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RUSTING and introduction to OXIDATION & REDUCTION

Doc Brown's Chemistry KS4 Science GCSE/IGCSE Chemistry Revision Notes

2. Corrosion of Metals e.g. iron & introducing redox reactions

In particular, the rusting of iron and corrosion prevention

What is an oxidation reaction? What is a reduction reaction? What do we mean by a REDOX reaction? How do we write oxidation/reduction equations? Why is rusting an oxidation? What is an oxidising agent? What is a reducing agent? What do you mean by the corrosion of metals? What is chemically happening when iron rust? How can we prevent iron from rusting? What is stainless steel? What is galvanising? What is an oxidation reaction? What is a reduction reaction? This page also includes an introduction to REDOX reactions.

Its a good idea to study the theory of oxidation & reduction before reading the technical details of rusting and rust prevention.

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

SEE ALSO 1. The Reactivity Series of Metals * 3. Metal Reactivity Series Experiments–Observations and ...

Metal Extraction Fe, Cu, Al etc.  *  Transition Metals  *  Other notes on using metals eg Al & Ti  *  Metal Structure – bonding

reactivity

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

 

METAL CORROSION and the RUSTING of IRON
  • The RUSTING PROCESS of iron
  • Iron (or steel) corrodes more quickly than most other transition metals and readily does so ONLY 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.
    • (1) Put an iron nail in pure water, but exposed to air. Lots of rust after a few days. The nail is well exposed to water and the oxygen in air.
    • (2) Put an iron nail into boiled water in a sealed tube, and a layer of oil too. The boiling drives off dissolved air and the oil provides an extra barrier. Very little rusting after a few days. If you do it very carefully, it can be quite some time for any rusting to show up. This shows with the oxygen from air, rusting will not happen.
    • (3) Put an iron nail in a sealed test tube of air and a drying agent (e.g. anhydrous calcium chloride, absorbs any moisture), very little, if any rusting even after quite a few days. The absence of water prevents rusting taking place.
    • (4) Put a nail in a test tube, but just exposed to air. Very little, if any, rusting observed. BUT it will eventually rust because there is water vapour in the air.
    • (5) You can extend this experiment by repeating experiment (1) with salt solution and then you see even faster rusting. Next time you are at the seaside take a few seconds to examine any railings where the paint has flaked off and the corroding effects of sea–spray will be very evident.
  • Rusting of iron 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.
  • Chemically the rusting of iron overall is:
    • iron + oxygen + water ==> hydrated iron(III) oxide (RUST)
    • 4Fe(s) + 3O2(g) + xH2O(l) ==> 2Fe2O3.xH2O(s)
    • x is a variable amount of water – extent of hydration, it can be very dry rust or very soggy rust!
  • 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) + O2(g) ==> 2Fe(OH)2(s) 
      • (ii) iron(II) hydroxide + water + oxygen ==> iron(II) hydroxide
      • 4Fe(OH)2(s) + 2H2O(l) + O2(g) ==> 4Fe(OH)3(s) 
    • Rusting is an oxidation because ...
      • (i) it involves iron gaining oxygen (Fe ==> Fe2O3)
      • (ii) atoms of iron losing electrons (Fe – 3e ==> Fe3+)
      • to fit in both definitions of oxidation (oxygen gain or electron loss), so the iron is oxidised.
      • Oxygen atoms/molecules from air gain electrons to form the oxide ion, so the oxygen is reduced:
        • O + 2e– ==> O2–
        • or strictly speaking: 1/2O2 + 2e ==> O2–
      • so this fits in with two important definitions of oxidation explained below.
      • It might be best to study the theory of redox reactions and then come back to the technical details of rusting chemistry and the chemistry of rust prevention described below.
    • 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, but the assessment of potential problems, and the cost of countering rusting must be taken into account in the cost of manufacturing iron and steel objects.
    • Rust is soft and crumbly and readily flakes off exposing more metal to water and oxygen (in air) i.e. the rusting chemistry just keeps on eating the metal away.
    • Corroded components or structures weakened, adding further costs in rust treatment or replacement.
    • Iron or steel objects near the coast rust faster because sea spray of salt water accelerates the rusting chemistry.
      • Rusting is a redox reaction and the presence of ions (Na+ and Cl from sodium chloride) enables the oxidation and reduction reactions to go faster.
    • Conversely, in some extremely dry desert regions, iron and steel objects barely rust at all.
      • Remember, both oxygen (from air) and water are needed for iron and steel to rust.
  • RUST PREVENTION
  • Iron and steel (alloy of iron) are most easily protected by paint (of any colour you want) which provides a physical barrier between the metal and air and water in the atmosphere or in contact with water containing dissolved oxygen.
  • Moving parts on machines can be protected by a water repellent oil or grease layer i.e. this keeps the water from reaching the iron or steel surface.

reactivityreactivityAn 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.
    • Theoretically, any iron ions formed by oxidation would be reduced by electrons from the oxidation of the more reactive 'sacrificed' metal.
    • The picture above illustrates what might be seen after a few days.
    • All the methods used ensure the iron or steel corrodes less readily than pure iron.
    • Blocks of a more reactive metal like magnesium can be bolted to the steel hulls of ships or underground iron pipes and the more reactive magnesium atoms preferentially lose electrons rather than the iron, i.e. the magnesium stops the iron rusting.
    • So the magnesium corrodes away leaving the iron intact.
    • Sacrificial corrosion is NOT a displacement reaction.
  • Steel, an alloy of iron and carbon, can also be protected by mixing in other metals (e.g. chromium) to make non–rusting steel alloys called stainless steels. The chromium, like aluminium, forms a protective oxide layer.
  • Coating iron or steel with a thin zinc layer is called 'galvanising'. Typical examples are steel buckets, steel nails and corrugated iron roofing.
    • 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 acts as a barrier between the iron/steel and air/water AND it has a second protective effect because it is higher in the metal reactivity series than iron ...
    • ... and the more reactive 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.
    • Galvanising is another example of sacrificial corrosion.
    • The zinc is preferentially oxidised by electron loss to the oxygen molecules in air.
      • Zn ==> Zn2+ + 2e       (the oxidation, electron loss)
      • which occurs preferentially to
      • Fe ==> Fe2+/Fe2+ + 2/3e       (the oxidation, electron loss)
  • TOPSteel 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.
    • The thin layer of tin plating acts as a barrier between the iron and water/air and tin is relatively unreactive metal.
    • 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!)
      • This is actually the reverse situation to sacrificial corrosion described above! Think it through for yourself!
      • Don't buy bashed tin cans from shops or supermarkets, you never know what it may taste like! hmm like rust!!!
      • You can demonstrate this effect by setting up an experiment by showing a more reactive metal like zinc protects iron wire by sacrificial corrosion, but less reactive metals like silver or copper actually accelerate the rate of corrosion i.e. the iron rusts faster!
  • ALUMINIUM CORROSION
  • 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.
    • The aluminium oxide layer doesn't flake off like rust does from iron or steel exposing more aluminium to corrosion.
  • 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. This is a combination of the hydroxide Cu(OH)2 and carbonate CuCO3 e.g. seen as corroded green copper roofs 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

(a) Oxidation is the gain or addition of oxygen by an atom, molecule or ion e.g. ...

(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]

(b) Reduction is the loss or removal of oxygen from a compound etc.  e.g.  ...

(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]

(c) Oxidation is the loss or removal of electrons from an atom, ion or molecule e.g.

(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]

(d) Reduction is the gain or addition of electrons by an atom, ion or molecule e.g. ...

(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]

(e) An oxidising agent is the species that gives the oxygen to a molecule

OR

an oxidising agent accepts electrons i.e. an oxidising agent removes electrons from some atom, ion or molecule.

In either process the oxidising agent gets reduced.

(f) A reducing agent is the species that removes the oxygen from a molecule

OR

a reducing agent acts as the electron donor i.e. it gives electrons to some atom, ion or molecule.

In either case the reducing agent gets oxidised.

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

(g) Redox reaction analysis based on the oxygen definitions of oxidation and reduction

  • (1) copper(II) oxide + hydrogen ==> copper + water
    • CuO(s) + H2(g) ==> Cu(s) + H2O(g)
    • Copper oxide reduced to copper by hydrogen, hydrogen is oxidised to water,
    • hydrogen is the reducing agent (removes O from CuO) and gets oxidised to water in the process,
    • copper oxide is the oxidising agent (donates O to hydrogen).
  • (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide
    • Fe2O3(s) + 3CO(g) ==> 2Fe(l) + 3CO2(g)
    • The iron(III) oxide is reduced to iron by the CO and the carbon monoxide is oxidised to carbon dioxide,
    • CO is the reducing agent (removes oxygen from Fe2O3) and gets oxidised in the process to CO2,
    • the Fe2O3 is the oxidising agent (O donator to CO).
  • (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide
    • 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g)
    • The nitrogen monoxide is reduced to nitrogen by oxygen loss,
    • carbon monoxide is oxidised to carbon dioxide by oxygen gain,
    • CO is the reducing agent (accepts oxygen) and gets oxidised in the process,
    • and the NO is the oxidising agent (donates oxygen to the CO) and is reduced to nitrogen in the process.
  • (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the Thermit reaction)
    • Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s)
    • Iron(III) oxide is reduced, oxygen loss, and is the oxidising agent, it oxidises aluminium,
    • aluminium is oxidised, oxygen gain, and Al is the reducing agent, it removes oxygen from the iron oxide.

(h) Redox reaction analysis based on the electron definitions of oxidation and reduction

  • (1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron
    • Mg(s) + FeSO4(aq) ==> MgSO4(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 SO42–(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.
    • You can think of it as two 'half–reactions'
      • Mg ==> Mg2+ + 2e     (the oxidation, electron loss)
      • Fe2+ + 2e ==> Fe       (the reduction, electron gain)
      • The electron loss and gain cancel out, so you don't see them in the full equation.
    • 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.
    • In metal – soluble metal salt displacement reactions, the metal atom always loses electrons (oxidation) and the metal ion always gains electrons (reduction).
  • (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen
    • Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)
    • the chloride ion Cl is the spectator ion
    • zinc + hydrogen ion ==> zinc ion + hydrogen
    • Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(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
    • Again, you can think of it as two 'half–reactions'
      • Zn ==> Zn2+ + 2e       (the oxidation, electron loss)
      • 2H+ + 2e ==> H2        (the reduction, electron gain)
      • again, the electron loss and gain cancel out, so you don't see them in the full equation.
  • (3) copper + silver nitrate ==> silver + copper(II) nitrate
    • Cu(s) + 2AgNO3(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
    • Again, you can think of it as two 'half–reactions'
      • Cu ==> Cu2+ + 2e       (the oxidation, electron loss)
      • 2Ag+ + 2e ==> 2Ag        (the reduction, electron gain)
      • again, the electron loss and gain cancel out, so you don't see them in the full equation.
  • (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)
    • X2(aq) + 2KY(aq) ==> 2KX(aq) + Y2(aq)
    • X2(aq) + 2Y(aq) ==> 2X(aq) + Y2(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 for more details
  • (6) If chlorine is bubbled into iron(II) chloride solution, iron(III) chloride is formed
    • Or you can just mix chlorine water with iron(II) chloride solution, a bit safer!
    • 2FeCl2 + Cl2 ==> 2FeCl3
    • 2FeCl2(aq) + Cl2(g/aq) ==> 2FeCl3(aq)
    • The iron(II) is oxidised to the iron(III) ion via the chlorine by electron loss: Fe2+ – e ==> Fe3+   (an oxidation)
    • The chlorine molecule is reduced to the chloride ion by electron gain: 1/2Cl2 + e– ==> Cl   (a reduction)
    • Chlorine is the oxidising agent because it gains electrons, the electron acceptor or remover,
    • and the iron(II) ion is the reducing agent because it donates or give electrons to the chlorine molecule.
    • If you put two of the two 'half–equations' together, you get the following full redox–ionic equation to match the 'molecular equation above. The original chloride ions are effectively spectator ions.
    • 2Fe2+ + Cl2 ==> Fe3+ + 2Cl
    • 2Fe2+(aq) + Cl2(g/aq) ==> 2Fe3+(aq) + 2Cl(aq)
  • (7) 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 and gain by the cation (positive ion)
      • or 2H+(aq) + 2e ==> H2(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: 2O2– – 4e ==> O2 or 2O2– ==> 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) + CuSO4(aq) ==> FeSO4(aq) + Cu(s)
      • iron + copper(II) ion ==> iron(II) ion + copper
      • Fe(s) + Cu2+(aq) ==> Fe2+(aq) + Cu(s)
      • Sulphate, SO42–(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, Cr2O72– 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.

    • In simple cases the oxidation number equals the number of electrons to be added (+) or electrons removed (–) to give the neutral element atom.

    • 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 and its not appropriate to explain them here.

        • 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!)

OTHER ASSOCIATED PAGE LINKS

reactivity

SEE ALSO (c) doc b 2. RUSTING & Introducing REDOX reactions

and 3. (c) doc b Metal Reactivity Series Experiments–Observations

  Easy KS3 science multiple choice quiz start on metal reactivity and (c) doc b KS3 word–fills

and GCSE/IGCSE m/c QUIZZES on metal reactivity

Foundation–tier Level (easier) multiple choice quiz on the Reactivity Series of Metals

or Higher–tier Level (harder) multiple choice quiz on the Reactivity Series of Metals

and (c) doc b GCSE/IGCSE reactivity gap–fill worksheet or (c) doc b Rusting word–fill worksheet

KS4 Science GCSE/IGCSE/O level Chemistry revision notes pages:

(c) doc b The Periodic Table  *  (c) doc b Group 1 Alkali Metals  *  (c) doc b Methods of Metal extraction

(c) doc b Transition Metals  *  (c) doc b Alloys–uses of metals  *  (c) doc b Electrochemistry–Electrolysis

(c) doc b Rates of Reactions Experiments (e.g. metal–acid)


Revision notes on rusting causes prevention oxidation reduction definitions reactions equations KS4 Science GCSE/IGCSE/O level Chemistry Information on rusting causes prevention oxidation reduction definitions reactions equations for revising for AQA GCSE Science, Edexcel Science chemistry IGCSE Chemistry notes on rusting causes prevention oxidation reduction definitions reactions equations OCR 21st Century Science, OCR Gateway Science notes on rusting causes prevention oxidation reduction definitions reactions equations WJEC gcse science chemistry notes on rusting causes prevention oxidation reduction definitions reactions equations CIE O Level chemistry CIE IGCSE chemistry notes on rusting causes prevention oxidation reduction definitions reactions equations CCEA/CEA gcse science chemistry (revise courses equal to US grade 8, grade 9 grade 10) science chemistry courses revision guides explanation chemical equations for rusting causes prevention oxidation reduction definitions reactions equations educational videos for rusting causes prevention oxidation reduction definitions reactions equations


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