Doc Brown's Chemistry  2. Corrosion of Metals e.g. iron & introducing Redox Reactions

Revision Notes KS4 Science IGCSE/O level/GCSE Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel 360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science  (revise courses equal to US grades 9-10)

SEE ALSO 1. The Reactivity Series of Metals (separate page)

3. Metal Reactivity Series Experiments-Observations (separate page)

Associated KS4 Science GCSE/IGCSE Chemistry notes pages: The Periodic Table * Group 1 Alkali Metals * Metal extraction * Transition metals * Alloys-uses of metals * Electrochemistry * Rates of Reactions (e.g. metal-acid) * Easy KS3 science multiple choice quiz start on metal reactivity and KS3 word-fills and GCSE m/c QUIZ on metal reactivity : Foundation Level or Higher Level & GCSE/IGCSE reactivity word-fill or Rusting word-fill  * EMAIL query?comment

• Iron (or steel) corrodes more quickly than most other transition metals and readily does so in the presence of both oxygen (in air) and water to form an iron oxide. You can do simple experiments to show that BOTH oxygen and water are needed. Put an iron nail into (1) boiled water in a sealed tube; (2) a tube of air and a drying agent; (3) an open test tube with water. Rusting appears overnight with (3) only.
• Rusting is speeded up in the presence of salt or acid solutions because of an increased concentration of ions. Corrosion is a redox process involving redox electron transfer and ion movement. The rusting metal behaves like a simple cell and more ions enable the current, and hence the electron transfer, to occur more readily.
• Rusting is overall:
  • iron + oxygen + water ==> hydrated iron(III) oxide
  • 4Fe_(s) + 3O_2(g) + xH₂O_(l) ==> 2Fe₂O₃.xH₂O_(s)
• i.e. rust is an orange-brown solid hydrated iron(III) oxide formed from the reaction with oxygen and water (the equation is not meant to be balanced and the amount of water x is variable, from dry to soggy!).
  • The reaction proceeds via (i) iron(II) hydroxide Fe(OH)₂ which is (ii) oxidised further to the hydrated Fe₂O₃, or if very soggy, it amounts to the formation of iron(III) hydroxide!
    • The reactions can be summarised in terms of hydroxide formation e.g.
    • (i) iron + water + oxygen ==> iron(II) hydroxide
    • 2Fe _(s) + 2H₂O _(l) + O_{2 (g)} ==> 2Fe(OH)_{2 (s)} 
    • (ii) iron(II) hydroxide + water + oxygen ==> iron(II) hydroxide
    • 4Fe(OH)_{2 (s)} + 2H₂O _(l) + O_{2 (g)} ==> 4Fe(OH)_{3 (s)} 
  • Rusting is an oxidation because it involves iron gaining oxygen (Fe ==> Fe₂O₃) or atoms of iron losing electrons (Fe - 3e^- ==> Fe³⁺).
  • See more examples of oxidation and reduction below.
• The rusting of iron is a major problem in its use as a structural material.
  • Preventing rusting adds cost to manufacturing things.
  • Corroded components or structures weakened, adding further costs in rust treatment or replacement.
• Iron and steel (alloy of iron) are most easily protected by paint which provides a barrier between the metal and air/water.
• Moving parts on machines can be protected by a water repellent oil or grease layer.

An experiment to investigate sacrificial corrosion

• This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc or magnesium). This is referred to as sacrificial protection or sacrificial corrosion, because the more reactive protecting metal is preferentially oxidised away, leaving the protected metal intact.
• The picture above illustrates what might be seen after a few days.* Iron or steel can also be protected by mixing in other metals (e.g. chromium) to make non-rusting alloys called stainless steel. The chromium, like aluminium, forms a protective oxide layer.
• * Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the more reactive 'sacrificed' metal.
• Coating iron or steel with a thin zinc layer is called 'galvanising'. The layer is produced by electrolytic deposition by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc (more details on the Extra Industrial Chemistry page). The zinc preferentially corrodes or oxidises to form a zinc oxide layer that doesn't flake off like iron oxide rust does. Also, if the surface is scratched, the exposed zinc again corrodes before the iron and continues to protect it. 
• Steel tin cans are protected by relatively unreacted tin and works well as long as the thin tin layer is complete and acts as an inert barrier between the steel and oxygen (air)/water HOWEVER, if a less reactive metal is connected to the iron, it then the iron rusts preferentially (try scratching a 'tin' can and leave out in the rain and note the corrosion by the scratch!)

• Aluminium does not oxidise (corrode) as quickly as its reactivity would suggest. Once a thin oxide layer of Al₂O₃ has formed on the surface, it forms a barrier to oxygen and water and so prevents further corrosion of the aluminium.
• Aluminium is a useful structural metal. It can be made harder, stronger and stiffer by mixing it with small amounts of other metals (e.g. magnesium) to make alloys.
• This property makes it a useful metal for out-door purposes e.g. aluminium window frames, greenhouse frames.

• Copper and Lead are both used in roofing situations because neither is very reactive and the compounds formed do not flake away as easily as rust does from iron. Lead corrodes to a white lead oxide or carbonate and copper corrodes to form a basic green carbonate (combination of the hydroxide Cu(OH)₂ and carbonate CuCO₃ e.g. seen as green roof on buildings).

• Both metals have been used for piping but these days lead is considered too toxic and copper is usually used as the stronger, but equally unreactive alloy with zinc, brass. Now of course, most piping is flowing in the plastic direction which doesn't corrode at all!

• Jewellery Metals

  • Silver's very low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its attractive silvery appearance.

  • Gold has an extremely low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its shiny attractive yellow appearance.

  • Platinum also has a very low reactivity makes it a valuable jewellery metal as it doesn't corrode easily and retains its attractive silvery appearance.

• The Group 1 Alkali Metals rapidly corrode in air and need to be stored under oil.

• Apart from their structural weakness they would hardly used for any outside purpose!

OXIDATION & REDUCTION - REDOX REACTIONS - INTRODUCTION for GCSE

Advanced Level Chemistry Redox Reaction Notes (it repeats this introduction and then moves on!)

OXIDATION - definition and examples

REDUCTION - definition and examples

(1) S ==> SO₂ [burning sulphur - oxidised]

(2) CH₄  ==> CO₂ + H₂O [burning methane to water and carbon dioxide, C and H gain O]

(3) NO ==> NO₂ [nitrogen monoxide oxidised to nitrogen dioxide]

(4) SO₂ ==> SO₃ [oxidising the sulphur dioxide to sulphur trioxide in the Contact Process for making sulphuric acid]

(1) CuO ==> Cu [loss of oxygen from copper(II) oxide to form copper atoms]

(2) Fe₂O₃ ==> Fe [iron(III) oxide reduced to iron in blast furnace]

(3) NO ==> N₂ [nitrogen monoxide reduced to nitrogen, catalytic converter in car exhaust]

(4) SO₃ ==> SO₂ [sulphur trioxide reduced to sulphur dioxide]

(1) Fe ==> Fe²⁺ + 2e^- [iron atom loses 2 electrons to form the iron(II) ion, start of rusting chemistry]

(2) Fe²⁺ ==> Fe³⁺ + e^- [the iron(II) ion loses 1 electron to form the iron(III) ion]

(3) 2Cl^- ==> Cl₂ + 2e^- [the loss of electrons by chloride ions to form chlorine molecules]

(1) Cu²⁺ + 2e^- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms, electroplating or displacement reaction)

(2)  Fe³⁺ + e^- ==> Fe²⁺  [the iron(III) ion gains an electron and is reduced to the iron(II) ion] 

(3) 2H⁺ + 2e^- ==> H₂ [hydrogen ions gain electrons to form neutral hydrogen molecules, electrolysis of acids or metal-acid reaction]

REDOX REACTIONS - in a reaction overall, oxidation and reduction must go together

(g) Redox reaction analysis based on the oxygen definitions

• (1) copper(II) oxide + hydrogen ==> copper + water
  • CuO_(s) + H_2(g) ==> Cu_(s) + H₂O_(g)
  • copper oxide reduced to copper, hydrogen is oxidised to water
  • hydrogen is the reducing agent (removes O from CuO)
  • copper oxide is the oxidising agent (donates O to hydrogen)
• (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide
  • Fe₂O_3(s) + 3CO_(g) ==> 2Fe_(l) + 3CO_2(g)
  • the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide
  • CO is the reducing agent (O remover from Fe₂O₃)
  • the Fe₂O₃ is the oxidising agent (O donator to CO)]
• (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide
  • 2NO_(g) + 2CO_(g) ==> N_2(g) + 2CO_2(g)
  • nitrogen monoxide is reduced to nitrogen
  • carbon monoxide is oxidised to carbon dioxide
  • CO is the reducing agent and NO is the oxidising agent
• (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the Thermit reaction)
  • Fe₂O_3(s) + 2Al_(s) ==> Al₂O_3(s) + 2Fe_(s)
  • iron(III) oxide is reduced and is the oxidising agent
  • aluminium is oxidised and is the reducing agent

(h) Redox reaction analysis based on the electron definitions

• (1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron
  • Mg_(s) + FeSO_4(aq) ==> MgSO_4(aq) + Fe_(s)
  • this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below.
  • The sulphate ion SO₄^2-_(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!!
  • magnesium + iron(II) ion ==> magnesium ion + iron
  • Mg_(s) + Fe²⁺_(aq) ==> Mg²⁺_(aq) + Fe_(s)
  • the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms.
  • Mg is the reducing agent (electron donor)  and the Fe²⁺ is the oxidising agent (electron remover or acceptor)
  • Displacement reactions involving metals and metal ions are electron transfer reactions.
• (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen
  • Zn_(s) + 2HCl_(aq) ==> ZnCl_2(aq) + H_2(g)
  • the chloride ion Cl^- is the spectator ion
  • zinc + hydrogen ion ==> zinc ion + hydrogen
  • Zn_(s) + 2H⁺_(aq) ==> Zn²⁺_(aq) + H_2(g)
  • Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor)
  • hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules
• (3) copper + silver nitrate ==> silver + copper(II) nitrate
  • Cu_(s) + 2AgNO_3(aq) ==> 2Ag + Cu(NO₃)_2(aq)
  • the nitrate ion NO₃^- is the spectator ion
  • copper + silver ion ==> silver + copper(II) ion
  • Cu_(s) + 2Ag⁺_(aq) ==> 2Ag_(s) + Cu²⁺_(aq)
  • copper atoms are oxidised by the silver ion by electron loss
  • electrons are transferred from the copper atoms to the silver ions, which are reduced
  • the silver ions are the oxidising agent and the copper atoms are the reducing agent
• (4) iron(II) chloride + chlorine ==> iron(III) chloride
• (5) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive)
  • X_2(aq) + 2KY_(aq) ==> 2KX_(aq) + Y_2(aq)
  • X_2(aq) + 2Y^-_(aq) ==> 2X^-_(aq) + Y_2(aq)
  • where halogen X is more reactive than halogen Y, F > Cl > Br > I
  • X is the oxidising agent (electron acceptor)
  • KY is the reducing agent (electron donor)
  • See GCSE Group 7 The Halogens - displacement reaction notes
• (6) Electrode reactions in electrolysis are electron transfer redox changes
  • at the negative cathode positive ions are attracted:
    • metal ions are reduced to the metal by electron gain:
    • Mⁿ⁺ + ne^- ==> M
    • n = the numerical charge of the ion and the number of electrons transferred
    • or 2H⁺_(aq) + 2e^- ==> H_2(g) (for the discharge of hydrogen)
  • at the positive anode negative ions are attracted:
    • negative non-metal ions are oxidised by electron loss e.g.
    • for oxide ions: 2O^2- - 4e^- ==> O₂ or 2O^2- ==> O₂ + 4e^-
    • for hydroxide ion: 4OH^- - 4e^- ==> O₂ + 2H₂O or 4OH^- ==> O₂ + 2H₂O + 4e^-
    • for halide ions (X = F, Cl, Br, I): 2X^- - 2e^- ==> X₂ or 2X^- ==> X₂ + 2e^-

Miscellaneous Extra Redox Notes

• Redox changes can often be observed as significant colour changes e.g.

  • iron + copper(II) sulphate ==> iron(II) sulphate + copper
    • Fe_(s) + CuSO_4(aq) ==> FeSO_4(aq) + Cu_(s)
    • iron + copper(II) ion ==> iron(II) ion + copper
    • Fe_(s) + Cu²⁺_(aq) ==> Fe²⁺_(aq) + Cu_(s)
    • Sulphate, SO₄^2-_(aq), is colourless BUT a blue to pale green colour change is observed in the solution as the blue copper(II) ion is replaced by the pale green iron(II) ion as well as the pink-dark precipitate of copper metal.

  • Potassium manganate(VII) is a powerful oxidising agent and an intense purple colour in water due to the MnO₄^- ion. In acidified solution it changes to an almost colourless* manganese(II) ion, Mn²⁺ when it oxidises something (* which actually is a very pale pink transition metal ion).

  • Potassium dichromate(VI) is another strong oxidising agent and is orange due to the dichromate(VI) ion, Cr₂O₇^2- ion. When it oxidises something it changes to the green chromium(III) ion, Cr³⁺.

  • Potassium iodide is a colourless salt dissolving in water to form a colourless solution. If it is oxidised e.g. with chlorine a yellow==>orange==>brown colour develops as iodine is formed from the colourless iodide ion.

• The use of Roman Numerals in names:

  • This indicates what is called the oxidation state of an atom in a molecule or ion.

  • It is easy to follow for simple metal ions because it equals the charge on the ion

    • e.g. the oxidation state of copper in the copper(II) ion is referred to as +2

    • the more electrons removed from the atom or ion by oxidation, the higher its oxidation state

    • e.g. Fe²⁺ - e^- ==> Fe³⁺, gives iron the oxidation state of +3 in the iron(III) ion

      • (via a suitable oxidising agent).

    • but for more complex ions things are not so simple.

      • in manganate(VII) ion, the Mn is in the +7 oxidation state

      • in dichromate(VI) ion, the Cr is in the +6 oxidation state

  • This topic is dealt with more thoroughly at AS-A2 advanced level chemistry.

Advanced Level Chemistry Redox Reaction Notes (it repeats this introduction and then moves on!)