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!

 

 

 
Alkali metals, group 1.
 
 

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

2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

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

 

 

 
The halogens, group 7.
 
 

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) 2NaCl(s)

Mg(s)  + Cl2(g) MgCl2(s)

2Fe(s) + 3Cl2(g) 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)   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) Br2(aq) + 2Cl (aq)

Both chlorine and bromine will oxidise iodide ions to iodine:

2I (aq)  + Cl2(g)   I2(aq) + 2Cl (aq)

2I (aq) + Br2(aq) I2(aq) + 2Br (aq)
 

 

 

 
Transition metals
 
 

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

      Iron in the Haber process for the manufacture of ammonia:

N2(g) + 3H2(g) 2NH3(g)

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

2SO2(g) + O2(g) 2SO3(g)
 

 

 

 
Chemicals from salt
 
 
Electrolysis of concentrated sodium hydroxide solution is used to manufacture chlorine, hydrogen and sodium hydroxide.

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    Cl2(g)

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

2H+(aq) + 2e   H2(g)

The sodium and hydroxide ions remain in solution.

 

 

 
Chemicals from calcium carbonate
 
 

Calcium oxide decomposes on strong heating:

CaCO3(s)   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)   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.

 

 

 
Extraction and uses of metals.
 
 

The reaction of a substance with oxygen is obviously oxidation:

2Mg(s) + O2(g) 2MgO(s)

CH4(g) + 2O2(g) CO2(g) + 2H2O(l)

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

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

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

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

 
Extraction of aluminium:
 
 

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:

2O2   O2 + 4e

C + O2 CO2
 

 

Aluminium ions are reduced to molten aluminium at the cathode:

Al3+ +  3 e   Al
 

 
Reduction of ores using carbon monoxide or carbon: making metals useful
 
 

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

 

Fe2O3 + 3C 2Fe + 3CO

Fe2O3 + 3CO 2Fe + 3CO2

CuO + C Cu + CO

CuO + CO 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 CO2     + heat; an exothermic reaction.

CO2 + C 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   2Fe  +  3CO2

Fe2O3 +  3C   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   CaO  + CO2

CaO + SiO2 CaSiO3
 

 

 

 
Useful products from crude oil
 
 

Combustion of alkanes:

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

C8H18(l)  +  12 O2(g)    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 O2(g)    CO(g)  +  2H2O(g)

CH4(g)  +  O2(g)   C(s) + 2H2O(g)
 

 
Note in all cases the water is shown as (g) because initially that is what it will be. The hydrocarbon is shown in its normal state at r.t., though in fact it is only the vapour of the hydrocarbon that burns.
 
 

Test for CO2: forms a white precipitate with limewater.

CO2(aq)  +  Ca(OH)2(aq)    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)   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)    BrCH2CH2Br (l)
 

 

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

n CH2=CH2   (CH2=CH2)n
 

 

 

 
Energy transfers accompanying reactions
 
 

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

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

C8H18(l)  +  12 O2(g)    8CO2(g)  +  9H2O(g)  +  heat

C6H12O6  +  6O2   6CO2 + 6H2(respiration)

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

 

Endothermic reactions take in heat:

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

6CO2 + 6H2  C6H12O6  +  6O2 (photosynthesis)

Photosynthesis is an endergonic process it takes in light rather than heat as its energy source, and it isnt a single reaction so is a process. 

 

 

 
Changing materials the environment
 
 

Haber process:

N2(g)  +  3H2(g)    2NH3(g) 
 

 

Production of ammonium nitrate fertiliser:

NH3(aq)  +  HNO3(aq)    NH4NO3(aq)
 

 

 

 
Preparing and analysing
 
 

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)   AgCl(s)

Pb2+(aq)  +  2Cl (aq)   PbCl2(s)

Sulphates: lead, barium and calcium:

Pb2+(aq)  +  SO42 (aq)   PbSO4(s)

Ba2+(aq)  +  SO42 (aq)   BaSO4(s)

Ca2+(aq)  +  SO42 (aq)   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)  +  CO32 (aq)    CaCO3(s)

M2+(aq)  +  CO32 (aq)    MCO3(s)

Cu2+(aq)  +  2OH (aq)    Cu(OH)2(s)

Mx+(aq)  +  xOH (aq)    M(OH)x(s)
 

 

 

 
Tests for gases
 
 
H2

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

2H2(g) + O2(g)    2H2O(g)
 

 
O2 Relights a glowing splint a good example of the effect of concentration on reaction rate.
 
 
CO2

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

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

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

 
HCl

      Turns moist blue litmus red;

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

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

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

NH3(g)  +  HCl(g)    NH4Cl(s)
 

 
NH3

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

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

 
SO2 If moist orange potassium dichromate paper is held in SO2 the paper turns green. Alternatively the gas may be bubbled through a solution of potassium dichromate(VI)  

 

 
Tests for ions
 
 
The tests need to have simple visual results they work because the number of ions that you analyse is quite small. As the number of possible ions increases, so do the complexity of the tests.
 
 

H+(aq):

      turns blue litmus red

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

      liberates hydrogen from magnesium ribbon

 

 
Na+, K+, Ca2+ and Cu2+ all give colours to a flame:
 
  Na+ yellow or orange
  K+ lilac
  Ca2+ brick red or orange-red
  Cu2+ green with blue centre
 
 
The tests do not involve chemical reactions so there are no equations for them.
 
 
Sodium hydroxide solution reacts with solutions of ions as follows:
 
Al3+ Al3+(aq) +  3OH (aq)   Al(OH)3(s) white ppt which dissolves in excess NaOH
Ca2+

Ca2+(aq) +  2OH (aq)   Ca(OH)2(s)

white ppt insoluble in excess NaOH
Cu2+

Cu2+(aq) +  2OH (aq)   Cu(OH)2(s)

blue ppt
Fe2+ Fe2+(aq) +  2OH (aq)   Fe(OH)2(s) dirty green ppt
Fe3+

Fe3+(aq) +  3OH (aq)   Fe(OH)3(s)

foxy-red ppt
NH4+ NH4+ (aq) + OH (aq)   NH3(g)  +  H2O(l) on warming evolves a gas that turns red litmus blue
 
   

Carbonates give CO2 with dilute hydrochloric acid and fizz vigorously:

CO32 (aq)  +  2H+(aq)   CO2(g)  +  H2O(l)
 

 

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

SO32 (aq)  +  2H+(aq)   SO2(g)  +  H2O(l)
 

 

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

Ba2+(aq)  +  SO42 (aq)    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)    AgX(s)

Cl
 
Ag+(aq)  +  Cl (aq)    AgCl(s)
 
white ppt soluble in dilute ammonia
 
Br
 
Ag+(aq)  +  Br (aq)    AgBr(s)
 
cream ppt soluble in conc but not dilute ammonia
 
I

 
Ag+(aq)  +  I (aq)    AgI(s)

 
yellow ppt insoluble in ammonia

 
 
Hydroxide, OH  ions can be tested for using any of the reactions above as well as turning red litmus blue. Note that these tests only show the presence of hydroxide ions, not that the substance is a hydroxide. Any alkaline solution will do the same thing, e.g. a carbonate.  

 

 
Hard water  
   
Hardness in water is due to the presence of Ca2+ and Mg2+ ions. They almost always occur together in the equations that follow magnesium could equally well be used instead of calcium.
 
 

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)    (C17H35COO)2Ca (s)
 

 

Hardness is of two varieties usually hard water has both:

      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)   Ca(HCO3)2(aq)

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

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

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

CaSO4(s)  +  aq    Ca2+(aq)  +  SO42 m(aq)
 

 

Water can be softened

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

Ca2+(aq)  +  CO32 (aq)      CaCO3(s)

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

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

 

 

 
Sulphuric acid  
   

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

S(l)  +  O2(g)    SO2(g)

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

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

 

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

SO2(g)  + O2(g)    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)    H2SO4(l)
 

 

 

 
Organic chemistry
 
 

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

C6H12O6   2CH3CH2OH +  2CO2
 

 

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

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

 

Ethanol is used as a fuel:

CH3CH2OH(l)  +  3 O2(g)    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)    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.

      With reactive metals:

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

      With carbonates:

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

      With sodium hydroxide:

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

 

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

CH3COOH + CH3CH2OH   CH3COOCH2CH3 + H2O
 

 
Homologous series: the names and structures of..
 
 
  .the first four alkanes:
 
   
 

methane

CH4

 
 

ethane

CH3CH3

 
 

propane

CH3CH2CH3

 
 

butane

CH3CH2CH2CH3

 
       
  .the first four alkenes:
 
   
 

ethene

CH2=CH2

 
 

propene

CH2=CHCH3

 
 

but-1-ene

CH2=CHCH2CH3

 
 

pent-1-ene

CH2=CHCH2CH2CH3

 
       
  the first four (primary) alcohols:
 
   
 

methanol

CH3OH

 
 

ethanol

CH3CH2OH

 
 

propan-1-ol

CH3CH2CH2OH

 
 

butan-1-ol

CH3CH2CH2CH2OH

 
   

 

 
  the first four carboxylic acids:
 
   
 

methanoic acid

HCOOH

 
 

ethanoic acid

CH3COOH

 
 

propanoic acid

CH3CH2COOH

 
 

butanoic acid

CH3CH2CH2COOH

 
 

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