candle combustion, example of exothermic burning reaction (c) doc bTYPES OF CHEMICAL REACTIONS or PROCESSES

Doc Brown's Science-Chemistry GCSE/IGCSE/AS Revision Notes



More than one 'descriptor' word can apply to a reaction, there is an 'alphabetical keyword' list below.

If you would like a brief description for a type of reaction which is not listed please EMAIL me

USE this alphabetical list of REACTION or PROCESS KEYWORDS for this page

Addition * Anodising * Burning-Combustion also (fast/slow/spontaneous combustion-fire triangle) * Catalytic converter * Contact Process * Cracking * Decomposition * Deliquescent * Dehydration * Displacement * Double decomposition * Electrolysis * ElectroplatingEndothermic reaction * Equilibrium * Esterification * Exothermic reaction * Fermentation * Galvanising * Haber Process * Hydration * Hygroscopic * Irreversible reaction * Neutralisation * Oxidation * Photosynthesis * Polymerisation (addition or condensation) * Precipitation * Redox * Reduction * Respiration * Reversible reaction * Rusting * Substitution * Synthesis * Thermit reaction * Thermal decomposition

Electrolysis , anodising and electroplating

  • electrolysis apparatus to collect two gases, at cathode and anode (c) doc bWhen substances which are made of ions are dissolved in water, or melted, they can be broken down (decomposed) into simpler substances by passing an electric current through them. This process is called electrolysis. The electrical conducting solution or melt of ions is called the electrolyte.

  • When an ionic substance is melted or dissolved in water the ions are free to move about and move to the electrical contacts called electrodes. The electrodes are usually inert e.g. carbon or platinum. The electron rich or negative electrode is called the cathode. The electron deficient or positive electrode is called the anode.

  • During electrolysis ions move to their oppositely charged electrode.

    • Positively charged ions, usually hydrogen or metal ions move to the negative electrode. Depending on the voltage, the positive ions may be reduced (e.g. Cu2+) by electron gain to deposit a metal (eg. Cu) or release hydrogen gas (H2) from hydrogen ions (H+).

    • At the same time negatively charged ions move to the positive electrode. The negative ions may be oxidised by electron loss. This usually results in the release of a non-metallic gas e.g. oxygen (O2) from hydroxide ions (OH-) or chlorine from chloride ions (Cl-)etc.

  • Electroplating is the process of coating a conducting material with a layer of metal using the process of electrolysis. The object to be coated is made the negative cathode and dipped into a salt solution of the metal ions of the metal to  form the coating. On passing a low d.c. voltage the metal is deposited on the conducting negative cathode. For more details see 4th link below.

  • Anodising a metal, like aluminium, is done by making it the positive anode in an electrolysis system. When the electrolyte is sulphuric acid solution, the oxygen formed at the anode oxidises the metal surface to make a thicker metal oxide layer.

  • For more detailed examples see ..

Exothermic and Endothermic Reactions or Changes
  • candle combustion, example of exothermic burning reaction (c) doc bEXOTHERMIC CHANGES

    • Heat is released or given out to the surroundings by the materials involved, so the temperature rises.

    • chemical change examples involve a new substance being formed and lots of examples (i) to (vi) below (but they are not always exothermic - see endothermic below).

    • physical change examples e.g. condensation, freezing etc. all require the removal of energy from the material e.g. water, to the surroundings to produce the change in state (its the same as releasing heat, but it doesn't seem like it!).

      • Note: Dissolving substances in water can release heat giving a warm/hot solution e.g. diluting concentrated sulphuric acid.

        • At KS3-GCSE level this exothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!

  • At a higher level of thinking for exothermic chemical changes: The net energy change when the energy needed to break bonds in the reactants is less than the energy released when new bonds are formed in the products.

  • A burning or combustion reaction usually means a very fast exothermic reaction where a flame is observed. It involves a highly energetic oxidation of 'fuels' where the temperature generated is so high the atoms give off light from the luminous flame zone  e.g.

    • (i) bunsen flame as methane gas fuel burns ...

      • methane + oxygen ==> carbon dioxide + water

      • CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(l) 

      • This is complete combustion with a pale blue flame and the products cannot react any further with oxygen.

      • If the oxygen supply is limited the flame is more yellow and can be 'smokey' due to soot formation (C) and dangerous since carbon monoxide (CO) can be formed.

      • These are examples of incomplete combustion.

        • methane + oxygen ==> carbon monoxide + water

        • 2CH4(g) + 3O2(g) ==> 2CO(g) + 4H2O(l) (carbon monoxide formation)

          • Most people who die in house fires are poisoned by carbon monoxide (and other toxic gases) in the thick smoke rather than from burns.

        • or

        • methane + oxygen ==> carbon (soot) + water

        • CH4(g) + O2(g) ==> C(s) + 2H2O(l) (soot formation)

        • The sooty carbon particles e.g. in a candle flame, are heated to such a high temperature they become incandescent and give out yellow light, but as far as I know virtually no carbon monoxide is formed!

      • See also KS4 Science GCSE/IGCSE Chemistry Notes on fossil fuel combustion

    • (ii) passing chlorine over hot aluminium metal to make aluminium chloride, the aluminium burns to form the chloride ...

      • aluminium + chlorine ==> aluminium chloride

      • 2Al(s) + 3Cl2(g) ==> 2AlCl3(s) 

    • (iii) burning magnesium ribbon with a bright white flame ...

      • magnesium + oxygen ==> magnesium oxide

      • 2Mg(s) + O2(g) ==> 2MgO(s) 

    • (i) to (iii) are all oxidation reactions, as are all 'fuel' burning reactions.

  • Continuous combustion requires the 'fire triangle' of heat + fuel + oxidant (oxidants like oxygen, air or other reactive gases like chlorine or fluorine and in rockets liquids like hydrogen peroxide)

    • Very fast or explosive combustion:

      • A roaring bunsen flame (of methane burning) is an example of fast combustion and when the air (oxygen) - methane (natural gas) mixture is first ignited it is a small explosion! (equation above). It seems contradictory, but a source of ignition is needed because the C-H and O=O bonds are very strong giving a high activation energy. However, once ignited, the heat from the flame keeps the burning going.

      • Another explosive example is the 'squeaky pop test for hydrogen'. When a lit splint is applied there is a faint blue flame for a fraction of a second as the two gases explode to form water + heat, light and sound energy!

        • hydrogen + oxygen ==> water

        • 2H2(g) + O2(g) ==> 2H2O(l)

      • In all these cases the high temperature reaction zone is seen as flame and an initial high energy source for ignition is needed to initiate the reaction e.g. a match or an electrical discharge.

    • Slow or smouldering combustion:

      • In these cases no flame is seen, but a high temperature heat source is still required to start the reaction and the reaction zone is still at a high temperature e.g. the red hot slow burning of charcoal (mainly carbon), but the main combustion product is still carbon dioxide. You can only get this slow/smouldering combustion with solid combustible reactants.

      • Gases will tend to explode unless controlled in a burner and liquids will vaporise in the heat from the flame and so will also burn very fast with a flame.

        • carbon + oxygen ==> carbon dioxide

          • C(s) + O2(g) ==> CO2(g)

          • This is an example of complete combustion.

        • BUT quite often, with limited air/oxygen supply, carbon monoxide is readily formed,

        • carbon + oxygen ==> carbon monoxide

          • 2C(s) + O2(g) ==> 2CO(g)

          • This is an example of incomplete combustion.

    • Spontaneous combustion:

      • This is when combustion occurs without any application of a high energy ignition source, sometimes described as self-ignition, though in some cases heat is generated in some way which triggers the reaction.

      • For example, it is possible to prepare a very finely divided black powder form of iron(II) oxide. When the powder is dropped through air lots of tiny flashes of light are seen as it burns to form another iron oxide (probably Fe3O4). The reaction is triggered by heat from friction. The powder has such a large surface area that the friction caused by just falling in air produces enough heat to initiate the reaction. Powdered coal dust or very fine flour can behave in the same way and both have been responsible for serious accidents in industry.

      • Other substances can spontaneously ignite in air because the activation energies required are so low and the kinetic energy of the particles is sufficient for the reaction to happen without help! e.g. the highly reactive Group 1 Alkali Metal caesium and the silicon-hydrogen compounds called silanes (SiH4, Si2H6 etc. which are like organic alkanes with the C's replaced by Si and far less stable).

      • Potassium and all the alkali metals below it ignite in water (Rb and Cs explosively) because the reaction is so exothermic and ignites metal vapour and the hydrogen gas produced.

  • See other web page for more detailed notes on energy changes and calculations.

  • BUT many exothermic reactions are not as dramatic as burning with a flame! e.g.

    • (iv) Respiration: the relatively slow 'burning' of carbohydrates in animals/plants, but it releases plenty of energy at 37oC!

      • glucose + oxygen ==> carbon dioxide + water + energy

      • C6H12O6(aq) + 6O2(g) ==> 6CO2(g) + 6H2O(l) + energy

    • (v) Neutralisation: acid + alkali ==> salt + water

      • e.g. hydrochloric acid + sodium hydroxide ==> sodium chloride + water

      • which is one of the fastest reactions in water, but the mixture only warms up by 5 to 10oC! A bunsen flame reaches 1200oC in the main combustion zone!

      • More details further down.

    • (vi) Rusting in which iron slowly reacts with water and oxygen (from air) to form the orange-brown hydrated iron oxide we call rust.


    • Heat is absorbed or taken in by the materials involved from the surroundings, the system cools or has to be heated to effect the change.

    • Chemical change examples of endothermic reactions e.g. thermal decomposition of limestone, cracking oil fractions, decomposition by electrolysis etc.

      • (i) Photosynthesis: input of energy from sunlight needed

        • carbon dioxide + water ==> glucose + oxygen

        • 6CO2 + 6H2O + sunlight energy ==> C6H12O6 + 6O2 

      • (ii) Making lime by heating limestone to over 900oC where a net input/absorption of energy is needed to bring about this thermal decomposition ...

        • limestone ==> quicklime + carbon dioxide

        • calcium carbonate ==> calcium oxide + carbon dioxide

        • CaCO3(s) ==> CaO(s) + CO2(g) 

      • (iii) Cracking hydrocarbon molecules from oil to make smaller molecules, also requires this absorption of heat by the reactant molecules to break em' up', or to put it 'poshly', another example thermal decomposition e.g.

        • hexane => ethene + butane

        • C6H14 ==> C2H4 + C4H10 

    • Physical change examples e.g. melting, boiling, evaporation etc. all require the input of energy to effect the change of state of the material.

      • Note: Dissolving substances in water can absorb heat giving a cool solution e.g. dissolving ammonium nitrate salt in water.

        • At KS3-GCSE level this endothermic 'dissolving' is considered a physical change, but at AS-A2 level they may be considered a chemical change too!

  • At a higher level of thinking for endothermic chemical changes. The net energy change when the energy needed to break bonds in the reactants is more than the energy released when new bonds are formed in the products.

  • See also KS4 Science GCSE/IGCSE Chemistry Notes on fossil fuel combustion

Decomposition and Thermal Decomposition
  • Decomposition in general means to break down into small species e.g. natural organic matter decomposes with enzymes into carbon dioxide, water and nitrogen etc.

    • Fermentation is form of biological degradation, catalysed by enzymes, to break down glucose sugar into the smaller molecules of ethanol ('alcohol') and carbon dioxide ...

      • C6H12O6(aq) ==> 2C2H5OH(aq) + 2CO2(g)

  • Light can cause decomposition e.g. in photography is a sort of photo-decomposition.

    • silver chloride + light ==> silver + chlorine

    • 2AgCl ==> 2Ag + Cl2

  • Thermal decomposition means to break down substances into two or more substances by heat (usually endothermic reactions at temperatures well above room temperature) e.g.



OXIDATION - definition and examples

REDUCTION - definition and examples

The gain or addition of oxygen by an atom, molecule or ion e.g. ...

(1) S + O2 ==> SO2 [burning sulphur - oxidised to sulphur dioxide]

(2) CH4 + O2 ==> CO2 + 2H2O [burning methane to water and carbon dioxide, methane oxidised as the C and H atoms gain O]

(3) 2NO + O2 ==> 2NO2 [nitrogen monoxide is oxidised to nitrogen dioxide by gaining oxygen]

(4) SO32- + [O] ==> SO42- [oxidising the sulphite ion to the  sulphate ion]

The loss or removal of oxygen from a compound etc.  e.g.  ...

(1) CuO + H2 ==> Cu + H2O [loss of oxygen from copper(II) oxide shows it to be reduced to copper atoms]

(2) Fe2O3 + 3CO ==> Fe + 3CO2 [iron(III) oxide ore is reduced to iron metal by oxygen loss in the blast furnace]

(3) 2CO + 2NO ==> CO2 + N2 [nitrogen monoxide reduced to nitrogen by losing oxygen]

(4) CuO + Mg ==> Cu + MgO [loss of oxygen from copper(II) oxide shows it to be reduced to copper atoms]

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]

(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 in electrolysis of chlorides or halogen displace]

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 e.g. in electrolysis or metal displacement reactions)

(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]

(4) Cl2 + 2e- ==> 2Cl- [chlorine molecules gain electrons to form chloride ions

An oxidising agent is the species that gives the oxygen or removes the electrons A reducing agent is the species that removes the oxygen or acts as the electron donor

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

Redox reaction analysis based on the oxygen definitions

  1. copper(II) oxide + hydrogen ==> copper + water
    • CuO(s) + H2(g) ==> Cu(s) + H2O(g)
    • copper oxide reduced to copper, hydrogen is oxidised to water, hydrogen is the reducing agent (removes O from CuO) and 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, the carbon monoxide is oxidised to carbon dioxide, CO is the reducing agent (O remover from Fe2O3) and the Fe2O3 is the oxidising agent (O donator to CO)
  3. 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g)
    • nitrogen monoxide is reduced to nitrogen, carbon monoxide is oxidised to carbon dioxide, CO is the reducing agent, NO is the oxidising agent
  4. iron(III) oxide + aluminium ==> aluminium oxide + iron
    • Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s)
    • iron(III) oxide is reduced and is the oxidising agent, aluminium is oxidised and is the reducing agent - incidentally, this is the 'thermit' reaction!
  5. For more details of some of these and similar reactions see ...

Redox reaction analysis based on the electron definitions

  1. magnesium + iron(II) sulphate ==> magnesium sulphate + iron
    • Mg(s) + FeSO4(aq) ==> MgSO4(aq) + Fe(s)
    • [this is the 'ordinary' equation 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!!]
    • 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 and the Fe2+ is the oxidising agent](2)(i) zinc + hydrochloric acid ==> zinc chloride + hydrogen
  2. zinc + hydrochloric acid ==> zinc chloride + hydrogen
    • Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g)
    • the chloride ion Cl- is the spectator ion]
    • Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g)
    • Zinc atoms are oxidised to zinc ions by electron loss and zinc is the reducing agent, 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) + 2AgNO3(aq) ==> 2Ag(s) + Cu(NO3)2(aq)
    • the nitrate ion NO3- is the spectator ion
    • Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq)
    • copper atoms are oxidised by the silver ion, electrons are transferred from the copper atoms to the silver ions, which are reduced. Silver ions are the oxidising agent and copper atoms are the reducing agent
  4. chlorine + potassium iodide ==> potassium chloride + iodine
    • Cl2(aq) + 2KI(aq) ==> 2KCl(aq) +  I2(aq)
    • or  Cl2(aq) + 2I-(aq) ==> 2Cl-(aq) +  I2(aq)
    • a halogen displacement reaction, the more reactive chlorine displaces the less reactive iodine.
    • Chlorine is reduced by electron gain and the iodide ions are oxidised by electron loss.
  5. For more details of similar reactions see the Metal Reactivity, Metal Extraction and Group 7 The Halogens notes.




means joining many small molecules called monomers into a long molecules of many units called a polymer and there are two principal types of polymerisation process

(1) Addition polymers are formed by (e.g. alkene) monomers adding together and forming no other products except the polymer e.g. two examples of addition polymerisation are

ethene ==> poly(ethene)

phenylethene ==> poly(phenylethene), old name polystyrene

(2) Condensation polymers are formed by one or more monomers add together, forming the polymer BUT in forming the polymer small molecules are eliminated 'between' the monomers e.g. two examples of condensation polymerisation are ...

dicarboxylic acid + diol ==> polyester + water

diamine + dicarboxylic acid ==> nylon + water

(1) Example equations showing addition polymerisation

diagram of polymerisation of ethene monomer to poly(ethene) polymer (c) doc b

(Ex. 1a) formation of poly(ethene) or 'polythene' from polymerising ethene to form an addition polymer. No other molecule is formed - just simple addition polymerisation.
diagram of polymerisation of chloroethene monomer to poly(chloroethene) PVC polymer (c) doc b (Ex. 1b) formation of poly(chloroethene) or 'PVC' from polymerizing chloroethene  to form an addition polymer. No other molecule is formed - just simple addition polymerization.
For more examples and details of addition polymers see Useful Oil Products Part 7
(2) Example equation illustrating condensation polymerisation

diagram of a Nylon made from two different monomers by condensation polymezisation (c) doc b

+ small molecules eliminated

In the case of Nylon, for each 'red' monomer - 'blue' monomer, a link is formed at each end of each monomer molecule by eliminating a water molecule e.g. where [R] = 'rest of molecule' a single link formation reaction can be shown as

[R]-COOH + HO-[R] ==> [R]-CO-O-[R] + H2O

(Example of 2) representation of a Nylon made from two different monomers (shown as red and green + linking atoms) joining by eliminating a small molecule between the two monomers, therefore Nylon is a condensation polymer.
For more examples and details of condensation polymers see Useful Oil Products Part 11


  • NEUTRALISATION usually involves mixing an acid (pH <7 if soluble) with a base or alkali (pH > 7 if soluble) which react to form a neutral salt solution of around pH7

  • Two situations are common:

  • (1) Water soluble bases, called alkalis and often hydroxides, are mixed with a soluble acid such as hydrochloric, citric, sulphuric or nitric acid.

    • acid + base/alkali ==> salt + water

    • e.g. sodium hydroxide + hydrochloric acid ==> sodium chloride + water

      • NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l)

    • certain carbonates like sodium carbonate, are also soluble to form alkaline solutions, and they will be similarly neutralised with 'fizzing' as carbon dioxide is formed as a 3rd product

    • e.g. sodium carbonate + hydrochloric acid ==> sodium chloride + water + carbon dioxide

      • Na2CO3(aq) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l) + CO2(g) 

  • (2) Dissolving a water insoluble base (often an oxide) in an acid

    • e.g. copper oxide + sulphuric acid ==> copper sulphate + water

      • CuO(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l)

    • the acid can also be neutralised with a metal or a carbonate to give a salt solution

    • metal + acid ==> salt + hydrogen (this is also a redox reaction)

    • e.g. zinc + hydrochloric acid ==> zinc chloride + hydrogen

      • Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g) 

    • insoluble carbonate + acid ==> (often soluble) salt + water + carbon dioxide

    • e.g. magnesium carbonate + sulphuric acid ==> magnesium sulphate + water + carbon dioxide

      • MgCO3(s) + H2SO4(aq) ==> MgSO4(aq) + H2O(l) + CO2(g) 

  •  More details  of these types of reactions involving acids and bases-alkalis


  • A precipitation reaction is generally defined as 'the formation of an insoluble solid either by mixing two solutions or bubbling a gas into a solution.

  • The silver halide salts are used in photography and can be made by precipitation on mixing solutions of two soluble salts e.g.

    • silver nitrate + potassium chloride ==> silver chloride + potassium nitrate

    • AgNO3(aq) + KCl(aq) ==> AgCl(s) + KNO3(aq)

    • Ag+(aq) + Cl-(aq) ==> AgCl(s)

      • is the ionic equation (NO3- and K+ are spectator ions - they don't take part in the reaction)

  • Insoluble barium sulphate can be made in a similar manner by mixing dilute sulfuric acid with barium chloride solution

    • barium chloride + sulphuric acid ==> barium sulphate + hydrochloric acid

    • BaCl2(aq) + H2SO4(aq) ==> BaSO4(s) + 2HCl(aq)

    • Ba2+(aq) + SO42-(aq) ==> BaSO4(s)

      • is the ionic equation (Cl- and H+ are spectator ions - they don't take part in the reaction)

  • Calcium carbonate is precipitated when carbon dioxide gas is bubbled into calcium hydroxide solution

    • calcium hydroxide + carbon dioxide ==> calcium carbonate + water

    • Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l)


Reversible Reactions

  • Generally speaking most chemical reactions are irreversible, that mean they go 100% one way to the products and that's that! e.g. magnesium fizzing away in hydrochloric acid to form magnesium chloride and hydrogen. There is no feasible way of reacting hydrogen and magnesium chloride to re-form magnesium metal and hydrochloric acid!

  • However a reversible reaction is a chemical reaction in which the products can be converted back to reactants under suitable conditions and there are quite a few examples of them.

  • A reversible reaction is shown by the sign (c) doc b i.e. a half-arrow to the right (forward reaction) and a half-arrow to the left (backward reaction).

  • As pointed out above, most reactions are not reversible and have the usual complete arrow (c) doc b only pointing to the right i.e. in the direction the reaction goes 100%.

Two examples of reversible reactions are given below:

(a) The thermal decomposition of ammonium chloride

On heating strongly, the white solid ammonium chloride, decomposes into a mixture of two colourless gases - ammonia and hydrogen chloride. On cooling the reaction is reversed and solid ammonium chloride reforms.

Ammonium chloride + heat (c) doc b ammonia + hydrogen chloride

NH4Cl(s) (c) doc b NH3(g) + HCl(g)

(b) The thermal decomposition of hydrated copper(II) sulphate

  • On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed (left to right reaction).

  • When the white solid is cooled and water added, blue hydrated copper(II) sulphate is reformed (right to left reaction).

blue hydrated copper(II) sulphate + heat (c) doc b white anhydrous copper(II) sulphate + water

CuSO4.5H2O(s) (c) doc bCuSO4(s) + 5H2O(g)

 For more details of reversible reactions for IGCSE/GCSE science-chemistry

Reversible reactions and Chemical Equilibrium (more details for GCSE)
  • When a reversible reaction occurs in a closed system an equilibrium is formed, in which the original reactants and products formed coexist.

  • In an equilibrium there is a state of balance between the concentrations of the reactants and products.

  • At equilibrium the rate at which the reactants change into products is exactly equal to the rate at which the products change back to the original reactants.

  • The result is that that the concentrations of the reactants and products remain the same BUT the reactions don't stop!

  • However the relative amounts of the reactants and products depend on the reaction conditions e.g. the temperature and pressure.

  • You can change the conditions to favour a particular reaction direction e.g. in a limekiln the carbon dioxide is vented out so all the limestone changes to lime avoiding an.

(a) The formation of calcium oxide (lime) from calcium carbonate (limestone)

calcium carbonate (limestone) (c) doc b calcium oxide (lime) + carbon dioxide

CaCO3(s) (c) doc b CaO(s) + CO2(g)

(b) The formation of ammonia

nitrogen + hydrogen (c) doc b ammonia

N2(g) + 3H2(g) (c) doc b 2NH3(g)

Hydration and dehydration (often reversible under the right conditions)
  • Hydration means the addition of water or combining with water.

  • Dehydration means the losing or removal of water.

  • Example 1: On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed (left to right reaction is a dehydration, water lost).

  • When the white solid is cooled and water added, blue hydrated copper(II) sulphate is reformed (right to left reaction is a hydration, water gained).

    • blue hydrated copper(II) sulphate + heat (c) doc b white anhydrous copper(II) sulphate + water

      • CuSO4.5H2O(s) (c) doc b CuSO4(s) + 5H2O(g)  

      • The water in the blue crystals (the 5H2O) is called water of crystallisation and becomes part of the crystal structure when a concentrated solution of copper sulphate is slowly evaporated to allow crystallisation e.g. in a salt preparation.

      • The left to right dehydration reaction also happens if blue hydrated copper sulphate is treated with cold/warm concentrated sulphuric acid, which is a powerful dehydrating agent (see sugar reaction further down in this section).

      • Summing up for this example:

        • From left to right is dehydration and from right to left is hydration and this connection will always be so for this type of reversible reaction.

  • Example 2: If the alcohol ethanol is heated with concentrated sulphuric acid it is dehydrated to form the alkene gas ethene plus water.

  • The same reaction happens if ethanol vapour is passed over very hot aluminium oxide which also catalyses the dehydration of the ethanol.

    • ethanol ==> ethene + water

    • a dehydration reaction - also a particular type of elimination reaction

    • ethanol, 'alcohol' (c) doc b ==> ethene (c) doc b + H2O

  • Example 3: If ethene gas is dissolved in concentrated sulphuric acid and diluted with water, on gentle boiling the alcohol ethanol is formed (the reverse of example 2). The same reaction occurs if ethene gas and water vapour are passed over a silica gel-phosphoric acid catalyst at 300oC. In both cases the ethene is hydrated to form ethanol by water addition.

    • ethene + water ==> ethanol

    • a hydration reaction - also a particular type of an addition reaction

    • ethene (c) doc b + H2O ==> ethanol, 'alcohol' (c) doc b

  • Example 4: If white crystalline sugar is heated with concentrated sulphuric acid a spongy mass of carbon is formed.

    • The acid dehydrates the sugar, i.e. it removes the equivalent of 'H2O' and leaves the 'C' atoms!

      • e.g. glucose ==> carbon + water

      • C6H12O6 ==> 6C + 6H2O

      • a dehydration reaction - the removal or elimination of water



DISPLACEMENT REACTIONS (usually with a 'reactivity series')
  • Displacement essentially means one part of a compound is replaced by another.

  • The reactants and products are similar BUT one 'bit' has been swapped with another, a sort of 'swap around'.

  • Displacement reactions occur when a more reactive metal displaces a less reactive metal from one of its compounds e.g. from the oxide, sulphate, chloride or nitrate.

  • Displacement reactions also occur when a more reactive non-metal displaces a less reactive non-metal from one of its compounds e.g. from the metal salt.

  • These experiments can be used to establish a so called 'Reactivity Series'.

Metal displacements

Non-metal displacements
  • e.g. the spectacular very exothermic 'thermit reaction' which shows aluminium is more reactive than iron.  When a mixture of aluminium and iron(III) oxide is ignited with a magnesium fuse (high activation energy), the mixture burns furiously with a shower of sparks to leave a red hot blob of iron and a white as of aluminium oxide ...

  • iron(III) oxide + aluminium ==> aluminium oxide + iron

  • Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s)  

  • or more reactive copper 'gently' displaces silver from silver nitrate solution to make silvery plated copper

  • copper + silver nitrate ==> copper(II) nitrate + silver

  • Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq)

  • more details in Metal Reactivity Notes

  • e.g. halogen displacement, more reactive chlorine displaces bromine from potassium bromide

  • chlorine + potassium bromide ==> potassium chloride + bromine

  • Cl2(aq) + 2KBr(aq) ==> 2KCl(aq) + Br2(aq) 

  • or more reactive bromine displaces iodine from sodium iodide

  • bromine + sodium iodide ==> sodium bromide + iodine

  • Br2(aq) + 2NaI(aq) ==> 2NaCl(aq) + I2(aq)

  • more details in KS4 Science GCSE/IGCSE Chemistry Notes on Group 7 Halogen Notes


Other processes and reactions


This means to coat iron or steel with a layer of zinc to stop it rusting (more details on Metal Reactivity page)

 Haber Process

The synthesis of ammonia by combining nitrogen and hydrogen using high temperature, pressure and an iron catalyst. (all the details)
 Contact Process
  • Part of the manufacture of sulphuric acid from:  (Full details of Contact Process)
  • sulphur ==> sulphur dioxide ==V catalyst==> sulphur trioxide* ==> sulphuric acid.
  • S(s) + O2 (g) ==> SO2 (g), * 2SO2 (g) + O2 (g) ==> 2SO3 (g),
  • followed by, indirectly, SO3 + H2O ==> H2SO4
  • where the sulfur trioxide combines with water to form sulfuric acid..
 Double decomposition Double decomposition is chemical reaction that takes place between two compounds, in which the first part of one compound combines with the second part of another compound. The bits left over combine to form the second compound. One of the compounds is usually insoluble.
  • e.g. if you mix solutions of sodium chromate with lead nitrate you get a yellow precipitate of lead chromate and sodium nitrate is left in solution.
  • sodium chromate + lead nitrate ==> lead chromate + sodium nitrate
  • Na2CrO4(aq) + Pb(NO3)2(aq) ==> PbCrO4(s) + 2NaNO3(aq)
  • This is the yellow pigment Van Gogh used in his paintings and you see it as the road markings you don't park on!
Catalytic Conversion (car exhaust)
  • One way of reducing pollutants from a car exhaust is to use transition metal catalysts (Pt+Rh set on a heat resistant base = the catalytic converter). A platinum/rhodium catalyst converts toxic carbon monoxide and nitrogen monoxide gases into harmless carbon dioxide and nitrogen gases.
  • 2CO (g) + 2NO (g) == Pt/Rh ==> 2CO2 (g) + N2 (g) 
  • Combining an organic carboxylic acid with an alcohol, produces an pleasant smelling ester (which are used in perfumes and flavourings) and water.
    • e.g. ethanoic acid + ethanol (c) doc b ethyl ethanoate + water
    • ethanoic acid (c) doc b + ethanol (c) doc b  the ester ethyl ethanoate (c) doc b + H2O
    • Its an equilibrium, 2/3 rds conversion, and the reaction is catalysed by a few drops of concentrated sulphuric acid.
  • 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. This means rusting is an oxidation reaction.
  • iron + oxygen + water ==> hydrated iron(III) oxide
  • 4Fe(s) + 3O2(g) + xH2O(l) ==> 2Fe2O3.xH2O(s)
  • i.e. rust is an orange-brown solid of 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!).
  • For more details of the chemistry of rusting and its prevention go to the corrosion section on the Metal Reactivity Series page.
  • A substitution reaction is where one part of a molecule is replaced by something else.
  • e.g. when chlorine reacts with methane, a hydrogen atom in methane is replaced by a chlorine atom from the chlorine molecule.
  • methane + chlorine ==> chloromethane + hydrogen chloride
  • CH4 + Cl2 ==> CH3Cl + HCl
  • An addition reaction is when one molecule adds to another molecule resulting in a single product e.g.
  • ethene + bromine ==> 1,2-dibromoethane
  • C2H4 + Br2 ==> C2H4Br2
  • ethene doc b oil notes bromine doc b oil notes 1,2-dibromoethane


Chemical Synthesis
  • The word synthesis in chemistry, usually means to build up a larger molecules from simpler molecules or atoms ...

  • e.g. iron sulphide is synthesised by heating together iron and sulphur ...

    • Fe(s) + S(s) ==> FeS(s)

  • or The Haber Synthesis of ammonia from its constituent elements  ...

    • nitrogen + hydrogen ==> ammonia

    • N2(g) + 3H2(g) ==> 2NH3(g)

  • In organic chemistry, usually at more advanced level than GCSE, complex molecules are made by doing a whole sequence of steps, modifying one molecule to another.

  • This would be described as a multi-stage synthesis.



Deliquescent and Hygroscopic

Two overlapping terms involving physical changes rather than chemical changes


  • A substance is described as showing 'deliquescence' if it absorbs so much water from the air it forms a solution i.e. dissolves in the absorbed water.

  • Examples: the salts calcium chloride and potassium fluoride.

  • It is the extreme case of being hygroscopic (below).

  • At a higher level, technically it has a lower water vapour pressure than the surrounding air, so there is a net gain in water. 


  • A substance is described as hygroscopic if it absorbs water vapour from the air to form a damp or moist solid or even a solution - see deliquescence above.

  • These substances can be used as drying agents (dessicants) for air or liquids.

  •  e.g. the anhydrous salts calcium chloride and sodium sulphate both readily absorb water to form the hydrated salt.


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