Causes chemistry of rusting rust prevention introduction to oxidation reduction REDOX reactions gcse igcse KS4 science chemistry revision notes revising

2A METAL CORROSION and the RUSTING of IRON and its PREVENTION

Introduction

Corrosion is the destruction of materials (metal, stone etc.) by chemical reactions with substances in the environment.

Most metals corrode - chemically attacked, when in contact with oxygen and water.

(Reminder: Always say oxygen or oxygen in air, NOT air on its own, air is only ~20% oxygen)

The rusting of iron is an example of corrosion.

Rusting, like any other corrosion, involves oxidation and reduction - redox reactions.

The more reactive a metal, the more easily it is oxidised - the more easily the metal atoms lose electrons to form a positive ion.

This means the more reactive a metal the more likely it is to be corroded - means the metal is more susceptible to be attacked by a combination of oxygen (from air) and water - water is an important medium for most corrosion chemistry.

e.g. consider two extremes - when sodium is exposed to air at room temperature, it is rapidly coated in sodium oxide (oxidation-corrosion) whereas gold or platinum are very inert and do not corrode at all.

Both air (oxygen source) and water are necessary for iron to rust - need to know experiment to show this.

Corrosion by rusting can be prevented by applying a coating that acts as a barrier to air and water, such as greasing, painting or electroplating.

Again, you need to know about experiments to investigate these anti-corrosion methods.

In the case of the metal aluminium, this has an oxide coating that protects the metal from further corrosion.

Some rust prevention coatings are reactive and may contain corrosion inhibitors or a more reactive metal

e.g. magnesium or zinc. Zinc is used to galvanise iron and when scratched provides sacrificial protection because zinc is more reactive than iron and oxidises away first (preferentially).

Magnesium blocks can be attached to steel ships to provide sacrificial protection, but the blocks have to replaced when corroded away.

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The RUSTING PROCESS of iron - experiment to show to show what is needed

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 to rust iron and other metals.
- In corrosion the metal is oxidised
- The products of oxidation can be hydrated metal oxides/hydroxides (iron) and carbonates and maybe a mixture of both (the green surface of copper roofs).
- (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:
- The overall word and symbol equations for rusting are:
- 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, but is usually described as hydrated iron(III) oxide, which can be quite dry or quite soggy!
  - 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)
    - Initially iron(II) hydroxide is formed, which is further oxidised to iron(III) hydroxide - RUST!
  - (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 ...
  - (a) it involves iron gaining oxygen (Fe ==> Fe₂O₃)
  - (b) atoms of iron losing electrons
    - overall oxidation half equations are :
      - (i) Fe – 2e^– ==> Fe²⁺ (or Fe ==> Fe²⁺ + 2e^–)
      - (ii) Fe²⁺ – e^– ==> Fe³⁺ (or Fe²⁺ ==> Fe³⁺ + e^–)
      - overall (i) + (ii) = (iii) Fe – 3e^– ==> Fe³⁺ (or Fe ==> Fe³⁺ + 3e^–)
  - 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:
    - reduction half equation: O + 2e– ==> O^2–
    - or strictly speaking for the oxygen molecule from air: ¹/₂O₂ + 2e^– ==> O^2–
  - so this fits in with two important definitions of oxidation explained in detail below in the next section.
  - 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.

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RUST PREVENTION

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, eventually, completely!
  - This doesn't happen with aluminium, the initial aluminium oxide layer formed, remains intact and doesn't flake off, so the bulk of the aluminium doesn't corrode away like iron does.
- 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.
Most anti–corrosion methods to stop iron or steel rusting involve the exclusion of oxygen (in air) and water i.e. some sort of barrier is applied or formed on the surface of the metal to be protected.
Other methods do involve sacrificial corrosion, where some metal, more reactive than iron, is bolted onto the steel and corrodes first i.e. preferentially sacrificed to the attack of oxygen and water (more details on this method later).
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.
- You can also use a thin layer of plastic which acts as a water repellent barrier.
- Paint also has its decorative effect.
Moving parts on machines can be protected by a water repellent oil or grease layer i.e. this barrier keeps the water from reaching the iron or steel surface.

reactivity An experiment to investigate sacrificial corrosion

This 'rusting' corrosion can be prevented by connecting iron to a more reactive metal (e.g. zinc on iron or magnesium on steel).
- 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.
  - BUT you must use a metal that is more reactive than iron.
  - If you use a less reactive metal, it increases the rate of corrosion.
- 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 but it is a preferential oxidation reaction.
- The block of metal e.g. magnesium must be replaced when the bulk of it has corroded away.
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 (electroplating) by making the iron/steel the negative cathode or by dipping the iron/steel object in molten zinc.
  - For examples and full explanations see Electroplating surfaces with metals including silver and copper.
  - In the plating process, the iron/steel object is made the negative cathode in a bath of a zinc salt solution.
  - When the d.c. current flows, the zinc ions are reduced to zinc atoms which coat the cathode object.
  - Reduction is electron gain: Zn²⁺ + 2e^- ==> Zn(s)
- The zinc acts as a physical barrier between the iron/steel and oxygen/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 (a similar effect to the aluminium oxide layer that forms on aluminium).
- Also, if the surface is scratched, the exposed zinc again corrodes before the iron and continues to protect it by reforming the protective zinc oxide layer.
- 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 half equation, electron loss from zinc atoms)
  - which occurs preferentially to the oxidation of iron (and in most steel alloys too)
  - Fe ==> Fe²⁺+ 2e^– and Fe²⁺ ==> Fe³⁺ + e^– (the oxidation half equations, electron loss from iron)
  - For lots more on electroplating see Electrolysis and applications of electroplating
Steel tin cans are protected by a relatively unreactive thin layer of TIN
- This works well as long as the thin layer of tin is complete and acts as a relatively inert physical barrier between the steel and any 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!
  - The least reactive metals (compared to iron), copper and silver accelerate the corrosion process of iron.
  - The metals more reactive than iron, like zinc or magnesium, are sacrificially corroded first and protect the iron.

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Corrosion and the use of other metals

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, its sticks on hard to the aluminium metal surface, and so NOT exposing more aluminium to corrosion.
- This means aluminium structures last a lot longer than those made of iron or steel.
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!

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2B OXIDATION & REDUCTION – REDOX REACTIONS –

An INTRODUCTION for GCSE/IGCSE/O Level students

OXIDATION

The oxidation reactions you are most likely to come across at first in your chemistry course are the reaction of metals with oxygen e.g. by heating a strip of metal in a strongly in a bunsen flame.

The metal is considered to be oxidised because the metal gains and combines with oxygen atoms from the oxygen molecules in air e.g.

(i) magnesium + oxygen ===> magnesium oxide

2Mg + O₂ ===> 2MgO

The magnesium ribbon burns brightly to form the white oxide powder, oxidised – oxygen gain.

(ii) copper + oxygen ===> copper oxide

2Cu + O₂ ===> 2CuO

The copper doesn't burn like magnesium (not as reactive), but the black oxide forms on the copper surface.
The chemical species that gains oxygen is oxidised.
In both cases you can consider the oxidation in terms of electron loss
For Mg or Cu: M ==> M²⁺ + 2e^-
The metal atoms are oxidised by electron loss to give the metal ion.
Oxidation is electron loss by something (atom, ion or molecule).
These are called half equations and show the movement or transfer of electrons in a chemical reaction.
Lots more on the electron definitions of oxidation & reduction with examples explained.

REDUCTION

The opposite of oxidation is reduction and one simple definition of it is oxygen loss (opposite of oxidation above!).

You most likely to come cross this first when studying the extraction of metals from a metal oxide (maybe as an ore)

For example, if you strongly heat a mixture of copper oxide with carbon (charcoal or graphite powder) you find bits of brown–orange copper are formed.

copper oxide + carbon ===> copper + carbon dioxide

2CuO + C ==> 2Cu + CO₂

The copper oxide has lost oxygen, so copper oxide has been reduced.
Any chemical species that loses oxygen has been reduced.
Anything that removes oxygen is called the reducing agent (carbon here).

The carbon gains oxygen and so is oxidised.
Anything that gains oxygen is oxidised.
Anything that donates oxygen is called the oxidising agent (copper oxide here)

Another way to look at the reduction is from the electron transfer point of view.

Reduction is electron gain by something (atom, ion or molecule).
The carbon donates 2 electrons to the copper ion in the copper oxide compound.
Cu²⁺ + 2e^- ==> Cu

For more details see of some of these reactions see notes on the

Metal Reactivity Series,
Metal Reactivity Series Experiments–Observations
Extraction of Metals

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

Definition reminders and lots more EXAMPLES of OXIDATION and REDUCTION

In the examples in the first double column section below, the equations are not meant to be complete or balanced.

What I've highlighted, is the chemical change that fits the definition of oxidation and reduction,

in other words the transfer-movement of oxygen or electrons.

OXIDATION – definition and examples

(based on oxygen gain OR electron loss transfer)

REDUCTION – definition and examples

(based on oxygen loss OR electron gain transfer)

(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 (oxidation) of 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. the balanced half equations

(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, also part of rusting chemistry]

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

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

(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 reactions]

(e) An oxidising agent is the species that gives the oxygen to an atom, ion or molecule

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 an atom, ion or molecule

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 (oxygen loss) 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 (oxygen loss) 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 Thermite reaction)
- Fe₂O₃(s) + 2Al(s) ==> 2Fe(s) + Al₂O₃(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) Examples metal – metal salt solution displacement reactions are described and explained in detail.
- (a) 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
- The fully balanced symbol ionic equation is ...
- Mg(s) + Fe²⁺(aq) ==> Mg²⁺(aq) + Fe(s)
- the magnesium atom, Mg, loses 2 electrons (oxidation) to form the magnesium ion.
- The iron(II) ion, Fe²⁺, gains 2 electrons (reduced) to form iron atoms.
- You can think of it as two 'half reactions' involving the movement and transfer of electrons
  - Mg ==> Mg²⁺ + 2e^– (the oxidation half equation, electron loss, magnesium atom is oxidised)
  - Fe²⁺ + 2e^– ==> Fe (the reduction half equation, electron gain, iron ion is reduced)
  - The electron loss and gain cancel out, so you don't see them in the full equation because you have balanced the movement and transfer of electrons (e.g. 2 lost, 2 gained).
  - This involves balancing the electron losses and gains, so that no electron shows up in the full ionic equation.
- Mg is the reducing agent (electron donor) and the Fe²⁺ is the oxidising agent (electron remover or acceptor).
- The more reactive metal (magnesium) always displaces the less reactive metal (iron).
- 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).

reactivity

(b) Exactly the same reaction occurs if you add iron filings to copper sulfate solution.
- A brown precipitate of copper forms on the surface of the iron filings and the blue colour fades as the less reactive copper is displaced by the more reactive iron.
- The iron sulfate formed in solution is a very pale green so the final solution is almost colourless if excess iron is added.
- iron + copper(II) sulfate ==> iron(II) sulfate + copper
- Fe(s) + CuSO_4(aq) ==> FeSO_4(aq) + Cu(s)
- iron + copper(II) ion ==> iron(II) ion + copper
- The fully balanced symbol ionic equation is ...
- Fe(s) + Cu²⁺(aq) ==> Fe²⁺(aq) + Cu(s)
- The sulfate ion is a spectator ion and is NOT shown in the ionic equation.
- Again, to fully understand what is happening, you can think of it as two 'half–reactions'' involving the movement and transfer of electrons
  - Fe ==> Fe²⁺ + 2e^– (the oxidation half equation, electron loss, iron atom is oxidised)
  - Cu²⁺ + 2e^– ==> Cu (the reduction half equation, electron gain, copper ion is reduced)
  - The electron loss and gain cancel out, so you don't see the electrons in the full equation.
  - The full redox ionic equation is derived from combining the half equations in the right proportions - this involves balancing the movement and transfer of electrons.
  - This involves balancing the electron losses and gains, so that no electron shows up in the full ionic equation.
- Iron (Fe) is the reducing agent (electron donor) and the copper(II) ion (Cu²⁺) is the oxidising agent (electron remover or acceptor).
- -

reactivity

(c) Adding magnesium to blue copper(II) sulfate solution
- The blue colour fades as colourless magnesium sulfate is formed and brown bits of copper metal form a precipitate around the magnesium ribbon:
- magnesium + copper sulfate ==> magnesium sulfate + copper
  - Mg(s) + CuSO₄(aq) ==> MgSO₄(aq) + Cu(s)
  - The sulfate ion is a spectator ion.
  - magnesium atom + copper ion ==> magnesium ion + copper atom
  - Mg(s) + Cu²⁺(aq) ==> Mg²⁺(aq) + Cu(s)
  - ionic equation: Mg(s) + Cu²⁺(aq) ==> Mg²⁺(aq) + Cu(s)
    - Mg ==> Mg²⁺ + 2e^– (the oxidation half equation, electron loss, Mg atom is oxidised)
    - Cu²⁺ + 2e^– ==> Cu (the reduction half equation, electron gain, Cu²⁺ is reduced)
  - The magnesium atoms transfer electrons to the copper(II) ion.
  - Magnesium (Mg) is the reducing agent (electron donor) and the copper(II) ion (Cu²⁺) is the oxidising agent (electron remover or acceptor).
  - The full redox ionic equation is derived from combining the half equations in the right proportions' - this involves balancing the movement and transfer of electrons.
  - This involves balancing the electron losses and gains, so that no electron shows up in the full ionic equation.
- -
(d) copper + silver nitrate ==> silver + copper(II) nitrate
- Even some of the less reactive metals still facilitate displacement reactions.
- Cu(s) + 2AgNO₃(aq) ==> 2Ag + Cu(NO₃)₂(aq)
- the nitrate ion NO₃^– is the spectator ion
- copper atom + two silver ions ==> two silver atoms + copper(II) ion
- The fully balanced symbol ionic equation (redox equation) is ...
- Cu(s) + 2Ag⁺(aq) ==> 2Ag(s) + Cu²⁺(aq)
- The more reactive 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, copper atom oxidised)
  - 2Ag⁺ + 2e^– ==> 2Ag (the reduction, electron gain, silver ion reduced)
  - again, the electron loss and gain cancel out, so you don't see them in the full equation.
  - The full redox ionic equation is derived from combining the half equations in the right proportions.
  - Note that in this case, you need two lots of the silver half equation to one of the copper half equation to give the correctly balanced full ionic equation. Note 2 electrons lost balances 2 x 1 electron gained
- The more reactive metal (copper) always displaces the less reactive metal (silver).
- -

(e) Displacement of iron from its salts by more reactive metals.
- Iron forms more than one series of compounds - a characteristic of transition metals.
- e.g. if you add more reactive zinc or magnesium to iron(II) sulfate OR iron(III) sulfate solution, iron is deposited on the zinc or magnesium surface.
- You should also see that the magnesium reacts faster than the zinc.
- I am assuming you can write out the word equations.
- (i) Iron(II) sulfate displacement reactions
  - Mg(s) + FeSO₄(aq) ===> MgSO₄(aq) + Fe(s) (already done in (a) above)
  - Zn(s) + FeSO₄(aq) ===> ZnSO₄(aq) + Fe(s)
  - The half equations are
  - oxidations - electron loss: Zn ==> Zn²⁺ + 2e^-
  - reduction - electron gain: Fe²⁺ + 2e^- ==> Fe
  - These can now be combined to give the complete balanced ionic equation (redox equation)
  - Zn(s) + Fe²⁺(aq) ===> Zn²⁺ + Fe(s)
- - -
- (ii) iron(III) sulfate displacement reactions: (these are a bit more tricky to balance!)
  - 3Mg(s) + Fe₂(SO₄)₃(aq) ===> 3MgSO₄(aq) + 2Fe(s)
  - 3Zn(s) + Fe₂(SO₄)₃(aq) ===> 3ZnSO₄(aq) + 2Fe(s)
  - The half equations are
  - oxidations - electron loss: Mg ==> Mg²⁺ + 2e^- and Zn ==> Zn²⁺ + 2e^-
  - reduction - electron gain: Fe³⁺ + 3e^- ==> Fe
  - Again, these can now be combined to give the complete balanced ionic equations (redox equations)
  - 3Mg(s) + 2Fe³⁺(aq) ===> 3Mg²⁺ + 2Fe(s)
  - 3Zn(s) + 2Fe³⁺(aq) ===> 3Zn²⁺ + 2Fe(s)
  - -

(2) Metal - acid reaction:
- The reaction between a metal and an acid is also, technically, what is called a REDOX reaction.
- This means the reaction takes place in two parts, an oxidation involving electron loss and a reduction involving electron gain.
- If a metal is above hydrogen in reactivity, it can displace hydrogen from the acid - theoretically anything at least reactive as lead - which reacts VERY slowly with acids AND you never add sodium or potassium to acid - check out the chart !!!
- Here the metal atoms like Zn lose electrons to form positive ions (oxidation).
- The hydrogen ions, H⁺, gain electrons to form hydrogen gas molecules (reduction).
- The two simultaneous changes occur on the surface of the metal where the positive hydrogen ions hit the metal surface and pinch electrons from the metal and so metal ions pass into solution e.g.
- zinc + hydrochloric acid ==> zinc chloride + hydrogen
- Zn(s) + 2HCl(aq) ==> ZnCl_2(aq) + H_2(g)
- the chloride ion Cl^– is the spectator ion which doesn't take part in the chemical changes.
- zinc + hydrogen ion ==> zinc ion + hydrogen
- The fully balanced symbol ionic equation (redox equation) is ...
- Zn(s) + 2H⁺(aq) ==> Zn²⁺(aq) + H₂(g)
- The zinc transfers electrons to the hydrogen ions.
- 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 half equation, electron loss, zinc atom is oxidised)
  - 2H⁺ + 2e^– ==> H₂ (the reduction half equation, electron gain, hydrogen ion reduced)
  - again, the electron loss and gain cancel out, so you don't see them in the full equation.
- More on acid reactions including with metals

(3) Halogen displacement reactions
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)
- The potassium ion is a spectator ion, so the ionic equation is ...
- X₂(aq) + 2Y^–(aq) ==> 2X^–(aq) + Y₂(aq)
- where halogen X is more reactive than halogen Y, F > Cl > Br > I
- Halogen X is reduced by electron gain and halide ion Y^– is oxidised by electron loss.
  - in principle: X + e^– ==> X^– (electron gain, reduction of halogen atom X)
  - but more correctly: X₂ + 2e^– ==> 2X^– (electron gain, reduction of halogen molecule X₂)
  - in principle: Y^– ==> Y + e^– (electron loss, oxidation of halide ion Y^–)
  - but more correctly: 2Y^– ==> Y₂ + 2e^– (electron loss, oxidation of halide ion Y^– to give molecule Y₂)
- X₂ is the oxidising agent (electron acceptor)
- KY (via Y^–) is the reducing agent (electron donor)
- See GCSE Group 7 The Halogens – displacement reaction notes for more details

(4) 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!
- iron(II) chloride + chlorine ==> iron(III) chloride
- 2FeCl₂ + Cl₂ ==> 2FeCl₃
- 2FeCl₂(aq) + Cl₂(aq) ==> 2FeCl₃(aq)
- The iron(II) is oxidised to the iron(III) ion via the chlorine by electron loss:
  - Fe²⁺ – e^– ==> Fe³⁺ (an oxidation half equation)
- The chlorine molecule is reduced to the chloride ion by electron gain:
  - Cl₂ + 2e^– ==> 2Cl^– (a reduction half equation)
- 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.
- The full balanced symbol ionic equation is ...
- 2Fe²⁺ + Cl₂ ==> Fe³⁺ + 2Cl^–
- 2Fe²⁺(aq) + Cl₂(aq) ==> 2Fe³⁺(aq) + 2Cl^–(aq)

(5) 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 gained 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. the half equations are
  - 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^–

2C 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, 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.
- 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 at GCSE level.
    - 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!)

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Advanced Level Chemistry Redox Reaction Notes

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Index of notes on reaction of acids

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OTHER ASSOCIATED PAGE LINKS to do with metal reactivity

and 3. Metal Reactivity Series Experiments–Observations

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

and GCSE/IGCSE reactivity gap–fill worksheet or Rusting word–fill worksheet

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

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