METALS: EXTRACTION PROPERTIES AND USES - Form 4 Chemistry notes

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Introduction

  • Only most unreactive metals occur naturally in their elementary form.
    Examples: - Gold, Silver, Platinum.
  • Other elements occur as ores i.e. metal bearing rocks.
    Examples:
    • Oxides
    • Sulphides
    • Carbonates
    • Chlorides.
  • Note:
    An ore is a mineral deposit with reasonable composition of a desired metal.

Methods of Extraction

  • Depend on position of the metal in the reactivity series.
  • Main methods are:
    1. Electrolytic Method:
      • Used for metals high up in the reactivity series
        E.g.
        - Sodium and Potassium
        - Calcium and Magnesium
        - Aluminium.
      • These metals occur in very stable ores
    2. Reduction method:
      • For less reactive metals.
        E.g. Iron, Zinc, and Copper.
      • Is achieved using;
        1. Carbon in form of coke.
        2. Carbon (II) oxide
        3. Hydrogen
      • Oxidation is also used followed by reduction.

Preliminary Steps Before Extraction

  • Minerals (mineral) are usually mined with several impurities which lower the concentration of the metal per given mass or volume.
  •  Thus the ore is first concentrated before the actual extraction.
  • Concentration is possible due to difference in properties between the mineral compound and the earthy materials.

Methods of Ore Concentration

  1. Physical methods.
    1. Optical sorting.
      • Used to separate ore particles that have sufficiently different colours to be detected by the naked eye.
      • It involves physical handpicking of the desired particles.
      • Mainly used for minerals containing transition elements such as chromium.
    2. Hydraulic washing.
      • Also called sink and float separation.
      • Utilizes the difference in density between the minerals and the unwanted materials
      • The ore is washed with streams of water.
      • The denser ore particles will sink to the bottom of the washing container and can then be collected.
      • Examples in ores of tin and lead.
    3. Magnetic separation.
      • Is used when either the ore particles or the earthy materials (unwanted materials) are magnetic
      • A strong magnet is used to attract the magnetic components and leaving the non-magnetic materials behind.
      • Examples: in ores like magnetite (Fe3O4) and chromite which are magnetic.
    4. Electrostatic separation.
      • Used to separate particles which have different electric charges.
      • The particles are subjected into an electric field.
      • The oppositely charged particles follow different paths and can then be separated.
    5. Froth floatation
      • Is mainly use for sulphide ores.
      • Takes advantage of two facts.
        • Oil can wet the surfaces of ores.
        • Oil floats on water
      • The process:
        • The ore is ground into a fine powder; to increase the surface area for upcoming reactions.
        • It is then mixed with water and a suitable oil detergent e.g pine or eucalyptus;
        • The mixture is then agitated by blowing compressed air through it;
        • Small air bubbles attach to the oiled ore particles; which are thenn buoyed up and carried to the surface where they float.
        • A froth rich in mineral is formed at the top while impurities sink at the bottom.
        • The froth is skimmed off and dried.
        • Froth floatation process is used for copper, lead and zinc metals;
          froth floatation apparatus
          Diagram: Froth Floatation Apparatus
  2. Chemical concentration.
    • Involves the use of chemical reactions to concentrate the ores.
      Examples:
    • Bauxite, the main aluminium ore is chemically concentrated by a process known as Bayer s process.
    • This takes advantage of the amphoteric nature of aluminium oxide, which can thus react with both acids and bases.
    • Chemical concentration can also be done by leaching.
    • This involves reacting the ore with a compound such as sodium cyanide;
    • The cyanide ions form complex ions with the metal.
    • The complex ions formed are water soluble, and can be separated by filtration, leaving the unwanted materials in the residue.

The Metals

Sodium

  • Main ores;
    1. Rock salt / sodium chloride; NaC
    2. Chile saltpetre / sodium Nitrate; NaNO3
    3. Soda ash/sodium carbonate; Na2CO3
  • Other ores include;
    1. Borax; Na2B4O7.10H2O
    2. Sodium Sulphate, Na2SO4;

Extraction

  • Sodium is obtained by the electrolysis of fused sodium chloride in the electrolytic cell.
  • Calcium chloride and calcium fluoride are added to the electrolyte.
    Reasons;
    To lower the melting point of sodium chloride from 800oC to 600oC;
  • Once molten, the electrical resistance within the cell is sufficient to maintain the temperature without external heating.
  • Steel or iron is used as the cathode , while carbon/graphite is used as the anode .
  • Thus steel is not used as the anode.
    Reason;
    At high temperatures, steel would react with chloride formed at the anode, but graphite is inert even at high temperatures.
  • Steel wire gauze separates the electrodes.
    Reason;
    To prevent products of electrolysis (sodium and chlorine) from mixing and reacting to form sodium chloride.
  • The electrolytic apparatus used in sodium extraction is called the Downs cell.
    downs cell
    Diagram: The Downs Cell.
  • During electrolysis, fused sodium chloride dissociates according to the equation;
    NaCl(l) → Na+(l) +Cl-(l)

    At the cathode:
    Observation;

  • Soft silvery metal
    Explanation
  • Na+ ions are attracted and undergo reduction (accept electrons) to form/ produce molten sodium metal.
    Equation;
    Na+(l) + e- → Na(l)
  • Molten sodium is lighter than fused sodium chloride and floats on the surface where it overflows into a separate container/sodium reservoir.
    Note;
  • The resultant sodium is usually collected in liquid/molten state, floating on top of the electrolyte.
    Reasons;
    - Less dense than molten sodium chloride
    - Has a low melting point.

    At the anode;
    Observations;

  • Evolution of a green-yellow gas.
  • Chlorine gas is evolved as a by product and collected separately.
  • Negatively charged Cl - ions migrate to the positive anode and undergo oxidation to form chlorine gas
    Equation:
    2Cl-(l) → Cl2(g) + 2e-

Properties of Sodium

Physical properties

  • Is a soft silvery metal with low density; 0.979gcm-3
  • Has a low melting point, 97oC, and a low boiling point of 883oC

Chemical reactions

  1. With air
    • Na is very reactive and tarnishes in moist air to form an oxide layer.
      4Na(s) + O2(g) → 2Na2O(s);
    • The oxide layer reacts with more air moisture to form hydroxide
      Na2O(s) + CO2(g) → Na2CO3(s) + H2O(l)
      Note
    • Due to those series of reactions sodium is stored under a liquid hydrocarbon e.g. petroleum, kerosene
    • Sodium burns in oxygen with a golden yellow flame to form sodium peroxide
      Equation:
      2Na(s) + O2(g) → Na2O2(s) (White)
  2. With water
    • Na reacts vigorously with water to form NaOH and Hydrogen.
      Equation:
      2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
    • The resulting solution is highly alkaline with a PH of 14.
    • Sodium is stored under oil to prevent contact with moisture from the atmosphere.
      Note:
    • The reaction between Na and dilute acids would be explosive and not safe to investigate
  3. With chlorine
    • Sodium burns in chlorine gas;
      2Na(s) + Cl2(g) → 2NaCl(s)
  4. With ammonia gas;
    Sodium forms hydrogen and a solid;
    2Na(s) + NH3(g) → 2NaNH2(s) (Sodamide) + H2(g)
    And;
    NaNH2(s) + H2O(l) → NaOH(aq) + NH3(g)


Uses of Sodium

  1. Is alloyed with lead in the preparation tetraethyl (IV) lead, which is added to petrol as an anti-knock.
  2. Provides the glow in sodium vapours lamps, for street lighting (orange-yellow street lights).
  3. Is an excellent conductor of heat and electricity with low melting point hence used;
    • In nuclear reactors to conduct away heat.
    • Modern aeroplane engines.
  4. Manufacture of sodium peroxide, and sodium cyanide used in the extraction of silver and gold.




    Question;
  1. Although electrolysis is an expensive way of obtaining metals, it must be used for some metals. Explain.
    Answer
    • Group 1 and 2 metals together with Al are themselves such powerful reducing agents that their oxides cannot be reduced by chemical reducing agents.

    Worked example
  2. Below is a simplified diagram of the Downs cell in which sodium metal is manufactured.
    worked example sodium
    1.  
      1. Identify; -
        • Electrolyte Y: - Molten sodium chloride
        • Gas X: -Chlorine gas
      2. Write an equation for the reaction at the cathode.
        Na+(l) + e- → Na(l).
      3. In what state is sodium collected?
        • Molten state/liquid state.
      4. Give two properties of Na that makes it possible to be collected as in (b) (iii) above.
        • Its less dense than molten sodium chloride.
        • Has a low melting point.
      5. The cathode is made of steel but the anode is made of graphite. Why is this yet steel is a better conductor?
        • At high temperature steel would react with chlorine formed but graphite is inert even at high temperatures.
      6. In this process, the naturally occurring, raw material is usually mixed with another compound. Identify the compound and state its use.
        • Compound; - Calcium chloride
        • Use; -To lower melting point of NaCl2 from 800oC to 600oC
      7. What is the function of the steel gauze cylinder?
        • Prevents sodium reacting with chlorine forming NaCl
      8. Give one industrial use of sodium
        • A coolant in nuclear reactors;
        • Alloy with lead in tetraethyl (IV) lead;
      9. Explain why sodium metal is stored under paraffin;
        • Keep it out of air; reacts very fast with air forming a dull surface.
        • Can react with water.
    2.  
      1. State an industry that can be built next to a sodium extracting plant.
      2. A current of 100 Amperes flows through an electrolyte of molten sodium chloride for 15 hours. Calculate the mass of sodium produced in kg (Na = 23; 1F = 96500 C)
        Solution:
        Q = It
        =100 x 15 x 60 x 60
        =5400000C.
        Cathode equation
        Na+(l) + e- → Na(l)
        96500C = 23g of Na
        5400000 C = 23 x 5400000/96500 =1287.04g
        =1.287kg
      3. For the same quantity of electricity as in (c) above ; calculate the volume of the gaseous product produced in the cell at 15oC and 800mmHg.(Molar gas volume at s.t.p = 22.4dm3)
  3. The diagram below shows the extraction of sodium metal using the Downs cell. Study it and answer the questions that follow:
    downs cell
    1. Explain why in this process the sodium chloride is mixed with calcium chloride.
    2. Why is the anode made of graphite and not steel?
    3. State the properties of sodium metal that makes it possible for it to be collected as shown in the diagram.
    4. What is the function of the steel guaze cylinder?

Aluminium

  • Forms 7% of the earth s crust and is the most common metal

Main Ores

  • Bauxite; Al2O3.H2O
  • Mica; K2Al2S6O16.
  • China clay;Al2S2O7.2H2O
  • Corundum;Al2O3

Chemical Test

  • Crush the ore into a fine powder;
  • Add dilute nitric (V) acid to the powder
  • Filter to obtain a solution of the ore;
  • To a solution of the ore add NaOH(aq) dropwise till in excess, and then repeat the same procedure using Ammonia solution, NH4OH.
    Observations:
    With NaOH(aq):
  • White precipitate in soluble in excess;

    With NH4OH(aq):
  • White precipitate insoluble in excess;

Extraction from Bauxite

  • Involves two main processes;-
    • Purification of Bauxite.
    • Electrolysis of purified bauxite (alumina

    Purification of Bauxite
  • Chief impurities are small quantities of silica and iron (III) oxide.
  • The oxide ore is ground and treated under pressure/ dissolved in hot aqueous sodium hydroxide.

    During the process;
  • The amphoteric bauxite dissolves in NaOH forming sodium aluminate;
    Equation:
    2NaOH(aq) + Al2O3.3H2O(s) → 2NaAl(OH)4(aq)
    Ionically:
    Al2O3(s) + 2OH-(aq) + 3H2O(l) → 2[Al(OH)4]-(aq);
  • Silica impurities also dissolve forming sodium silicate
    Equation:
    SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l)
  • The iron impurities (mainly iron (III) oxide) DO NOT dissolve.
  • This mixture is then filtered, during which iron (III) oxide remain as residue of red mud while a filterate of sodium aluminate and sodium silicate is collected.
  • Carbon (IV) oxide is bubbled through the filterate, followed by dilution then addition of a little aluminium hydroxide to cause precipitation (seeding) of Aluminium hydroxide.
    Ionically:
    2[Al(OH)4]-(aq) + CO2(g) → 2Al(OH)3(s) + H2O(l);
    Alternatively:
    Al(OH)4-(aq) → 2Al(OH)3(s) + OH-(aq);
    General equation:     hydrolysis
    NaAlO2(aq) + 2H2O(l)          →      NaOH(aq) + Al(OH)3(s)
  • The precipitated Aluminium hydroxide is then filtered off, washed and ignited to give pure aluminium oxide (Alumina);
    Equation:
    2Al(OH)3(s) → Al2O3(s) (Alumina) + 3H2O(l)

    Electrolysis of Purified Bauxite (alumina)
  • The Alumina (Al2O3), has a high melting point, 2015o C and a lot of heat would be required to melt it.
  • Additionally the molten compound is a very poor conductor of electricity.
  • Consequently, cryolite (Na3AlF6) is mixed with the oxide.
    Reason;
    To lower the melting temperature of Al2O3 from 2015oC to around 800oC;
  • At this lower temperature the molten oxide also conducts well.
  • The molten alumina mixed with bauxite is then electrolysed in a steel cell lined with carbon graphite as the cathode.
    Note;
  • Other than being an electrolyte the graphite cathode lining also prevents alloy formation, as it ensures no contact between the resultant aluminium and the steel tank;
  • The anodes also made of Graphite dip into the steel tank at intervals.
    electrolytic steel cell for the extraction of aluminium
    Diagram: electrolytic steel cell for the extraction of Aluminium.

    Electrolytic reactions;
  • The Aluminium oxide dissociates to give constituent ions;
    Equation:
    Al2O3(l) → 2Al3+(l) +3O2-(l)
    At the cathode;
    Observation;
  • A silvery white metal which quickly becomes dulled.
    Explanation:
  • Aluminium ions move to the cathode and are reduced to form aluminium metal.
    Equation;
    2Al3+(l) + 6e→ 2Al(l)
    At the anode;
    Observation;
  • Effervescence of a colourless gas.
    Explanations:
  • Oxygen ions migrate to the anode and get oxidized to form oxygen gas.
  • The resultant oxygen gas reacts further with the graphite anode to form carbon (IV) oxide.
  • This is due to the high temperatures involved during the process.
    Note;
  • Consequently the carbon anode should be replaced from time to time.
    Equations
    3O2-(l) → 3O2(g) + 6e-
    Then;
    C(s) (anode) + O2(g) → CO2(g)

    Note:
  • Cryolite usually adds Na+ ; and F- ions into the electrolyte.
  • Thus the anions are O2- and F- ions into the electrolyte.
  • However oxygen is discharged in preference to Fluorine.
    Reason;
  • Fluorine is a stronger oxidizing agent than oxygen. Thus oxygen easily gives electrons than fluorine, hence discharge.
  • Aluminium is discharged in preference to sodium.

Summary: - Flow Chart on the Extraction of Aluminium from Bauxite .
flowchart for the extraction of aluminium

Properties of Aluminium

Physical properties

  • Is a silvery white metal which quickly becomes dulled with a thin oxide layer.
  • Has very low density (2.7gcm-3), with ability to be rolled into wires/foil.
  • Is a good conductor of heat and electricity.


Chemical properties.

Reaction with Air;

  • In air it acquires a continuous very thin coating of oxide, which resists further reaction.
  • Removal of this protective cover renders the metal reactive.
  • Consequently steel wool or wood ash should NOT be used in aluminium utensils.
  • Usually, salty water attacks the oxide film allowing the aluminium to corrode and for this reason, ordinary aluminium is not used for marine purposes .
  • Aluminium will burn in air at 800 o C to form is oxide and nitrate.
    Equations:
    4Al(s) + 3O2(g) → 2Al2O3(s);
    2Al(s) + N2(g) → 2AlN(s)

Reaction with Acids.

  • Note:
    The protective Aluminium oxide (being covalent and insoluble) layer makes its reactivity with acids less than expected.

    With nitric (V) acid;
  • Has hardly any effect on the metal, at any concentration.
    Reason:
  • Being a powerful oxidizing agent, it simply thickens the oxide layer thereby preventing further reaction.

    With sulphuric (VI) acid;
  • Only hot concentrated sulphuric (VI) acid breaks down the oxide layer and reacts with the metal.
    Equation:
    2Al(s) + 6H2SO4(l) → Al2(SO4)3(aq) + 6H2O(l) + 3SO2(g)

    With Hydrochloric acid;
  • Dilute HCl dissolves aluminium slowly; liberating hydrogen.
    Equation:
    2Al(s) + 6HCl(l) → 2AlCl3(aq) + 3H2(g)
  • With concentrated HCl the rate of reaction is increased.


Reaction with Chlorine;

  • Hot aluminium burns in chlorine gas with a white light, forming dense white fumes of Aluminium (III) Chloride.
  • The white fumes cool and collect on the cooler parts of the apparatus as a white solid.
    Equation:
    2Al(s) + 3Cl2(g) → 2AlCl3(g)
    Note:
  • The apparatus for the preparation of AlCl3 is kept very dry.
    Reason:
  • Aluminium chloride is readily/easily hydrolysed by water/moisture, and so it fumes in damp air with the evolution of hydrogen chloride gas.
    Equation:
    AlCl3(s) + 3H2O(l) → Al(OH)3(s) +3HCl(g)

Reaction with Water

  • Aluminium does not react with cold water, due to the formation of an insoluble coating of Aluminium oxide.
    Note;
  • If the oxide film is removed, the metal reacts slowly with cold water.

Reaction with Caustic Soda.

  • The metal, especially in powder form, reacts with caustic soda solution, liberating hydrogen and leaving sodium aluminate in solution.
  • The reaction is exothermic and once started, it is very vigorous.
    Equation:
    2NaOH(aq) + 2Al(s) + 2H2O(s) → 2NaAlO2(aq) + 3H2(g)
    Ionically:
    2Al(s) + 2OH-(aq) + 2H2O(l) → 2AlO-2 (g) + 3H2(g)
    Note:
  • Thus aluminium has an amphoteric nature as it reacts with both acids and alkalis.

Uses of Aluminium

  1. Making parts of airplanes, railway, trucks, buses, tankers, furniture, and car e.t.c.
    Reason;
    - It is Very light due to a very low density.
  2. Making cooking vessels/utensils such as sufurias.
    Reason:
    - It is a good conductor of heat and electricity.
    - It is not easily corroded by cooking liquids due to the unreactive coating of aluminium oxide.
  3. Making overhead cables
    Reason:
    - It is a good conductor of electricity.
    - It is light hence can easily be supported by poles and ductile to be rolled into wires (cables).
  4. Aluminium powder mixed with oil is used as a protective paint.
  5. Making Aluminium foils due to its high malleability. The foil is used in cooking; packaging and for milk bottle tops.
  6. Making alloys, which have high tensile strength and yet light.
    Examples:
     Alloy  Component
     Duralumin  Aluminium, copper, manganese and magnesium
     Magnalium  Aluminium (70%) and magnesium (30%)
  7. As a reducing agent in the Thermite process in the production of some elements such as chromium, cobalt manganese and titanium.
    Example:
    Cr2O3(s) + 2Al(s) → 2Cr(s) + Al2O3(s)
    Note:
    The thermite process
    • Is a process of reducing oxides of metals which are ordinarily difficult to reduce using Aluminium powder.
    • Examples:
      • Iron (III) Oxide (Fe2O3)
      • Chromium (III) Oxide (Cr2O3)
      • Compound oxide of manganese (Mn3O4)
    • The oxide and the Al powder are well mixed together, forming Thermite.
    • The thermite is ignited using magnesium ribbon fuse, since the reaction will not start at low temperatures.
    • The high heat of formation of Aluminium oxide , results into a vigorous exothermic reaction that leads to a molten metal.
      Example: -In the reduction of chromium (III) Oxide.
      Cr2O3(s) + 2Al(s) → 2Cr(l) + Al2O3(s) + Heat

Sample Question

  1. The extraction of aluminium from its ore takes place in two stages, purification stage and electrolysis stage. The diagram below shows the set up for the electrolysis stage.
    aluminium sample question
    1. Name the ore from which aluminium extracted 
    2. Name one impurity which is removed at the purification stage

Zinc

  • Main ores;
    1. Zinc blende; ZnS
    2. Calamine; ZnCO3 .

Qualitative Analysis/Test for Presence of Zn2+ in an Ore Sample

  • The ore is crushed and then dilute nitric or hydrochloric or sulphuric acid added to dissolve the ore.
  • It is then filtered to obtain Zn2+ filtrate.
  • The filtrate is divided into 2 different test tubes.
  • To the first portion sodium hydroxide solution is added dropwise till in excess, formation of a white precipitate soluble in excess confirms presence of either Zn2+ ; Al3+ or Pb2+ .
  • To the second sample aqueous ammonium hydroxide is added dropwise till in excess; formation of a white precipitate soluble in excess NH4OH(aq) confirms presence of Zn2+ only;
    Equations

    With little NH4OH:-
    Zn2+(aq) + 2OH-(aq) → Zn(OH)2(s) (White ppt)

    In excess:
    Zn(OH)2(aq) + 4NH3(aq) → [Zn(NH)4]2+(aq) (Colourless solution) + 2OH-(aq)

Extraction

  • Is done by electrolysis or reduction of its oxide using carbon.


    Preliminary steps:
  • The ore is first concentrated by froth floatation.
  • The ore is roasted in air to convert it to the oxide.

    Equations:
    From Zinc blende:
    2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g)

    From Calamine:
    ZnCO3(s) → ZnO(s) + CO2(g)
  • After obtaining the oxide the metal is extracted by either reduction or electrolysis:
    1. The reduction method.
      • The oxide is mixed with coke and limestone and heated in a furnace.
        furnace for zinc extraction
        Diagram: Furnace for zinc extraction by reduction:
      • The limestone (CaCO3) decomposes to liberate CO2 which is then reduced by coke to form carbon (II) oxide.
        Equations:
        CaCO3(s) → CaO(s) + CO2(g)
        Then:
        CO2(g) + C(s) → 2CO(g)
      • The resultant carbon (II) oxide and coke are the reducing agents in the furnace, at about 1400 o C.
      • They reduce the oxide to the metal; which is liberated in vapour form.
        Equations:
        ZnO(s) + C(s) → Zn(g) + CO(g)
        ZnO(s) + CO(s) → Zn(g) + CO2(g)
      • At the furnace temperatures zinc exists in vapour form, and leaves at the top of the furnace.
      • Liquid zinc being lighter settles above molten lead and is run off;
      • The vapour is condensed in a spray of molten lead to prevent re-oxidation of zinc.
      • The resultant zinc is 98-99% pure and can be further purified by distillation.
      • SO2 is a by-product and is the main source of pollution in the extraction of zinc.
      • Usually it is channeled to a contact process plant for the manufacture of sulphuric acid.
      • Alternatively it can be scrubbed off to prevent pollution of the environment.
      • Less volatile impurities remain in the furnace.
      • The silica impurities combine with the quicklime/ calcium oxide (CaO) from limestone to form calcium silicate.
        Equation:
        CaO(s) + SiO(s) → CaSiO3(s)
      • The silicates together with other less volatile impurities form slag , at the bottom of the furnace from where it is run off.

        Summary: Flow chart and the extraction of zinc
        flowchart for the extraction of zinc by electrolysis
    2.  Electrolytic extraction of zinc.
      • Zinc metal is obtained from the oxide via a series of steps:

        Step I: Preparation of electrolyte:
      • The ZnO obtained from roasting the ore is converted to zinc sulphate by reacting it with dilute sulphuric (VI) acid.
        Equation:
        ZnO(s) + H2SO4(aq) → ZnSO4(aq) + H2O(l)
      • Any lead (II) oxide impurity present in the zinc oxide reacts with the acid to form lead (II) sulphate.
        Equation:
        PbO(s) + H2SO4(aq) → PbSO4(s) + H2O(l)
      • The insoluble lead (II) sulphate is then precipitated and separated by filtration;
      • The zinc sulphate is then dissolved in water and the solution electrolysed.




        Step II: The electrolytic process:
        Electrolyte:
      • Zinc (II) sulphate solution;

        Ions present:
      • Zn2+ and H+ as cations; and SO42- and OH- as anions;

        Cathode:
      • Lead containing 1% silver.

        Anode:
      • Aluminium sheets;

        Chemical reactions:
        Cathode:
        Observations:
      • Deposits of a grey solid.

        Explanations:
      • Zn2+ and H+ migrate to the cathode.
      • The Zn2+ are discharged in preference to H+;
        Reason:
      • The cathode is relatively reactive. Thus since zinc is more reactive thn hydrogen, its ions undergo reduction faster;
        Equation:
        Zn2+(aq) + 2e- → Zn(s);
        Note:
      • If graphite electrodes were used, hydrogen gas would have been evolved instead;

        Anode:
        Observations:
      • Evolution of a colourless gas that relights a glowing splint;

        Explanations:
      • OH- and SO42- migrate to the cathode.
      • The OH- are discharged in preference to SO42- ; giving off oxygen gas
        Reason:
      • The OH- ions have a higher oxidation potential than SO42- and therefore easily giving electrons for reduction at the cathode
        Equation:
        4OH-(aq) → 2H2O(l) + O2(g) + 4e-
        Note:
      • Over 80% of zinc is extracted by the electrolytic method.
      • Zinc extracted by the electrolytic method is much more pure.
        Note: - Industrial plants that can be set up near the zinc extraction plant.
      • Contact process plant, to make use of the SO2 by-product.
      • Lead accumulators factories, to utilize the zinc produced.
      • Paper factory using, SO3 and hence SO2 in bleaching.
      • Brass factory for alloying zinc and copper.
      • Steel factory to use zinc in galvanization.

Properties of Zinc

Physical Properties

  • Is a blue-grey lustrous metal with:
    • Density of 7.1gcm-3
    • Melting point of 420o C
    • Boiling point of 907oC

Chemical reactions

  1. With air;-
  • Zinc tarnishes slowly forming a protective layer which prevents further reaction i.e. oxide layer or basic carbonate.
    Equations:
    2Zn(s) + O2(g) → 2ZnO(s)
    Then;
    ZnO(s) + CO2(g) → ZnCO3(s)
  • It burns with a blue-green flame when strongly heated in air to form an oxide which is yellow when hot and white when cold.
    Equation:
    2Zn(s) + O2(g) → 2ZnO(s)
  1. With water
  • Zinc does not react with water
  • Steam reacts with red-hot zinc, forming zinc oxide and liberating hydrogen gas.
    Equation:
    Zn(s) + H2O(g) → ZnO(s) + H2(g)
  1. With dilute acids
  • Zinc is above hydrogen in the reactivity series hence displaces hydrogen from steam (water) and dilute acids like H2SO4 and HCl.
    Equation:
    Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)
  • Pure zinc reacts slowly while impure zinc reacts faster/more quickly.
  • Copper (II) sulphate is used as a catalyst to speed up the reaction.
  1. With Concentrated acids.

    Concentrated sulphuric (VI) acid.
  • Equation:
    Zn(s) + 2H2SO4(l) → ZnSO4(aq) + 2H2O(l) + SO2(g)

    50% concentrated nitric (v) acid.
  • It reacts with 50% concentrated nitric (V) acid to liberate nitrogen (II) oxide.
    Equation:
    3Zn(s) + 8HNO3(l) → Zn(NO3)2(aq) + 2H2O(l) + 2NO(g)


    Concentrated nitric (V) acid.
  • It reduces concentrated nitric (V) acid to nitrogen (IV) oxide.
    Equation:
    Zn(s) + 4HNO3(l) → Zn(NO3)2(aq) + 4H2O(l) + 2NO2(g)
  1. With Alkalis
  • Zinc is amphoteric and dissolves in hot alkalis to give the zincate ion and hydrogen gas.
    Equation:
    Zn(s) + 2OH-(aq) + 2H2O(l) → H2(g) + [Zn(OH)4]2-(aq) (Zincate ion)
  1. Other reactions.
    1. Zinc burns in chlorine to give zinc chloride
      Zn(s) + Cl2(g) → ZnCl2(s)
    2. Zinc combines with sulphur
      Zn(s) + S(s) → 2ZnS(s)

Uses of Zinc

  1. Galvanization of iron sheets to prevent corrosion and rusting.
    Note:
    Rusting does not occur even when galvanized iron sheets are scratched and exposed.
    Reason:
    - The rest of the zinc protects the iron from rusting. This is because zinc being more reactive gets oxidized in preference to iron, and is hence sacrificed in the protection of iron.
    - This method is referred to as cathodic or sacrificial protection.
  2. Making brass; an alloy of copper and zinc.
  3. Making outer casings of dry batteries;
  4. Die-castings contain zinc and aluminium, and are used for making radio and car parts;
  5. Zinc cyanide is used for refining silver and gold;

Sample Question

  1. Study the flow chart below and answer the questions that follow.
    zinc sample question
    1. State the condition necessary for the reaction in step 2 to occur. 
    2. Name:
      1. Gas D
      2. One use of zinc

Iron

  • Is thesecond most abundant metal after aluminium, forming about 7% of the earth s crust.

    Main ores
  • Haematite, Fe2O3;
  • Magnetite, Fe3O4;
  • Siderite, FeCO3;

Qualitative Analysis for Presence Copper in an Ore Sample.

  • Crush the ore into fine powder;
  • Add dilute nitric (V) acid to the ore, to dissolve the oxide filter to obtain the filtrate.
  • To the filtrate add aqueous sodium hydroxide / ammonium hydroxide dropwise till in excess.
  • Formation of a red brown / brown precipitate in both cases indicates presence of Fe3+

Extraction from Haematite (Fe2O3)

Summary of the Process

  • The ore-haematite is crushed and mixed with coke and limestone.
  • The mixture is called charge.
  • The charge is loaded into the top of a tall furnace called blast furnace .
  • Hot air-the blast is pumped into the lower part of the furnace.
  • The ore is reduced to iron as the charge falls through the furnace.
  • A waste material called slag is formed at the same time.
  • The slag floats on the surface of the liquid iron produced.
  • Each layer can be tapped off separately.

Details of the Extraction Process.

The Blast Furnace

  • Is a tall, somewhat conical furnace usually made of silica and lined on the inside with firebrick .
    blast furnace for iron manufacture
    Raw materials.
  • Iron ore; i.e. haematite
  • Coke; C
  • Limestone; CaCO3 ;
  • Hot air;
    Conditions
  • Temperature at the bottom of furnace, 1400-1600oC
  • Temperature at the top of the furnace, 400oC

    Reactions and Processes
    Step-1 crushing and loading
  • The ore is crushed into powder form, to increase the surface area for the upcoming reduction/ redox reactions.
  • It is then mixed with coke and limestone and then fed at the top of the furnace using the double bell (double-cone devise) changing system
    Note :
  • The double bell charging system ensures that the furnace can be fed continuously from the top with very little heat loss, by preventing any escape of hot gases.
  • This in turn reduces production costs.

    Step 2: -Pre heating of the blast furnace.
  • Air that has been preheated to about 700 o C is blown/ fed into the base of the blast furnace through small pipes called tuyers.
  • This provides the required temperatures for the reactions in the blast furnace.
  • This results into highest temperatures, about 1600 o C at the hearth (bottom of the furnace) which then decreases upwards the furnace.

    Step 3: -Generation of reducing agents.
  • Two reducing agents are used in this process: Coke and carbon (II) oxide; with carbon (II) oxide being the main reducing agent.
    1. Oxidation of coke;
      • Coke burns in the blast at the bottom of the furnace.
      • The reaction temperatures is about 1600 o C and the product is Carbon (IV) oxide gas
      • This reaction is exothermic, producing a lot of heat in the blast furnace.
        Equation:
        C(s) + O2(g) → CO2(g)
    2.  Decomposition of limestone;
      • The limestone in the charge decomposes in the blast furnace to calcium oxide (Quicklime) and carbon dioxide.
        Equation:
        CaCO3(s) → CaO(s) + CO2(g)
      • The calcium oxide will be used in the removal of the main ore impurity/ silicates/ silica in the form of silicon (IV) oxide.
      • The CO2 then moves up the blast furnace to regenerate carbon (II) oxide, the chief reducing agent.
    3. Production of carbon monoxide
      • The CO 2 from oxidation of coke and decomposition of limestone (calcium carbonate) react with (excess) coke, to form carbon (II) oxide
      • The reaction occurs higher up in the blast furnace at about 700oC;
        Equation:
        CO2(g) + C(s) → 2CO(g)

    Step 4: The actual reduction process
  • Reduction of the ore is by either CO or coke, depending on temperatures.
    1. Reduction by coke
      • This occurs much lower down the furnace at higher temperatures of about 800oC and above .
      • This reaction is ordinarily slow and thus serves to only reduce the part of the ore reduced by CO at lower temperatures in the upper parts of the furnace.
        Equation:
        2Fe2O3(s) + 3C(s) → 4Fe(s) + CO2(g)
        Note:
      • The resultant CO2 is quickly reduced to CO by the white-hot coke to more carbon (II) oxide as per step 3(iii) above
    2. Reduction by carbon (II) oxide
      • This is the main reducing agent.
      • The reaction between CO and Fe2O3 is relatively faster and occurs at lower temperatures of 500o C-700o C, higher up the furnace.
        Equation:
        Fe2O3(s) + 3CO(g) → 2Fe(s) + CO2(g)
      • The resultant carbon (IV) oxide is also quickly recycled by being reduced to CO by coke to from more reducing agent
    3. Melting
      • The iron produced in both of the reduction processes is in solid state.
      • As the iron drops / falls down the furnace, it melts as it passes through the melting zone/ molten zone (1500o C-1800o C)
      • The molten iron runs to the bottom of the furnace.
      • Temperatures at the hearth (bottom of the furnace) is maintained at approx. 1400o C and yet pure iron melts at about 1525o C.
      • Consequently the molten iron would easily solidify at the base (Temp =1400o C)
      • However this is not usually the case;
        Reason:
      • Impurities absorbed by iron during melting ( mainly carbon ) reducing the melting point to below 1400o C.
      • The molten iron is then easily tapped off.

    Step 5: -Removal of earthy impurities.
  • The earthy impurities in the ore (mainly silica) react with calcium oxide from decomposition of limestone to form calcium silicate.
    Equation:
    CaO(s) + SiO2(s) → CaSiO3(s)
  • These earthy impurities form molten slag whose main component is calcium silicate .
  • The slag does not mix with iron but rather floats on top of it, at the base of the furnace.

    Importance of the slag
  • As it floats on top of molten iron it protects it from being re-oxidized by the incoming hot air.

    Uses/application of the slag
  • Light-weight building material.
  • Manufacture of cement.
  • Road building material.

    Step 6:- Removal of furnace (waste) gases.
  • Hot unreacted/waste gases leave at the top of the furnace.
  • Main components include Nitrogen, unreacted CO2, unreacted CO, oxygen and Argon (Noble gases)
  • Additionally they contain dust particles.
    Note:
  • Upon removal of dust particles, the furnace gases, being hot can be used to pre-heat the air blown in at the base.

Properties of Iron:

Physical properties:

  • It has a melting point of 420oC and a boiling point of 907o C;
  • Have a good thermal and electrical conductivity;
  • It is ductile and malleable;

Chemical properties.

  1. Reaction with air.
    • It readily rusts in presence of moist air hydrated brown iron (III) oxide; Fe2O3.H2O(s)
      Equation:
      4Fe(s) + 2H2O(l) + 3O2(g) → 2Fe2O3.H2O(l)
    • When heated it reacts with oxygen to form tri-iron tetroxide; Fe 3 O 4 ;
      Equation:
      3Fe(s) + 2O2(g) → Fe3O4(s)
  2. Reaction with water.
    • It does not readily react with cold water.
    • It however reacts with steam liberating hydrogen gas and forming tri-iron tetroxide.
      Equation:
      3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g)
  3. Reaction with chlorine.
    • Hot iron glows in chlorine without further heating, forming black crystals of iron (III) chloride;
    • Iron (III) chloride sublimes on heating and will thus collect on the cooler parts of the apparatus;
      Equation:
      2Fe(s) + 3Cl2(g) → 2FeCl3(s)
      Note:
    • Iron (III) chloride fumes when it is exposed to damp (moist) air;
      Reason:
    • It is readily hydrolysed by water with evolution of hydrogen chloride gas;
      Equation:
      FeCl3(s) + 3H2O(l) → Fe(OH)3(s) + 3HCl(g)
  4. Reaction with acids:
    Hydrochloric acid:
    • Iron reacts with hydrochloric acid to liberate hydrogen gas.
      Equation:
      2Fe(s) + HCl(aq) → FeCl2(aq) + H2(g)

    Sulphuric (VI) acid:
    • Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)
    • Note: With hot concentrated H 2 SO 4 ;
    • The iron reduces hot concentrated H 2 SO 4 to sulphur (IV) oxide and it is itself oxidized to iron (III) sulphate.
      Equation:
      2Fe(s) + 6H2SO4(l) → Fe2(SO4)3(aq) + 6H2O(l) + 3SO2(g)

    Nitric (V) acid.
    • Iron reacts with dilute nitric (V) acid to form nitrogen (IV) oxide and ammonia which then forms ammonium nitrate.
      Equation:
      10HNO3(aq) + 4Fe(s) → 4Fe(NO3)2(aq) + NH4NO3(aq) + 3H2O(l)
    • Warm dilute nitric (V) acid gives iron (II) nitrate.
    • Concentrated nitric (V) cid renders the iron unreactive.
      Reason:
    • Formation of iron oxide as a protective layer on the metal surface.
  5. Reaction with sulphur.
    • Iron when heated in sulphur forms iron (II) sulphide.
      Equation:
      Fe(s) + S(s) → FeS(s)

Uses of Iron

  • Iron exists in different types and alloys, depending on percentage composition of iron, and other elements.
  • Each type of alloy of iron has different uses depending on properties.
     Iron alloy or type  Properties  Uses
     Cast iron  - Refers to iron just after it has been produced in the blast furnace
     - Contains 3-5% carbon, 1% silicon and 2% phosphorus
    Disadvantage: very little hence easily breaks
    Advantage: It is extremely very hard

    - Making:
    ➢ Furnaces;
    ➢ Grates;
    ➢ Railings;
    ➢ Drainage pipes;
    ➢ Engine blocks;
    ➢ Iron boxes;
    Note: This is due to its very hard nature;
    - Manufacture of wrought iron and steel;

     Wrought  - Refers to cast iron with 0.1 % carbon
    - It is malleable hence can easily be moulded or welded
     - Making iron nails; horse shoe; agricultural implements like pangas;
    Note:
    Its use is declining due to increased use of mild steel
     Steel  - Are alloys whose main components is iron
    - Other components may be carbon; vanadium; manganese; tungsten; nickel and chromium
    Examples
    Mild steel
    - has about 0.3 % carbon, 99.75% iron
    Special steel
    - has a small percentage of carbon together with other small substances
    Stainless steel
    - Contains 74% iron, 18% chromium and 8% nickel
    Cobalt steel
    - Contains about 97.5% iron and 2.5% cobalt
    - Very tough and hard
    - Slightly magnetic

    - Mild steel is used for making:
    ➢ Nails; Car bodies;
    ➢ Railway lines; Ship bodies;
    ➢ Rods for reinforced concrete, pipes;
    Note: Advantage of mild steel: - It is easy to work on;
    - That with 10-12% chromium and some nickel is used to make: cutlery; sinks; vats;
    -Steel containing 5-18% tungsten is used for: making high speed cutting and drilling tools;
    - For making electromagnets;

Copper

  • Description: - A red brown metal.
  • Distribution: - Canada, USA, Zambia, and Tanzania.

Main Ores

  • Copper pyrites, CuFeS2 ;
  • Cuprite, CuO
  • Chalcocite, Cu2S
  • Malachite, CuCO3.Cu(OH)2

Qualitative Analysis/Test for Presence in an Ore Sample

  • Crush the ore and then add dilute nitric or hydrochloric or sulphuric acid to dissolve the ore.
  • Filter to obtain Cu2+ filtrate.
  • Divide filtrate into 2 different test tubes.
  • To one sample add aqueous Ammonium hydroxide dropwise till in excess formation of a pale blue precipitate soluble in excess NaOH to form a deep blue solution.
    Equations

    With little NaOH:-
    Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s) (Pale blue ppt)

    In excess:
    Cu(OH)2(aq) + 4NH3(aq) → [Cu(NH3)4]2+(aq)Deep blue solution + 2OH-(aq)
  • To the second portion add sodium hydroxide solution dropwise till in excess, formation of a pale blue precipitate insoluble in excess confirms presence of Cu2+.

Extraction - from Copper Pyrites

  1. Crushing the ore
    • The ore is crushed to increase the surface area for the succeeding chemical reactions.
    • The ore is then concentrated.
  2. Concentration of the ore.
    • The ore is concentrated by froth floatation .
    • The fine ore powder is mixed with water and oil, after which air is blown into the mixture, usually from below.
    • Bubbles of the air forms froth, resulting to concentration of the ore.
    • The lighter oil floats on top of the water, with the ore floating on top of the oil.
    • The denser water sediments the earthy impurities like soil particles.
    • The concentrated ore is then tapped off.
    • This process involves formation of an oil froth onto which the ore floats hence the name froth formation .
  3. First Roasting
    • The concentrated copper pyrite, CuFeS2 is then roasted in air to remove some of the sulphur impurities as sulphur (IV) oxide.
      Equation:
      2CuFeS2(s) + 4O2(g) → 3SO2(g) + 2FeO(s) + Cu2S(s);
      Note:
    • During 1 st roasting limestone and silica (SiO2) are added to the roasted ore and the mixture heated in the absence of air.

      Importance
    • Removal of iron impurities.
    • The iron (II) so formed during roasting is converted to iron (II) silicate, FeSiO3.
    • The iron (II) silicate constitutes the major portion/component of the slag .
      Equation:
      FeO(s) + SiO2(g) → FeSiO3(s)
    • The slag separates itself from the copper (I) sulphide.
    • The sulphur (IV) oxide escapes into the atmosphere and is the major pollutant in this process.

      Pollution control mechanisms.
    • Scrubbing the gas using calcium hydroxide;
      Equation:
      SO2(g) + Ca(OH)2(s) → CaSO3(s) + H2O(l)
    • Construction of a contact process nearby.
  4. Second Roasting.
    • The CuS is heated in a regulated supply of air where some of it is converted to Cu2O
      Equation:
      2Cu2S(s) + 3O2(g) → Cu2O(s) + 2SO2(g)
  5. Reduction of copper (II) oxide
    • Note: - Not all the Cu2S was oxidized to copper (I) oxide Cu2O.
    • The unreacted (unoxidized) Cu2S serves as the reducing agent in this step. i.e. The copper (I) oxide formed in step 4 is reduced to copper metal by the (unreacted) copper (I) sulphide
    • This is called blister copper .
      Equation:
      Cu2S(s) + 2Cu2O(s) → 6Cu(s) + SO2(g);
  6. Electrolysis
    • The copper metal from reduction in step 6 is impure and is thus purified by electrolysis.

      Main impurities
    • Traces of gold
    • Traces of silver
    • Iron
    • Sulphur

      Electrolytic apparatus
      Anode: Impure copper
      Cathode: Pure copper plates/ sheets;
      Electrolyte: Dilute copper (II) sulphate solution (containing Cu2+; H+; SO42- and OH-)
      electrolysis of copper
      Diagram of Electrolytic Apparatus.

      Electrolytic reactions;
      At the cathode.

      Observations:
    • Deposition of a brown solid.

      Explanations
    • The copper (II) ions, Cu2+ move to the cathode, where they accept electrons and undergo reduction.
    • Cations in the electrolyte are Cu2+ and H+ but Cu2+ are preferentially discharged due to their easy tendency to undergo reduction.
      Equation:
      Cu2+(aq) + 2e- → Cu(s);

      At the anode
      Observations:
    • Dissolution of the anode, hence the impure copper rod decreases in size.

      Explanation
    • Since the metal rod is dipped into a solution of its ions, the copper solid undergoes oxidation, losing electrons to form copper ions, Cu2+
    • Consequently as more copper ions, Cu2+ get reduced at the cathode; more are released by the dissolving anode.
      Equation:
      Cu(s) → Cu2+(aq) + 2e-
      Overall reaction
      Cu(s) + Cu2+(aq) → Cu2+(aq) + Cu(s)
    • The electrolytic product is 99.98% copper.
    • Traces of silver and gold collect as sludge at the bottom of the cell.
      Note :-To improve purity of the product of electrolysis, the following steps are advisable;
      1. Increase the dilution of the electrolyte/ use a very dilute electrolyte.
      2. Reduce the amount of current / use a low current.

 

Summary of Extraction of Copper from Copper Pyrites.
flowchart for the extraction of copper

Uses of Copper

  • Making copper wires and contacts in switches, plugs and sockets
    Reason: -Copper is a good conductor of electricity
    Note: -For this purpose, pure copper is necessary, since impurities increase electrical resistance.
  • Making soldering instruments.
    Reason: -Copper has a high thermal conductivity
  • Manufacture of alloys.
    Examples
     Alloy  Components
     Brass  Copper and zinc
     Bronze  Copper and tin
     German silver  Copper. zinc and nickel
  • Making coins and ornaments.
    Reason: -it is durable and aesthetic.

Properties of Copper

Physical properties

  • Soft red-brown metal.
  • Melting point of 1083oC and a boiling point of 2595oC
  • Density is 8.92gcm -3 and electrical conductivity is about 5.93 x 10-9 Ώ-1m-1

Chemical properties

  • It does not react with cold water or steam.
  • Heating in air
    • When heated in air it forms a black layer of copper (II) oxide on its surface.
    • Finely divided copper burns with a blue flame .
      Equation:
      2Cu(s)       +     O2(g)   →     2CuO(s)
      (Red brown)                               (Black)
  • Reaction with chlorine
    • Cu reacts with chlorine in presence of heat to form green copper (II) chloride.
      Equation:
      Cu(s) + Cl2(g) → CuCl2(s)
  • Reaction with Acids.
    • Copper does not react with dilute hydrochloric, nitric and sulphuric acids.
    • However it reacts with 50% Nitric acid, concentrated Nitric acid and concentrated sulphuric acid.

      With 50% Nitric (V) acid

    • Copper reduces the nitric acid to nitrogen monoxide.
      Equation;
      3Cu(s) + 8HNO3(l) → 3Cu(NO3)2(aq) + 4H2O(l) + 2NO(g) ;

      With concentrated Nitric (V) acid

    • Copper reduces the acid to brown nitrogen (IV) oxide/Nitrogen dioxide gas.
      Equation:
      Cu(s) + 4HNO3(l) → Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g)

      With concentrated sulphuric acid.

    • The sulphuric acid is reduced to sulphur (IV) oxide.
      Equation:
      Cu(s) + 2H2SO4(aq) → CuSO4(aq) + 2H2O(l) + SO2(g)

    Sample question
    1.  The flow chart below outlines some of the processes involved in the extraction of copper from copper pyrites. Study it and answer the questions that follow.
      worked example copper
      1. Give the name and formula of slag M. (1mark)
      2. Give the name of the reaction in chamber N. (1mark)
      3. Name the impure copper X. (1mark)
    2. Pure copper is obtained from impure copper by electrolysis.
      1. Name the anode the cathode and the electrolyte. (3 marks)
      2. Write equations for the reactions at the anode and cathode. (2 marks)
      3.  calculate the time taken for a current of 10 amperes to deposit 32kg of pure copper. (Cu = 64, 1F = 96000C) (3 marks)
    3. Draw a diagram to show how you would plate an aluminium spoon with copper

Lead

  • Is a transition element that combines with other elements to form compounds with 2 oxidation states.
  • It is among the group 4 elements;

Main Ores

  • Galena, PbS (lead sulphide); the main ore;
  • Cerrusite, PbCO3 (lead carbonate)
  • Anglesite, PbSO4 (lead II) sulphate);

Extraction of Lead

  • Occur in three main steps:
    • Ore concentration;
    • Extraction by reduction;
    • Purification (refining) by electrolysis;
  1. Ore concentration:
    • Is done by selective froth floatation;
    • The ore is ground into a fine powder, then water and a suitable oil added;
    • Air is then blown into the mixture; facilitating formation of a low density froth that floats on top;
    • Additionally, chemicals such as sodium cyanide and zinc sulphate are added to facilitate separation of zinc sulphide present in the ore.
    • The separated PbS is then dried and broken into smaller pieces, then subjected to reduction;
  2. Reduction:

    Step I: Roasting the ore:
    • The crushed and concentrated ore is roasted in a furnace to convert it to lead (II) oxide;
      Equation:
      2PbS(s) + 3O2(g) → 2PbO(s) + 2SO2(g)
    • During roasting some of the lead (II) sulphide is converted to lead (II) sulphate;
      Equation:
      PbS(s) + 2O2(g) → PbSO4(s);
    • Any lead sulphate formed is converted to lead silicate by silicon (IV) oxide;
    • The fate of lead (II) silicate;
      Note:
    • Additionally the lead (II) sulphate may further react with lead sulphide to form lead metal;

    Step II: Ore reduction:

    • The lead oxide obtained is mixed with coke, limestone and silica and some scrap iron;
    • The mixture is fed into the top of the Imperial smelting furnace (ISF); where it is melted using hot air blasts introduced near the bottom of the furnace;
      lead blast furnace
      Diagram: The Imperial Smelting furnace for Lead extraction.
      Main reactions:
      1. The lead (II) oxide is reduced to lead by the coke.
        Equation:
        PbO(s) + C(s) → Pb(s) + CO(s);
      2. The resultant carbon (IV) oxide produced in reaction (i) above further reduces any remaining lead (II) oxide;
        Equation:
        PbO(s) + CO(s) → Pb(s) + CO2(g);
      3. The scrap iron is added so as to react with any lead sulphide that may be present.
        Equation:
        Fe(s) + PbS(s) → Pb(s) + FeS(s)
      4. The limestone undergoes decomposition to give calcium oxide and liberate carbon (IV) oxide;
        Equation:
        CaCO3(s) → CaO(s) + CO2(g);
        The carbon (IV) oxide gets reduced by coke to form more carbon (II) oxide for reduction as in reaction (ii);
        Equation:
        CO2(g) + C(s) → 2CO(g);
      5. The calcium oxide reacts with silica in form of SiO2 to form calcium silicate;
        Equation:
        CaO(s) + SiO2(g) → CaSiO3(l) ;

        Waste gases and residues.
        • The iron sulphide and calcium silicate form a molten slag which is less dense and floats on top of molten lead at the bottom of the furnace;
        • From here the slag is separately tapped off;
        • Excess gases and air that did not react in the blast furnace escape through outlets at the top of the furnace;
        • These waste gases can be trapped and recycled;
        • These gases include: excess CO; excess CO2 ; oxygen; nitrogen; some SO2; and argon;

        Pollution effects:
        • Main pollutant is sulphur (IV) oxide from roasting of the ore.
        • Pollution control:

        Polution control

        • It is directly fed into a contact process plant or scrubbed using calcium hydroxide forming calcium sulphite;
  3. Purification (refining) of lead:
    • The molten lead obtained in this process contains impurities such as gold, silver, copper, arsenic, tin and sulphur;
    • The impure lead is refined by electrolysis.

      Electrolysis of molten lead.

      Electrolyte:

    • Any aqueous solution containing lead ions;
      The anode:
    • Impure lead;
      The cathode:
    • Pure lead;

      Electrolytic reactions;

      At the cathode.
      Observations:

    • Deposition of a grey solid.

      Explanations:

    • The lead (II) ions, Pb 2+ move to the cathode, where they accept electrons and undergo reduction.
    • Cations in the electrolyte are Pb 2+ and H + but Pb 2+ are preferentially discharged due to their easy tendency to undergo reduction.
      Equation:
      Pb2+(aq) + 2e- → Pb(s);

      At the anode
      Observations:

    • Dissolution of the anode, hence the impure lead rod decreases in size.

      Explanation
    • Since the metal rod is dipped into a solution of its ions, the impure lead solid undergoes oxidation, losing electrons to form lead (II) ions, Pb2+
    • Consequently as more lead (II) ions, Pb2+ get reduced at the cathode; more are released by the dissolving anode.
      Equation:
      Pb(s) → Pb2+(aq) + 2e-
      Overall reaction
      reduction of lead equation

Summary: flow chart on extraction of lead.
lead extraction flow diagram

Properties of Lead

Physical properties;

  • Has a low melting point but a high density;
  • The unusually low melting point of lead is difficult to explain using simple metallic bonding theory;
  • It is rather soft and pliable;
  • Relatively malleable;

Chemical properties.
Note: - Lead is fairly unreactive to most other metals;

Uses of Lead

  • It is used in several alloys e.g. solder and also added to bronze alloys to make them stronger;
  • Lead ingots are used in the manufacture of accumulators;
  • Being so malleable and so chemically inert lead sheeting was used for roofing (look at roofs of old churches and cathedral)-cost and pollution effects have however brought this to a stop;
  • Making lead pipes for water supply; this is also discouraged particularly in soft water areas due to threat of lead poisoning;
  • Making tetraethyl lead (IV) which for many years was used as a fuel additive to increase octane rating of fuels;
  • Used in weights, clock pendulums, plumb bobs etc; due to its high density;
  • It absorbs X-rays and hence lead aprons and lead glass are used to shield hospital radiographers;
  • Used for safe disposal or storage of radioactive substances since no radioactive emission has been known to pass through thick lead blocks;

Sample Question

The flow chart below illustrates the industrial extraction of lead metal. Study it and answer the questions that follow.
lead sample question

  1. Name the ore that is commonly used in this process.
  2. Explain what takes place in the roasting furnace.
  3. Identify gas P.

Download METALS: EXTRACTION PROPERTIES AND USES - Form 4 Chemistry notes.


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