This is a catalogue of many of the equations that you could be expected to know for the Edexcel Chemistry syllabus A 1530.

You need to have learned the formulae of the ions and the rules for forming covalent compounds before you can understand equations properly.

The reactions are listed in syllabus order. For some (for example the displacement reactions of halide ions by more reactive halogens) I have given an example that can easily be extended to the other reactions necessary.

There are some categories that could have thousands of examples – exothermic reactions or redox reactions, say – so I have given enough to enable the principles to be shown.

In most cases state symbols are included. However for some reactions or half-reactions this is difficult – thus the half-equations for the electrolysis of aluminium oxide in molten cryolite do not lend themselves to state symbols.

The only way to learn the reactions is to write them out from memory – if you do so you will not only understand your Chemistry much better since patterns will emerge, but you will also score significantly higher marks!

Reaction with water: all the group 1 metals behave similarly.

2Na(s) + 2H₂O(l) à 2NaOH(aq) + H₂(g)

The reactivity increases down the group, with increasing atomic size. The solution produced is alkaline.

Halogens react directly with many metals on heating to give the metal halide. In most cases the reactions are similar for chlorine, bromine or iodine.

2Na(s) + Cl₂(g) à 2NaCl(s)

Mg(s)  + Cl₂(g) à MgCl_2­(s)

2Fe(s) + 3Cl₂(g) à 2FeCl₃(s)
 

Hydrogen burns in halogens to give the hydrogen halide; the reaction is important in producing hydrogen chloride and hence, by dissolving this in water, hydrochloric acid:

H₂(g) + Cl₂(g) à  2HCl(g)
 

The halogens are oxidising agents; chlorine is stronger than bromine which is stronger than iodine. Chlorine is used commercially to extract bromine by oxidation of bromide ions in seawater:

2Br ^– (aq) + Cl₂(g) à Br₂(aq) + 2Cl ^– (aq)

Both chlorine and bromine will oxidise iodide ions to iodine:

2I ^– (aq)  + Cl₂(g) à  I₂(aq) + 2Cl ^– (aq)

2I ^– (aq) + Br₂(aq) à I₂(aq) + 2Br ^– (aq)
 

Transition metals or their compounds are used as catalysts for some important industrial processes:

Ø      Iron in the Haber process for the manufacture of ammonia:

N₂(g) + 3H₂(g) à 2NH₃(g)

Ø      Vanadium pentoxide in the Contact Process for the manufacture of sulphur trioxide and hence sulphuric acid:

2SO₂(g) + O₂(g) à 2SO₃(g)
 

The solution contains the ions Na⁺, H⁺, Cl ^– , OH ^– .

At the anode chloride ions are discharged in preference to hydroxide ions, and chlorine is liberated:

2Cl ^– (aq) + 2 e ^–  à  Cl₂(g)

At the cathode hydrogen ions are discharged in preference to sodium ions, and hydrogen is liberated:

2H⁺(aq) + 2e ^– à  H₂(g)

The sodium and hydroxide ions remain in solution.

Calcium oxide decomposes on strong heating:

CaCO₃(s) à  CaO(s)  +  CO₂(g)

This reaction is important in cement manufacture and in the blast furnace.

Calcium oxide reacts exothermically with water; steam will be evolved:

CaO(s)  +  H₂O(l) à  Co(OH)₂ (aq)

The solution produced is limewater.

Both calcium oxide and calcium hydroxide are bases and being relatively cheap are used to neutralise soil acidity.

The reaction of a substance with oxygen is obviously oxidation:

2Mg(s) + O₂(g) à 2MgO(s)

CH₄(g) + 2O₂(g) à CO₂(g) + 2H₂O(l)

C₆H₁₂O₆(s) + 6O₂(g) à 6CO₂(g) + 6H₂O(l)

Reduction can be seen as the inverse process where oxygen is removed:

CuO(s) + H₂(g) à Cu(s) + H₂O(g) on heating

Fe₂O₃(s) + 3CO(g) à 2Fe(l) + 3CO₂(g) in blast furnace
 

At the anode of the electrolytic cell, oxide ions are oxidised to oxygen which immediately reacts with the graphite anode, this having to be replaced from time to time:

2O^{2 –}  à O₂ + 4e ^–

C + O₂ à CO₂
 

Aluminium ions are reduced to molten aluminium at the cathode:

Al³⁺ +  3 e ^–  à Al
 

A wide variety of oxides of less reactive metals can be reduced by carbon or carbon monoxide:

Fe₂O₃ + 3C à 2Fe + 3CO

Fe₂O₃ + 3CO à 2Fe + 3CO₂

CuO + C à Cu + CO

CuO + CO à Cu + CO₂
 

For production of iron:

Coke is added to produce heat and the reducing agent, CO; these reactions occur near the air inlet.

C + O₂ à CO₂     + heat; an exothermic reaction.

CO₂ + C à 2CO   an endothermic reaction
 

CO reduces the iron ore; there is also some reduction with unburnt carbon higher up the furnace. The reaction with CO is endothermic and uses about half the energy in the furnace:

Fe₂O₃ +  3CO à  2Fe  +  3CO₂

Fe₂O₃ +  3C à  2Fe  +  3CO
 

The main impurity in the ore is silica, SiO₂. This is acidic and reacts with CaO (a base) to give molten slag CaSiO₃. CaO is produced from limestone in an endothermic reaction by heating :

CaCO₃ à  CaO  + CO₂

CaO + SiO₂ à CaSiO₃
 

Combustion of alkanes:

CH₄(g)  +  2O₂(g)  à  CO₂(g)  +  2H₂O(g)  +  heat

C₈H₁₈(l)  +  12½ O₂(g)  à  8CO₂(g)  +  9H₂O(g)  +  heat
 

Incomplete combustion in limited oxygen always gives several reactions; some carbon dioxide is always formed, together with carbon monoxide and possibly carbon:

CH₄(g)  +  1 ½ O₂(g)  à  CO(g)  +  2H₂O(g)

CH₄(g)  +  O₂(g) à  C(s) + 2H₂O(g)
 

Test for CO₂: forms a white precipitate with limewater.

CO₂(aq)  +  Ca(OH)₂(aq)  à  CaCO₃(s) + H₂O(l)
 

Cracking alkanes gives alkenes and smaller alkanes; in the lab aluminium oxide can be used as the catalyst, though the products are rather random. In industry conditions are chosen to give the products desired. A possible reaction is:

C₁₂H₂₆(g) à  C₈H₁₈(g) + 2 CH₂=CH₂ (g)
 

Alkenes react with bromine water to give a colour change from orange to colourless (NOT clear!):

CH₂=CH₂ (g)  +  Br₂(aq)  à  BrCH₂CH₂Br (l)
 

Alkenes can polymerise to give poly(alkenes) with very long carbon chains – 10,000 or more carbon atoms.

n CH₂=CH₂ à  (CH₂=CH₂)_n
 

Exothermic reactions give out heat; most reactions are exothermic. The examples can be extended indefinitely!

CH₄(g)  +  2O₂(g)  à  CO₂(g)  +  2H₂O(g)  +  heat

C₈H₁₈(l)  +  12½ O₂(g)  à  8CO₂(g)  +  9H₂O(g)  +  heat

C₆H₁₂O₆  +  6O₂ à  6CO₂ + 6H₂O  (respiration)

N₂(g)  +  3H₂(g)  à  2NH₃(g)  (Haber process)
 

Endothermic reactions take in heat:

N₂(g) + O₂(g)  à 2NO(g)  (in internal combustion engines)

6CO₂ + 6H₂O  à  C₆H₁₂O₆  +  6O₂ (photosynthesis)

Photosynthesis is an endergonic process – it takes in light rather than heat as its energy source, and it isn’t a single reaction so is a process. 

Haber process:

N₂(g)  +  3H₂(g)  à  2NH₃(g) 
 

Production of ammonium nitrate fertiliser:

NH₃(aq)  +  HNO₃(aq)  à  NH₄NO₃(aq)
 

Insoluble salts can be made using precipitation reactions – these also form the basis for qualitative analysis (below). The common insoluble salts are:

Chlorides: silver and lead

Ag⁺(aq)  +  Cl ^– (aq) à  AgCl(s)

Pb²⁺(aq)  +  2Cl ^– (aq) à  PbCl₂(s)

Sulphates: lead, barium and calcium:

Pb²⁺(aq)  +  SO₄^{2 –} (aq) à  PbSO₄(s)

Ba²⁺(aq)  +  SO₄^{2 –} (aq) à  BaSO₄(s)

Ca²⁺(aq)  +  SO₄^{2 –} (aq) à  CaSO₄(s)

Since CaSO₄ is not totally insoluble (hard water contains it) the latter reaction requires concentrated solutions.

Carbonates and hydroxides: all of these are insoluble apart from compounds of ammonium and group 1 cations. The examples given are of one specific reaction for each and a general reaction for each:

Ca²⁺(aq)  +  CO₃^{2 –} (aq)  à  CaCO₃(s)

M²⁺(aq)  +  CO₃^{2 –} (aq)  à  MCO₃(s)

Cu²⁺(aq)  +  2OH ^– (aq)  à  Cu(OH)₂(s)

M^x+(aq)  +  xOH ^– (aq)  à  M(OH)_x(s)
 

If lit burns with a squeaky pop if mixed with air, which it usually is when tested in the lab:

2H₂(g) + O₂(g)  à  2H₂O(g)
 

Gives a white precipitate when bubbled through limewater; prolonged bubbling causes the precipitate to disappear:

CO₂(aq)  +  Ca(OH)₂(aq) à  CaCO₃(s)  +  H₂O(l)

CaCO₃(s)  +  CO₂(aq)  +  H₂O(l) à  Ca(HCO₃)₂(aq)
 

Ø      Turns moist blue litmus red;

Ø      gives a white precipitate if a drop of silver nitrate solution is held in the gas:

AgNO₃(aq)  +  HCl(aq)  à  AgCl(s)  +  HNO₃(aq)

Ø      forms a white smoke if an open bottle of conc ammonia is held near to it:

NH₃(g)  +  HCl(g)  à  NH₄Cl(s)
 

Ø      Turns moist red litmus blue – it is the only common alkaline gas.

Ø      Forms a white smoke if an open bottle of conc HCl is held near to it – see above.
 

H⁺(aq):

Ø      turns blue litmus red

Ø      gives CO₂ (test with limewater) with sodium carbonate or sodium hydrogen carbonate

Ø      liberates hydrogen from magnesium ribbon

Carbonates give CO₂ with dilute hydrochloric acid and fizz vigorously:

CO₃^{2 –} (aq)  +  2H⁺(aq) à  CO₂(g)  +  H₂O(l)
 

Sulphites give SO₂ on warming with dilute hydrochloric acid; the SO₂ turns orange potassium dichromate(VI) paper green.

SO₃^{2 –} (aq)  +  2H⁺(aq) à  SO₂(g)  +  H₂O(l)
 

Sulphates give a white precipitate of barium sulphate with barium chloride or barium nitrate solution:

Ba²⁺(aq)  +  SO₄^{2 –} (aq)  à  BaSO₄(s)
 

Halide (chloride, bromide or iodide) gives a precipitate with silver nitrate solution, which may or may not dissolve in ammonia solution:

                                                                      Ag⁺(aq)  +  X ^– (aq)  à  AgX(s)

Soap is the sodium salt of various long-chain organic acids, e.g. stearic acid gives sodium stearate, C₁₇H₃₅COONa.  Scum is formed by precipitation of the insoluble calcium salts of these acids:

2C₁₇H₃₅COO ^– (aq)  +  Ca²⁺(aq)  à  (C₁₇H₃₅COO)₂Ca (s)
 

Hardness is of two varieties – usually hard water has both:

Ø      Temporary: formed by reaction between carbon dioxide dissolved in rainwater and limestone, chalk or marble, i.e. calcium carbonate, giving calcium hydrogen carbonate in solution:

CO₂(aq) + H₂O(l) + CaCO₃(s) à  Ca(HCO₃)₂(aq)

Ø      This reaction is reversed when the water is boiled, giving a precipitate of  calcium carbonate as fur or scale:

   Ca(HCO₃)₂(aq)   à   CO₂(g) + H₂O(l) + CaCO₃(s) 

Ø      Permanently hard water is caused by water dissolving the sparingly-soluble rock gypsum, which is calcium sulphate:

CaSO₄(s)  +  aq  à  Ca²⁺(aq)  +  SO₄^{2 – m}(aq)
 

Water can be softened

Ø      By reacting it with sodium carbonate which precipitates the calcium ions as calcium carbonate:

Ca²⁺(aq)  +  CO₃^{2 –} (aq)   à   CaCO₃(s)

Ø      By reacting it with an ion exchange resin, which exchanges sodium ions on its surface for calcium ions in solution:

2 resin-Na⁺(s)  +  Ca²⁺(aq)  à  (resin)₂-Ca²⁺(s)  +  Na⁺(aq)

Ø      By distillation – this is too expensive for large-scale use.
 

Sulphuric acid is made from sulphur or sulphide ores; firstly sulphur dioxide is obtained by burning liquid sulphur

S(l)  +  O₂(g)  à  SO₂(g)

or by roasting an ore such as iron pyrites FeS₂ in air:

4FeS₂(s)  + 11 O₂(g)  à 2 Fe₂O₃(s)  +  8 SO₂(g)
 

Sulphur dioxide is oxidised to sulphur dioxide using a temp of around 430^oC and a catalyst of vanadium(V) oxide:

SO₂(g)  + ½ O₂(g)  à  SO₃(g)
 

Sulphur trioxide is dissolved in 98% sulphuric acid and reacts with the water to make a more concentrated acid which is then diluted: the net effect is

SO₃(g) +  H₂O(l)  à  H₂SO₄(l)
 

Ethanol is produced by fermentation; the overall process from glucose is:

C₆H₁₂O₆ à  2CH₃CH₂OH +  2CO₂
 

Ethanol for solvent use is produced by hydration of ethane at 67atm, 350^oC, and a catalyst of H₃PO₄:

CH₂=CH₂(g)  +  H₂O(g)  à  CH₃CH₂OH(g)
 

Ethanol is used as a fuel:

CH₃CH₂OH(l)  +  3 O₂(g)  à  2CO₂(g)  +  3H₂O(l)  +  heat
 

Ethanol slowly oxidises in air to give ethanoic acid – this reaction spoils wine. All the states are (aq) since this reaction happens in aqueous solution/

CH₃CH₂OH(aq)  +  O₂(aq)  à  CH₃COOH(aq) +  H₂O(l)
 

Ethanoic acid is a typical acid although since it is only about 1% dissociated at room temperature its reactions are much slower than those of, say, hydrochloric acid of the same concentration.

Ø      With reactive metals:

Mg(s) + 2CH₃COOH(aq) à Mg(CH₃COOH)₂ (aq)  +  H₂(g)

Ø      With carbonates:

Na₂CO₃(aq) + 2CH₃COOH(aq) à 2CH₃COONa(aq) + H₂O(l) + CO₂(g)

Ø      With sodium hydroxide:

NaOH(aq) + CH₃COOH(aq)  à  CH₃COONa(aq) + H₂O(l)
 

Ethanoic acid reacts with ethanol to give an ester, ethyl ethanoate:

CH₃COOH + CH₃CH₂OH à  CH₃COOCH₂CH₃ + H₂O