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Oxford AQA International A
Level Chemistry
Reactions of Ions in Aqueous Solution
Contents
Reactions of Ions in Aqueous Solution
Reactions of Metal-Aqua Ions
Identifying Transition Metal Ions
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Reactions of Ions in Aqueous Solution
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Metal-Aqua Ions
Metal-aqua complex ions form when soluble transition metal salts dissolve in water
A metal-aqua complex ion is the complex of a central transition metal cation with co-ordinate
bonds to ligands
Dissolving copper(II) sulfate in water
This is often shown as:
CuSO4 (s) + aq → Cu2+ (aq) + SO42- (aq)
But, the copper(II) ions actually form the hexaaqua copper(II) ion
CuSO4 (s) + aq → [Cu(H2O)6]2+ (aq) + SO42- (aq)
The water molecules form co-ordinate bonds to central transition metal cation
A lone pairs on the oxygen molecule provides the electrons for the bond
The hexaaqua copper(II) ion, [Cu(H2O)6]2+ (aq), causes the blue colour of the solution
Other metal-aqua complex ions
Iron(II) salts also form hexaaqua complex ions
Fe(NO3)2 (s) + aq → [Fe(H2O)6]2+ (aq) + 2NO3- (aq)
The hexaaqua iron(II) ion, [Fe(H2O)6] 2+ (aq), causes the green colour of the solution
Two common metal-aqua ions with a 3+ charge are iron(III) and aluminium:
Fe(NO3)3 (s) + aq → [Fe(H2O)6]3+ (aq) + 3NO3- (aq)
The hexaaqua iron(III) ion, [Fe(H2O)6]3+ (aq), causes the yellow-orange colour of the solution
Al2(SO4)3 (s) + aq → 2[Al(H2O)6] 3+ (aq) + 3SO42- (aq)
The hexaaqua aluminium(III) ion, [Al(H2O)6]3+ (aq), causes the lack of colour of the solution
Metal-aqua 2+ and 3+ complex ions
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All of the example hexaaqua ions have 6 water ligands and a 2+ or 3+ charge
Acidity in Metal-Aqua Ions
Typically, when transition metal salts dissolve in water they form solutions that are not neutral
pH table for different transition metal solutions
0.1 mol dm-3 solution
pH
Iron(III) chloride, FeCl3
2.0
Aluminium chloride, AlCl3
3.0
Copper(II) nitrate, Cu(NO3)2
4.0
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Metal-aqua complex ions with a 3+ charge are more acidic than those with a 2+ charge
3+ ions, such as iron(III) and aluminium, are smaller than 2+ ions
This means that they have a higher charge density than +2 ions
The higher charge density attracts the lone pair from a water molecule more strongly
This weakens the O-H bonds
O-H bonds can dissociate into a hydroxide ligand, OH-, and a hydrogen ion, H+
The hydrogen ion causes the resulting solution to be more acidic
The metal ion polarises the water molecules
How metal(III) ions polarise water ligands
Metal(III) ions have a high charge density and polarise water molecules in the hexaaqua complexes
[Fe(H2O)6]3+ (aq) → [Fe(H2O)5(OH)]2+ (aq) + H+ (aq)
The loss of the hydrogen ion results in a pentaaqua 2+ complex ion
The remaining hydroxide ion from the water molecule has a negative charge
This cancels one of the charges on the original 3+ complex ion
The same reaction can be shown with a hydroxonium ion product
[Fe(H2O)6]3+ (aq) + H2O (l) → [Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq)
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These reactions can be called deprotonation reactions
Deprotonation reactions of [Fe(H2O)6]3+ (aq)
This usually occurs in several steps
The first two deprotonations of [Fe(H2O)6]3+ (aq) are:
[Fe(H2O)6]3+ (aq) → [Fe(H2O)5(OH)]2+ (aq) + H+ (aq)
[Fe(H2O)5(OH)]2+ (aq) → [Fe(H2O)4(OH)2]+ (aq) + H+ (aq)
The third deprotonation does not usually occur without the presence of a base
The base removes the third proton
This produces a red-brown precipitate of insoluble hydrated iron(III)hydroxide
[Fe(H2O)4(OH)2] + (aq) + OH- → Fe(H2O)3(OH)3 (s) + H2O (l)
Examiner Tips and Tricks
The splitting of water molecules can be called a hydrolysis reaction
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Reactions of Metal-Aqua Ions
Reactions of Metal-Aqua Ions with Bases
Reactions with Bases
The differences in the chemistry of +2 and +3 aqua ions can be seen in their reactions with:
Hydroxide ions, OHAmmonia, NH3
Carbonate ions, CO22-
Hexaaqua iron(II) reactions
With hydroxide ions
The hexaaqua iron(II) undergoes deprotonation reactions with dilute and excess hydroxide ions
This is a two-step process:
[Fe(H2O)6]2+ (aq) + OH– (aq) → [Fe(H2O)5(OH)]+ (aq) + H2O (l)
[Fe(H2O)5(OH)]+ (aq) + OH– (aq) → Fe(H2O)4(OH)2 (s) + H2O (l)
The green [Fe(H2O)6]2+ solution reacts to form a dark green precipitate of hydrated iron(II) hydroxide,
Fe(H2O)4(OH)2 (s)
The green hydrated iron(II) hydroxide will slowly turn orange-brown
This is due to an oxidation reaction forming hydrated iron(III) hydroxide
There is no further reaction with excess hydroxide ions
With ammonia solution
Ammonia behaves in the same way as sodium hydroxide because it is a base
It removes protons from the water ligands
The overall reaction with ammonia is:
[Fe(H2O)6]2+ (aq) + 2NH3 (aq) → Fe(H2O)4(OH)2 (s) + 2NH4+ (aq)
Again, the green [Fe(H2O)6]2+ solution reacts to form a dark green precipitate of hydrated iron(II)
hydroxide, Fe(H2O)4(OH)2 (s)
The green hydrated iron(II) hydroxide will oxidise to orange-brown hydrated iron(III) hydroxide
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There is no further reaction with excess ammonia
With carbonate ions
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With carbonate ions, iron(II) carbonate precipitates out:
[Fe(H2O)6]2+ (aq) + CO32- (aq) → FeCO3 (s) + 6H2O (l)
The green [Fe(H2O)6]2+ solution reacts to form a green precipitate of iron(II) carbonate, FeCO3 (s)
Hexaaqua copper(II) reactions
With hydroxide ions
The hexaaqua copper(II) undergoes deprotonation reactions with dilute and excess hydroxide ions
This is a two-step process:
[Cu(H2O)6]2+ (aq) + OH– (aq) → [Cu(H2O)5(OH)]+ (aq) + H2O (l)
[Cu(H2O)5(OH)]+ (aq) + OH– (aq) → Cu(H2O)4(OH)2 (s) + H2O (l)
The blue [Cu(H2O)6]2+ solution reacts to form a blue precipitate of copper(II) hydroxide, Cu(H2O)4(OH)2
(s)
There is no further reaction with excess hydroxide ions
With ammonia solution
Initially, ammonia behaves in the same way as sodium hydroxide because it is a base
It removes protons from the water ligands
This reaction with ammonia is
[Cu(H2O)6]2+ (aq) + 2NH3 (aq) → Cu(H2O)4(OH)2 (s) + 2NH4+ (aq)
A further reaction occurs with excess ammonia
Ammonia is a stronger ligand than water
So, ammonia partially substitutes for water
Cu(H2O)4(OH)2 (s) + 4NH3 (aq) → [Cu(NH3)4(H2O)2]2+ (aq) + 2OH- (aq) + 2H2O (l)
The blue Cu(H2O)4(OH)2 precipitate reacts to form a deep blue solution of
dihydroxytetraaminecopper(II), [Cu(NH3)4(H2O)2]2+ (aq)
With carbonate ions
With carbonate ions, copper(II) carbonate precipitates out:
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[Cu(H2O)6] 2+ (aq) + CO32- (aq) → CuCO3 (s) + 6H2O (l)
The blue [Cu(H2O)6]2+ solution reacts to form a blue / blue-green precipitate of copper(II) carbonate,
CuCO3 (s)
Hexaaqua iron(III) reactions
With hydroxide ions
The hexaaqua iron(III) undergoes deprotonation reactions with dilute and excess hydroxide ions
This is a three-step process:
[Fe(H2O)6]3+ (aq) + OH– (aq) → [Fe(H2O)5(OH)]2+ (aq) + H2O (l)
[Fe(H2O)5(OH)]2+ (aq) + OH– (aq) → [Fe(H2O)4(OH)2]+ (aq) + H2O (l)
[Fe(H2O)4(OH)2]+ (aq) + OH– (aq) → Fe(H2O)3(OH)3 (s) + H2O (l)
The yellow-orange [Fe(H2O)6]3+ solution reacts to form a red-brown precipitate of hydrated iron(III)
hydroxide, Fe(H2O)3(OH)3 (s)
There is no further reaction with excess hydroxide ions
With ammonia solution
Ammonia behaves in the same way as sodium hydroxide because it is a base
It removes protons from the water ligands
The overall reaction with ammonia is:
[Fe(H2O)6]3+ (aq) + 3NH3 (aq) → Fe(H2O)3(OH)3 (s) + 3NH4+ (aq)
Again, the yellow-orange [Fe(H2O)6]3+ solution reacts to form a red-brown precipitate of hydrated
iron(III) hydroxide, Fe(H2O)3(OH)3 (s)
There is no further reaction with excess ammonia
With carbonate ions
Hexaaqua 3+ ions are acidic in water
This means that they undergo neutralisation reactions with carbonate ions
The reaction between hexaaqua iron(III) ions and water exists in an equilibrium
[Fe(H2O)6]3+ (aq) + 3H2O (l) ⇌ Fe(H2O)3(OH)3 (s) + 3H3O+ (aq)
The hydronium ions, H3O+, react with carbonate ions to produce carbon dioxide
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2H3O+ (aq) + CO32- (aq) → CO2 (g) + 3H2O (l)
This reaction removes hydronium ions from the hexaaqua iron(III) equilibrium
This pushes the equilibrium to the right
Therefore, the hydrated iron(III) hydroxide precipitates out
The overall reaction equation is:
2[Fe(H2O)6]3+ (aq) + 3CO32− (aq) → 2Fe(H2O)3(OH)3 (s) + 3CO2 (g) + 3H2O (l)
Hexaaqua aluminium reactions
With hydroxide ions
The hexaaqua aluminium undergoes deprotonation reactions with dilute and excess hydroxide ions
This is a three-step process:
[Al(H2O)6]3+ (aq) + OH– (aq) → [Al(H2O)5(OH)]2+ (aq) + H2O (l)
[Al(H2O)5(OH)]2+ (aq) + OH– (aq) → [Al(H2O)4(OH)2]+ (aq) + H2O (l)
[Al(H2O)4(OH)2]+ (aq) + OH– (aq) → Al(H2O)3(OH)3 (s) + H2O (l)
The colourless [Al(H2O)6]3+ solution reacts to form a white precipitate of Al(H2O)3(OH)3 (s)
Further deprotonation reactions occur with excess hydroxide ions
Al(H2O)3(OH)3 (s) + OH- (aq) ⇋ [Al(H2O)2(OH)4]- (aq) + H2O (l)
[Al(H2O)2(OH)4]- (aq) + OH- (aq) ⇋ [Al(H2O)(OH)5]2- (aq) + H2O (l)
[Al(H2O)(OH)5]2- (aq) + OH- (aq) ⇋ Al(OH)63- (aq) + H2O (l)
The white Al(H2O)3(OH)3 (s) precipitate reacts to form a colourless solution of aluminium hydroxide,
Al(OH)6 (aq)
With ammonia solution
Ammonia behaves in the same way as sodium hydroxide because it is a base
It removes protons from the water ligands
The overall reaction with ammonia is:
[Al(H2O)6] 3+ (aq) + 3NH3 (aq) → Al(H2O)3(OH)3 (s) + 3NH4+ (aq)
Again, the colourless [Al(H2O)6]3+ solution reacts to form a white precipitate of aluminium hydroxide,
Al(H2O)3(OH)3 (s)
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There is no further reaction with excess ammonia
With carbonate ions
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Hexaaqua 3+ ions are acidic in water
This means that they undergo neutralisation reactions with carbonate ions
The reaction between hexaaqua aluminium ions and water exists in an equilibrium
[Al(H2O)6]3+ (aq) + 3H2O (l) ⇌ Al(H2O)3(OH)3 (s) + 3H3O+ (aq)
The hydronium ions, H3O+, react with carbonate ions to produce carbon dioxide
2H3O+ (aq) + CO32- (aq) → CO2 (g) + 3H2O (l)
This reaction removes hydronium ions from the hexaaqua iron(III) equilibrium
This pushes the equilibrium to the right
Therefore, the aluminium hydroxide precipitates out
The overall reaction equation is:
2[Al(H2O)6]3+ (aq) + 3CO32− (aq) → 2Al(H2O)3(OH)3 (s) + 3CO2 (g) + 3H2O (l)
Summary of reactions
Ion
Colour of
hexaaqua
solution
Reaction with OH-
Reaction with NH3
Reaction with CO32-
Fe2+
Pale green
Forms a dark green
precipitate
Forms a dark green
precipitate
Forms a green
precipitate
Insoluble in excess
Insoluble in excess
Cu2+
Al3+
Blue
Colourless
Forms a blue precipitate Forms a blue precipitate
Insoluble in excess
Redissolves in excess to
form a deep blue
solution
Forms a white
precipitate
Forms a white
precipitate
Redissolved in excess to
form a colourless
Insoluble in excess
Forms a blue
precipitate
Forms a white
precipitate and
bubbles of gas
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solution
Fe3+
Yellow-orange
Forms a red-brown
precipitate
Forms a red-brown
precipitate
Insoluble in excess
Insoluble in excess
Forms a red-brown
precipitate and
bubbles of gas
Examiner Tips and Tricks
Transition metals in the +3 state are acidic and do not form carbonate precipitates, unlike the +2
ions.
Amphoteric Hydroxides
Aluminium hydroxide is classified as an amphoteric hydroxide
Amphoteric means it reacts with both acids and bases
Aluminium hydroxide is insoluble in water but readily dissolves in dilute hydrochloric acid producing the
hexaaquaaluminium ion:
Al(OH)3(H2O)3 (s) + 3HCl (aq) → [Al(H2O)6]3+ (aq) + 3Cl- (aq)
Aluminium hydroxide dissolves in sodium hydroxide to form sodium tetrahydoxoaluminate
Al(OH)3(H2O)3 (s) + NaOH (aq) → Na[Al(OH)4] (aq) + 3H2O (l)
You need a strong base to carry out the reaction, so it is usually done with hot concentrated sodium
hydroxide
Examiner Tips and Tricks
You can also show the reactions with sodium hydroxide as:
Al(OH)3(H2O)3 (s) + OH- (aq) → [Al(OH)4]- (aq) + 3H2O (l)
Al(OH)3(H2O)3 + OH– → [Al(OH)4(H2O)2]– + H2O
Al(OH)3 + NaOH → NaAl(OH)4
Al(OH)3 + OH– → [Al(OH)4]-
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Identifying Transition Metal Ions
Required Practical 9: Transition Metal Ions in Aqueous
Solution
Objective
To carry out simple test–tube reactions to identify transition metal ions in aqueous solution.
Apparatus
Solution Q
Solution R
Solution S
1.0 mol dm-3 sodium hydroxide solution
1.0 mol dm-3 sodium carbonate solution
0.05 mol dm-3 silver nitrate solution
12 test tubes
Test-tube rack
7 dropping pipettes
250 cm3 beaker
Hot water
Distilled / deionised water
Method
Test 1 - part a
Note the initial colour of solution Q
Place 10 drops of solution Q in a test tube
Add sodium hydroxide solution, dropwise with gentle shaking, until in excess
Keep the test tube for part b
Record any observations in an appropriate table
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Repeat this test with solutions R and S
Test 1 - part b
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Half fill a 250 cm3 beaker with freshly boiled water
Stand the three test tubes from part a in the beaker of hot water for about 10 minutes
Record any observations in an appropriate table
Test 2
Place 10 drops of sodium carbonate solution in a test tube
Add 10 drops of solution Q
Shake the mixture gently
Record any observations in an appropriate table
Repeat this procedure with solutions R and S
Test 3
Place about 10 drops of solution Q in a test tube
Add about 10 drops of silver nitrate solution and shake the mixture gently
Repeat this procedure with solutions R and S
Allow the three test tubes to stand for about 10 minutes
Record any observations in an appropriate table
Practical Tip
Always use clean test tubes and pipettes as the tests are very sensitive and you don't want to crosscontaminate the solutions
It's a good idea to label your test tubes to avoid confusing ones that have the same appearance,
especially when they are in a water bath and not arranged in a test tube rack
To observe colour changes more clearly, using a white background such as a piece of paper can help
Results
The results for this required practical are the observations of solutions Q, R and S in tests 1, 2 and 3
A suitable results table could be:
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Test
Q
R
S
Your notes
1a - Initial colour
1a - Add NaOH (aq)
1a - Stand in water bath
2 - Add Na2CO3 (aq)
3 - Add AgNO3 (aq)
Evaluation
The results from all three tests are used to identify the ions present in solutions Q, R and S
Worked Example
The following tests were completed on solutions Q, R and S. The results are shown in the table.
Test
Q
R
S
Initial colour
yellow solution
light blue solution
pale green
solution
Add NaOH (aq)
orange/brown precipitate
blue precipitate
grey/green
precipitate
Add excess
NaOH (aq)
no visible change
no visible change
no visible change
Add Na2CO3 (aq)
orange/brown precipitate and
effervescence
blue green
precipitate
grey/green
precipitate
Add AgNO3 (aq)
no visible change
white precipitate
no visible change
Identify the anions and cations present, where possible, in solutions Q, R and S.
Answers:
Solution Q - 0.2 mol dm-3 iron(III) nitrate solution
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Solution R - 0.2 mol dm-3 copper(II) chloride solution
Solution S - 0.5 mol dm-3 ammonium iron(II) sulfate solution
Solution Q
Test 1 results
Yellow solution suggests a Fe3+ cation
Orange brown precipitate with NaOH (aq) suggests a Fe3+ cation
No further change with excess NaOH (aq) suggests a Fe3+ cation
Test 2 results
Orange/brown precipitate and effervescence suggests a Fe3+ cation
Test 3 results
No visible change suggests the anion is not a halide ion
Anion = cannot be determined
Cation = iron(III) / Fe3+
Solution R
Test 1 results
Light blue solution suggests a Cu2+ cation
Blue precipitate with NaOH (aq) = suggests a Cu2+ cation
No further change with excess NaOH (aq) suggests a Cu2+ cation
Test 2 results
Blue/green precipitate and effervescence suggests a Cu2+ cation
Test 3 results
White precipitate suggests a Cl- anion
Anion = chloride / ClCation = copper(II) / Cu2+
Solution S
Test 1 results
Pale green solution suggests a Fe2+ cation
Grey / green precipitate with NaOH (aq) suggests a Fe2+ cation
No further change with excess NaOH (aq) suggests a Fe2+ cation
Test 2 results
Grey/green precipitate suggests a Fe2+ cation
Test 3 results
No visible change suggests the anion is not a halide ion
Anion = cannot be determined
Cation = iron(II) / Fe2+
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