper and silver) and no chemical separation or extraction is needed. In fact all the metals below hydrogen can be found as the 'free' or 'native' element.

• Other methods are used in special cases using the displacement rule. A more reactive metal can be used to displace and extract a less reactive metal but these are costly processes since the more reactive metal also has to be produced in the first place! See Titanium or see at the end of the section on copper extraction.

• Sometimes electrolysis is used to purify less reactive metals which have previously been extracted using carbon or hydrogen (e.g. copper and zinc). Electrolysis is also used to plate one metal with another.

• The demand for raw materials does have social, economic and environmental implications e.g. conservation of mineral resources by recycling metals, minimising pollution etc.

• Historically as technology and science have developed the methods of extraction have improved to the point were all metals can be produced. The reactivity is a measure of the ease of compound formation and stability (i.e. more reactive, more readily formed stable compound, more difficult to reduce to the metal).

  • The least reactive metals such as gold, silver and copper have been used for the past 10000 years because the pure metal was found naturally.

  • Moderately reactive metals like iron and tin have been extracted using carbon based smelting for the past 2000-3000 years.

  • BUT it is only in the last 200 years that very reactive metals like sodium or aluminium have been extracted by electrolysis.

• Iron oxide ore is mined in many parts of the world. Examples are haematite Fe₂O₃ and magnetite Fe₃O₄.
• The solid mixture of haematite ore, coke and limestone is continuously fed into the top of the blast furnace.
• 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_(s) + O_2(g) ==> CO_2(g)
• at high temperature the carbon dioxide formed, reacts with more coke (carbon) to form carbon monoxide
  • carbon dioxide + carbon ==> carbon monoxide
  • CO_2(g) + C_(s) ==> 2CO_(g)
  • (note: CO₂ reduced by O loss, C is oxidised by O gain)
• The carbon monoxide is the molecule that actually removes the oxygen from the iron oxide ore. This a reduction reaction (Fe₂O₃ loses its O, or Fe³⁺ gains three electrons to form Fe) and the CO is known as the reducing agent (the O remover and gets oxidised in the process).
• This frees the iron, which is molten at the high blast furnace temperature, and trickles down to the base of the blast furnace. The main reduction reaction is ...
  • iron(III) oxide + carbon monoxide ==> iron + carbon dioxide
  • Fe₂O_3(s) + 3CO_(g) ==> 2Fe_(l) + 3CO_2(g)
  • note, as in the two reactions above, oxidation and reduction always go together!
    • Other possible ore reduction reactions are ...
    • iron(III) oxide + carbon ==> iron + carbon monoxide
    • Fe₂O_3(s) + 3C_(g) ==> 2Fe_(l) + 3CO_(g)
    • or
    • iron(III) oxide + carbon ==> iron + carbon dioxide
    • 2Fe₂O_3(s) + 3C_(g) ==> 4Fe_(l) + 3CO_2(g)
• 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 of e.g. calcium silicate.
  • calcium carbonate + silica ==> calcium silicate + carbon dioxide
  • CaCO₃ + SiO₂ ==> CaSiO₃ + CO₂
  • this is sometimes shown in two stages:
    • CaCO₃ ==> CaO + CO₂
    • CaO + SiO₂ ==> CaSiO₃
• 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_(g) + O_2(g) ==> 2CO_2(g)
  • Acidic gases like sulphur dioxide from sulphide 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. So it is converted into steel alloys for a wide range of uses.

Raw Materials:

• Iron Ore e.g. haematite ore [iron(III) oxide, Fe₂O₃]

• coke (carbon, C)

• hot air (for the O₂ in it)

• limestone (calcium carbonate, CaCO₃)

• Aluminium is obtained from mining the mineral bauxite.
• The purified bauxite ore of aluminium oxide is continuously fed in. Cryolite is added to lower the melting point and dissolve the ore.
• Ions must be free to move to the electrode connections called the cathode (-, negative), attracting positive ions e.g. Al³⁺, and the anode (+, positive) which attracts negative ions e.g. O^2-.
• When the d.c. current is passed through aluminium forms at the negative cathode (metal*) and sinks to the bottom of the tank.
• At the positive anode, oxygen gas is formed (non-metal*). This is quite a problem. At the high temperature of the electrolysis cell it burns and oxidises away the carbon electrodes to form toxic carbon monoxide or carbon dioxide. So the electrode is regularly replaced and the waste gases dealt with! 
• It is a costly process (6x more than Fe!) due to the large quantities of expensive electrical energy needed for the process.
• * Two general rules:
  • Metals and hydrogen (from positive ions), form at the negative cathode electrode.
  • Non-metals (from negative ions), form at the positive anode electrode.

Raw materials for the electrolysis process:

• Bauxite ore of impure aluminium oxide [Al₂O₃ made up of Al³⁺ and O^2- ions]

• Carbon (graphite) for the electrodes.

• Cryolite reduces the melting point of the ore and saves energy, because the ions must be free to move to carry the current

• Electrolysis means using d.c. electrical energy to bring about chemical changes e.g. decomposition of a compound to form metal deposits or release gases. The electrical energy splits the compound!

• At the electrolyte connections called the anode electrode (+, attracts - ions) and the cathode electrode (-, attracts + ions). An electrolyte is a conducting melt or solution of freely moving ions which carry the charge of the electric current.

• At the negative  (-) cathode, reduction occurs (electron gain) when the positive aluminium ions are attracted to it. They gain three electrons to change to neutral Al atoms.
  • Al³⁺ + 3e^- ==> Al
• At the positive (+) anode, oxidation takes place (electron loss) when the negative oxide ions are attracted to it. They lose two electrons forming neutral oxygen molecules.
  • 2O^2- ==> O₂ + 4e^- 
  • or 2O^2- - 4e^- ==> O₂ 
• Note: Reduction and Oxidation always go together!
• The overall electrolytic decomposition is ...
  • aluminium oxide => aluminium + oxygen
  • 2Al₂O₃ ==> 4Al + 3O₂
  • and is a very endothermic process, lots of electrical energy input!

• GENERAL NOTE ON ELECTROLYSIS:
• Any molten or dissolved material in which the liquid contains free moving ions is called the electrolyte.
• Ions are charged particles e.g. Na⁺ sodium ion, or Cl^- chloride ion, and their movement or flow constitutes an electric current, because a current is moving charged particles.
• What does the complete electrical circuit consist of?
  • There are two ion currents in the electrolyte flowing in opposite directions:
    • positive cations e.g. Al³⁺ attracted to the negative cathode electrode,
    • and negative anions e.g. O^2- attracted to the positive anode electrode,
    • BUT remember no electrons flow in the electrolyte, only in the graphite or metal wiring!
  • The circuit of 'charge flow' is completed by the electrons moving around the external circuit e.g. copper wire or graphite electrode, from the positive to the negative electrode
  • This e^- flow from +ve to -ve electrode perhaps doesn't make sense until you look at the electrode reactions, electrons released at the +ve anode move round the external circuit to produce the electron rich negative cathode electrode.
• Electron balancing: In the above process it takes the removal of four electrons from two oxide ions to form one oxygen molecule and the gain of three electrons by each aluminium ion to form one aluminium atom. Therefore for every 12 electrons you get 3 oxygen molecules and 4 aluminium atoms formed. This means you can do mole ratio product calculations. See section 13. on the Chemical Calculations page.

• The impure copper from a smelter is cast into a block to form the positive anode. The cathode is made of previously purified copper. These are dipped into an electrolyte of copper(II) sulphate solution. 
• When the d.c electrical current is passed through the solution electrolysis takes place.  The copper anode dissolves forming blue copper(II) ions Cu²⁺.
• These positive ions are attracted to the negative cathode and become copper atoms. The mass of copper dissolving at the anode exactly equals the mass of copper deposited on the cathode. The concentration of the copper(II) sulphate remains constant.
• Any impurities present in the impure copper anode fall to the bottom of the electrolysis cell tank. This 'anode sludge' is not completely mineral waste, it can contain valuable metals such as silver!
• See section below for extraction of impure copper from an ore.

Raw materials for the electrolysis process:

• Impure copper from a copper smelter.

• Electrolyte of aqueous copper(II) sulphate.

• A pure copper cathode.

Electrolysis is using d.c. electrical energy to bring about chemical changes at the electrolyte connections called the anode and cathode  electrodes.

An electrolyte is a conducting melt or solution of ions which carry the electric charge as part of the circuit.

Scrap copper can be recycled and purified this way too ,and is cheaper than starting from copper ore AND saves valuable mineral resources.

• At the positive (+) anode, the process is an oxidation, electron loss, as the copper atoms dissolve to form copper(II) ions.

Cu_(s) ==> Cu²⁺_(aq) + 2e^-

• at the negative (-) cathode, the process is a reduction, electron gain by the attracted copper(II) ions to form neutral copper atoms.

Cu²⁺_(aq) + 2e^- ==> Cu_(s)

• Note: Reduction and Oxidation always go together, hence the use of the term redox change or reaction.
• Electroplating is mentioned on the Industrial Chemistry and Electrochemistry pages.

The original extraction of copper from copper ores

• From copper carbonate ores* ...
  • The ore can be roasted to concentrate the copper as its oxide.
  • Water is driven off and the carbonate thermally decomposed.
  • copper(II) carbonate ==> copper oxide + carbon dioxide
  • CuCO_3(s) ==> CuO_(s) + CO_2(g)
  • The oxide can be smelted by heating with carbon (coke, charcoal) to reduce the oxide to impure copper, though this method isn't really used much these days (the 'bronze age' method archaeologically!).
  • copper(II) oxide + carbon ==> copper + carbon dioxide
  • 2CuO_(s) + C_(s) ==> 2Cu_(s) + CO_2(g)
  • The carbon acts as the reducing agent - the 'oxygen remover'.

• From copper sulphide ores ...
  • These include chalcocite/chalcosine = copper(I) sulphide Cu₂S and covellite = copper(II) sulphide CuS
    • and chalcopyrite CuFeS₂. which is one of the most important ores for the extraction of copper.
      • This can be roasted in air to produce copper(I) sulfide which is roasted again in a controlled amount of air so as not to form a copper oxide (see below).
      • 2CuFeS₂ +  4O₂ ==> Cu₂S + 3SO₂ + 2FeO
  • Copper sulphide ores can be rapidly roasted in heated air enriched with oxygen to form impure copper and this extraction process is called 'flash smelting'.
    • Nasty sulphur dioxide gas is formed, this must be collected to avoid pollution and can be used to make sulphuric acid to help the economy of the process.
    • copper(I) sulphide + oxygen ==> copper + sulphur dioxide
      • Cu₂S_(s) + O_2(g) ==> 2Cu_(s) + SO_2(g)
    • or copper(II) sulphide + oxygen ==> copper + sulphur dioxide
      • CuS_(s) + O_2(g) ==> Cu_(s) + SO_2(g)
• It is also possible to dissolve an oxide or carbonate ore in dilute sulphuric acid and extracting copper by ....
  • (1) using electrolysis see purification by electrolysis above, or
  • (2) by adding a more reactive metal to displace it e.g. scrap iron or steel is used by adding it to the resulting copper(II) sulphate solution.
    • iron + copper(II) sulphate ==> iron(II) sulphate + copper
    • Fe_(s) + CuSO_4(aq) ==> FeSO_4(aq) + Cu_(s)

The Extraction of Chromium and Titanium by Displacement

• Titanium ore is mainly the oxide TiO₂, which is converted into titanium tetrachloride TiCl₄ by heating with carbon and chlorine.

• The chloride is then reacted with sodium or magnesium to form titanium metal and sodium chloride or magnesium Chloride.

• This reaction is carried out in an atmosphere of inert argon gas so non of the metals involved becomes oxidised by atmospheric oxygen.

  • TiCl₄ + 2Mg ==> Ti + 2MgCl₂ or TiCl₄ + 4Na ==> Ti + 4NaCl

• Overall the titanium oxide ore is reduced to titanium metal (overall O loss, oxide => metal) and the magnesium or sodium acts as a reducing agent.

• Chromium ore is processed and purified into chromium(III) oxide. This is reacted, very exothermically, in a thermit style reaction, with aluminium (see reactions of aluminium) to free the chromium metal.

  • Cr₂O_3(s) + 2Al_(s) ==> Al₂O_3(s) + 2Cr_(s) 

  • The chromium(III) oxide is reduced to chromium by O loss, the aluminium is oxidised to aluminium oxide by O gain, and the aluminium is the reducing agent i.e. the O remover.

• These are examples of metal displacement reactions e.g. the less reactive chromium or titanium are displaced by the more reactive sodium, magnesium or aluminium.

The Extraction and Purification of Zinc

• Zinc is extracted from either zinc blende/sphalerite ore (zinc sulphide) or sometimes calamine/Smithsonite ore (zinc carbonate).
• (1) The zinc sulphide ore is roasted in air to give impure zinc oxide.
  • 2ZnS_(s) + 3O_2(g) ==> 2ZnO_(s) + 2SO_2(g)
  • Note: calamine ore can be used directly in a zinc smelter because on heating it also forms zinc oxide.
    • ZnCO_3(s)  ==> ZnO_(s) + CO_2(g) (endothermic thermal decomposition)
• (2) The impure zinc oxide can be treated in two ways to extract the zinc:
  • (a) It is roasted in a smelting furnace with carbon (coke, reducing agent) and limestone (to remove the acidic impurities). The chemistry is similar to iron from a blast furnace.
    • C_(s) + O_2(g) ==> CO_2(g) (very exothermic oxidation, raises temperature considerably)
    • C_(s) + CO_2(g) ==> 2CO_(g) (C oxidised, CO₂ reduced)
    • ZnO_(s) + CO_(g) ==> Zn_(l) + CO_2(g) (zinc oxide reduced by CO, Zn undergoes O loss)
    • or direct reduction by carbon: ZnO_(s) + C_(s) ==> Zn_(l) + CO_(g) (ZnO reduced, C oxidised)
    • The carbon monoxide acts as the reducing agent i.e. it removes the oxygen from the oxide.
    • The impure zinc is  then fractionally distilled from the mixture of slag and other metals like lead and cadmium out of the top of the furnace in an atmosphere rich in carbon monoxide which stops any zinc from being oxidised back to zinc oxide.
    • The slag and lead (with other metals like cadmium) form two layers which can be tapped off at the base of the furnace.
    • The zinc can be further purified by a 2nd fractional distillation or more likely by dissolving it in dilute sulphuric acid and purified electrolytically as described below.
  • (b)Two stages
    • (i) It is dissolved and neutralised with dilute sulphuric acid to form impure zinc sulphate solution.
    • ZnO_(s) + H₂SO_4(aq) ==> ZnSO_4(aq) + H₂O_(l)
    • or using calamine ore/zinc carbonate directly:
      • ZnCO_3(s) + H₂SO_4(aq) ==> ZnSO_4(aq) + H₂O_(l)+ CO_2(g)
    • (ii) Quite pure zinc is produced from the solution by electrolysis. It can be deposited on a pure zinc negative electrode (cathode) in the same way copper can be purified. The other electrode, must be inert e.g. for laboratory experiments, carbon (graphite) can be used and oxygen is formed.
      • Zn²⁺_(aq) + 2e^- ==> Zn_(s)
        • A reduction process, electron gain, as zinc metal is deposited on the (-) electrode.
      • You can't use solid zinc oxide directly because its insoluble and the ions must free to carry the current and migrate to the electrodes in some sort of solution.
      • For more details of the type of electrolysis system used, see purification of copper (just swap Zn for Cu in the method/diagram).
      • PLEASE note: In the industrial production of zinc by electrolysis (called electro-winning) the negative (-) cathode is made of aluminium (Al, where zinc deposits) and the positive (+) electrode is made of a lead-silver alloy (Pb-Ag, where oxygen gas is formed). Why these particular electrode metals are used in this 'electrowinning' process I'm not quite sure, but aluminium is so unreactive that it is effectively inert, and lead and silver are also of low activity, but ... ???

The electrolytic extraction of sodium

The positive sodium ions migrate to the negative cathode electrode and are reduced by electron gain to form liquid sodium atoms.

Na⁺ + e^- ==> Na

The negative chloride ions migrate to the positive anode electrode and get oxidised by electron loss to form green chlorine gas molecules.

2Cl^- ==> Cl₂ + 2e^-

Environmental Impact and Economics of Metal and other Mineral Extraction - Sociological, environmental issues etc.

• One of the problems of metal or mineral extraction is balancing ecological, environmental, economic, social advantage factors.

• It doesn't matter whether you are mining and processing iron ore or limestone, many of the advantages and disadvantages are common to these operations.

• Examples of advantages of a country exploiting it's own mineral resources:

  • Valuable revenue if the mineral or its products are exported.

  • Jobs for people, especially new sources of employment in poor countries or areas of high unemployment in developed countries.

  • Wages earned go into the local economy.

  • Increase in local facilities promoted e.g. transport systems, like roads, recreational and health social facilities.

• Examples of disadvantages of a country exploiting it's own mineral resources and reduction of its social and environmental impact:

  • Dust from mining-quarrying or processing can be reduced by air filter and precipitation systems and even hosing water on dusty areas or spoil heaps or carried away to somewhere else via tall chimneys.

  • Noise from process operation or transport of raw materials and products (lorries/trucks/wagons).

    • Difficult to deal with, sound-proofing often not practical, but operations can be reduced for unsociable hours e.g. evening movement.

  • Pollution can be reduced by cleaning the 'waste' or 'used' air, water and waste gases etc. of toxic or acidic materials e.g.

    • Toxic carbon monoxide from the blast furnace extraction of iron, it can be burnt as a fuel, but it must not be released into the air unless converted to biologically harmless carbon dioxide.

    • Sulphur dioxide gas from copper extraction of its sulphide ore is an irritating poisonous gas which can also cause acid rain, but it can be converted to the useful, therefore saleable, industrial chemical concentrated sulphuric acid, so you can remove a harmful pollutant and recover back some of the metal extraction costs, good green economics? Acidic gases like sulphur dioxide 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'.

  • Mining operations will disfigure the landscape BUT it can be re-claimed and 'landscaped' in an attempt to restore the original flora and fauna. However in the case of a limestone quarry, I'm afraid there is no way round the fact that huge chunks of beautiful hills get carted away if we want to use it as useful mineral.

• There is also an 'economics' section on the "Extra Industrial Chemistry" page too.

• The cost of extracting and purifying metals is quite varied for several reasons.

  • If the ore is plentiful it is cheaper e.g. iron ore, but silver ores and gold are much rarer and on that basis alone they would be a more valuable commodity.

  • Reduction of ores using coke (e.g. iron), made from cheap coal, is cheaper than the electricity bill for extracting aluminium from its molten oxide by electrolysis, but different metals have different properties best suited for particular and different uses.

  • Generally speaking, more reactive metals (like Al) are more costly to extract than less reactive metals (like Fe) because of the different energy demands and ease of extraction, which may sometimes be due to more costly technology.

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