|Home||Y10||Y11||AS Units||A2 Units|
Salts are prepared by reacting an acid with a metal or a base, such as a metal carbonate, hydroxide or oxide.
The acid provides the non-metal ion for the salt, e.g. chloride or sulphate or nitrate ions.
The metal or base provides the metal ion for the salt, e.g. sodium or copper.
The method used to produce a particular salt depends on two factors -
Since all acids are aqueous solutions the acid needed does not directly affect the method of preparation.
The following table summarises the solubilities of the various bases and salts,
|compound||metal compound soluble in water||insoluble in water|
|hydroxide||sodium, potassium, ammonium, calcium||all others|
|oxide||sodium, potassium, calcium (all dissolve to give hydroxide)||all others|
|carbonate||sodium, potassium, ammonium||all others|
|sulphate||all others||lead, barium|
|chloride||all others||silver, lead|
There are three main methods of preparing salts -
(i) for a metal or insoluble base reacting with an acid to produce a soluble salt filtration is used, e.g. reacting copper(II) oxide with sulphuric acid to make copper(II) sulphate.
(ii) for a soluble base reacting with an acid to produce a soluble salt titration is used, e.g. reacting sodium hydroxide with hydrochloric acid to make sodium chloride.
(iii) for an insoluble base reacting with an acid to produce an insoluble salt a two stage process involving filtration and then precipitation is used, e.g. reacting lead(II) oxide firstly with nitric acid to produce lead(II) nitrate and then reacting the lead(II) nitrate with aqueous iodide ions to produce lead(II) iodide.
A solubility curve is a graph showing the solubility of a particular compound in a solvent at various temperatures.
For example -
The data for this graph can be gathered by performing a series of experiments at different temperatures.
A beaker of water is heated to the required temperature and an excess of solid is added to the eater, with stirring. The excess is filtered off, dried and weighed and the mass dissolved can be calculated by subtraction.
The graph can be used to calculate how much solid is needed to make a saturated solution at a particular temperature.
For example -back to top
Experimental sheet for the preparation of copper(II) sulphate.
As shown by the diagrams above, the first stage is the addition of black copper(II) oxide to sulphuric acid. Mild heating is required for a full reaction to occur; however, care must be taken to ensure that the acid does not boil as this would be a great safety hazard.
The copper(II) oxide is added until no more visible reaction can be seen, i.e. the base no longer dissolves and a black solid is seen in the blue solution.
The mixture is then filtered (stage 2 above) to remove the excess black solid and leave a clear blue solution in the evaporating bowl. If the blue solution is heated gently, to remove some of the water and allowed to cool down slowly, crystals will appear. The slower this crystallization is allowed to occur, the larger the crystals that will be produced.
Try balancing the equation for this reaction,
Experimental sheet for the preparation of sodium chloride.
N.B.: This method of preparation is used when any and all sodium or potassium salts are made.
The problem with the reaction of a soluble base reaction with acid is that once all the acid has reacted any excess base will not be visible. This problem is overcome by adding a third chemical into the reaction mixture called an indicator.
Indicators are chemicals that change colour with a change in pH. So, if the indicator is added to an acid it will be one colour. As base is added the pH of the solution is raised until, once all the acid has reacted, i.e. been neutralized, and the base is now in excess, the indicator changes colour.
Exemplar indicators -
|Indicator||colour in acid||colour in base|
The skill involved in this technique is to add the base to the acid slowly enough so that the indicator just changes colour with one drop of excess base.
Specialist glassware is used to allow extreme precision in the volume measurements for this technique. The acid required to begin with is measured using a pipette into a conical flask and a few drops of an indicator is added to the flask.
A clamp stand is set up with a burette attached, see diagram above, and the burette is filled up with the solution of base.
The base is added from the burette slowly until the indicator changes colour. The volume of base required can be read from the markings on the side of the burette.
The experiment can be run a number of times to gain a more accurate average value for the volume of base that needs to be added to neutralise the acid used.
Once this accurate value has been determined, the experiment can be run one last time without any indicator. The volume of base that is required can be added carefully from the burette to produce a colourless solution, instead of the coloured solutions produced when indicator is added.
This clear colourless solution can be heated to encourage crystals to form and then left to cool down slowly so that large crystals can form.
Try balancing the equation for this reaction,
Experimental sheet for the preparation of lead(II) iodide.
There is no direct method of preparing an insoluble salt from an insoluble base. The problem us that as the base is added the salt is produced so filtration would yield a mixture of two solids. Titration won't work either as the base is not soluble in water.
Using the preparation of lead(II) iodide as an example, the solution is a two stage process. The first stage is the same as the method for making copper(II) sulphate (shown above). Lead(II) oxide is added to hot nitric acid and when an excess of oxide is present the mixture is filtered.
This gives a filtrate of lead(II) nitrate solution. An equal volume of potassium iodide solution is then added and the resulting mixture filtered to give solid lead(II) iodide as the residue.back to top
Experimental sheet for chemical ion tests.
When metal ions are reacted with hydroxide ions a displacement reaction generally occurs, since most metal hydroxides are insoluble in water. Group I hydroxides, e.g. sodium hydroxide, are all soluble in water and so if a solution of sodium hydroxide is mixed with a solution of another metal ion a precipitate is formed.
The transition metal hydroxides produced tend to be uniquely coloured and so they allow easy identification of the particular metal ion that is present, e.g. iron(II) hydroxide is green, iron(III) hydroxide is red and copper(II) hydroxide is blue.
Below is a table summarizing the precipitates formed with common GCSE metal ions -
|Cation||Reaction with NaOH(aq)||Reaction with NH4OH(aq)|
|aluminium, Al3+(aq)||white precipitate, soluble in excess||white precipitate, insoluble in excess|
|ammonium, NH4+(aq)||ammonia (SMELL !) produced on heating|
|calcium, Ca2+(aq)||white precipitate, insoluble in excess||no precipitate|
|copper, Cu2+(aq)||light blue precipitate, insoluble in excess||light blue precipitate, soluble in excess giving a dark blue solution|
|iron(II), Fe2+(aq)||green precipitate, insoluble in excess||green precipitate, insoluble in excess|
|iron(III), Fe3+(aq)||red-brown precipitate, insoluble in excess||red-brown precipitate, insoluble in excess|
|zinc, Zn2+(aq)||white precipitate, soluble in excess||white precipitate, soluble in excess|
Exemplar equation -
2NaOH(aq) + CuSO4(aq) → Na2SO4(aq) + Cu(OH)2(s)
These equations can also be represented only by the reacting ions, simplifying the equation -
2-OH(aq) + Cu2+(aq) → Cu(OH)2(s)
As can be seen in the table above if ammonium hydroxide, NH4OH, is used instead of sodium hydroxide some differences occur. Ammonium hydroxide is a weaker base than sodium hydroxide, i.e. there are fewer hydroxide ions in solution.
For calcium ions this means that there aren't enough hydroxide ions to form any calcium hydroxide, so no precipitate is seen.
The addition of excess hydroxide ions allows more differentiation between ions, such as aluminium and zinc. Normally hydroxides are bases, i.e. they react only with acids; however, some metal hydroxides are amphoteric, i.e. they react with both acids and bases. This is the case with aluminium hydroxide and zinc hydroxide. If excess sodium hydroxide is added to the precipitate a reaction occurs that forms a soluble product.
Exemplar equation -
sodium hydroxide + aluminium hydroxide → sodium aluminate
3NaOH(aq) + Al(OH)3(s) → Na3Al(OH)6(aq)
Ammonium hydroxide is too weak a base to allow this reaction to occur, so the aluminium hydroxide does not dissolve in excess ammonium hydroxide.
With zinc hydroxide and copper(II) hydroxide the precipitates do dissolve in excess ammonium hydroxide, not because of an acid-base reaction, but because the molecules of ammonia surround the metal ions in solution, making them soluble. This is an idea that is more fully dealt with in A2 chemistry (see transition elements in Trends and Patterns).
Try balancing the equations for these precipitation reactions,
There is no particular pattern for testing non-metal ions. The halide ions, chloride, bromide and iodide, are tested for by adding nitric acid followed by aqueous silver nitrate and noticing the colour of the precipitate formed (see halogens in Year 10 GCSE Chemistry).
Carbonate ions react with acids to liberate carbon dioxide gas, which can be tested for by passing the gas through limewater.
Sulphate ions will undergo a displacement reaction when added to acidified aqueous barium chloride, BaCl2. A white precipitate fo barium sulphate is formed.
Ammonium ions are also non-metal ions and are tested for by warming the ammonium compound with aqueous sodium hydroxide, or some other strong base. This produces ammonia gas, which can be tested for by placing damp red litmus paper in the gas. If it turns blue then a basic gas has been given off and that must be ammonia.
Exemplar ionic equation -
NH4+(aq) + -OH(aq) → NH3(g) + H2O(l)
Below is a table summarizing the observations for the common non-metal ion tests -
|carbonate, CO32-(aq)||CO2 liberated when reacted with dilute acids|
|chloride, Cl-(aq)||gives white precipitate with acidified silver nitrate, AgNO3(aq)|
|iodide, I-(aq)||gives yellow precipitate with acidified lead nitrate, Pb(NO3)2(aq)|
|nitrate, NO3-(aq)||ammonia, NH3, liberated on heating with NaOH(aq) and Al foil (see tests on ammonia gas below)|
|sulphate, SO42-(aq)||gives white precipitate with acidified barium chloride, BaCl2(aq)|
Below is a table summarizing the various gas tests encountered in GCSE -
|ammonia, NH3||turns damp red litmus paper blue (+ SMELL !)|
|carbon dioxide, CO2||gives a white precipitate with limewater|
|chlorine, Cl2||bleaches damp litmus paper|
|hydrogen, H2||explodes with a 'pop' with a lighted splint|
|oxygen, O2||relights a glowing splint|
written by Dr Richard Clarkson : © Saturday, 1 November 1997
Updated : Monday 2nd April, 2012
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