Relating Observations to Chemical Equations
- Chemical equations give information about the reaction that is taking place
- Balanced equations show the number of particles participating in the reaction and the number of products being formed
- Balanced equations can be used to calculate the number of moles involved in reactions
- Balanced equations can, also, be used to calculate masses and volumes involved in reactions
- Ionic equations only show the reacting particles
- Ionic equations allow you to identify spectator ion
Types of reaction
- Chemical equations can be used to determine the type of reaction taking place
Displacement reactions
Br2 (aq) + 2KI (aq) → I2 (aq) + 2KBr (aq)
- In this reaction, the more reactive bromine displaces the less reactive iodide in potassium iodide
- This can also be seen in the ionic equation for the reaction
Br2 (aq) + 2I- (aq) → I2 (aq) + 2Br- (aq)
Exam Tip
The use of chemical equations can help identify risks and hazards in the reaction and suggest appropriate precautions where necessary
For example, the use of aqueous bromine in the above example should suggest the potential use of a fume cupboard and nitrile gloves because:
- Bromine liquid is toxic, corrosive and harmful to the environment
- Bromine water with a concentration of 0.2 mol dm3 is corrosive
- With a concentration of between 0.06 mol dm3 and 0.2 mol dm3, bromine water is an irritant
- Only below concentrations of 0.06 mol dm3 is bromine water considered a low hazard
Neutralisation reactions
- These can be identified by the presence of reactant acids and bases as well as the formation of a neutral salt solution and water (and sometimes other compounds such as carbon dioxide)
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
Na2CO3 (aq) + 2HNO3 (aq) → 2NaNO3 (aq) + H2O (l) + CO2 (g)
- The ionic equations can more clearly demonstrate the neutralisation of an acid and a base:
H+ (aq) + OH- (aq) → H2O (l)
2H+ (aq) + CO32- (aq) → H2O (l) + CO2 (g)
Exam Tip
For neutralisation reactions, the main hazards are linked to:
- The concentration of the acid
- The type of acid
- This could mean strong, e.g. HCl, or weak, e.g. CH3COOH
- This could also mean monoprotic, e.g. HNO3, diprotic, e.g. H2SO4, or triprotic, e.g. H3PO4
- The concentration of the base
- The strength of the base
- The physical state of the base, e.g. NaOH (s) is arguably considered more corrosive than a high concentration solution of NaOH (aq)
Precipitation reactions
- These are shown by the reaction of two aqueous solutions to form products which include one solid
BaCl2 (aq) + Na2SO4 (aq) → BaSO4 (s) + 2NaCl (aq)
- The ionic equation shows the precipitation reactions more clearly as there are no other products considered
Ba2+ (aq) + SO42- (aq) → BaSO4 (s)
Exam Tip
A precipitation reaction is a clear example of where consideration for further practical procedures is most obvious
The formation of a solid product should tell you that any purification of the product should include filtering or decanting as a minimum
