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2.6Inorganic Chemistry

Reactions of Ions in Aqueous Solution

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Complete PDF with figures and original page layoutPage 1 Oxford AQA International Your notes A Level (IAL) Chemistry Reactions of Ions in Aqueous Solution Contents Reactions of Ions in Aqueous Solution Reactions of Metal-Aqua Ions Identifying Transition Me…

Section 1 of 14Metal-aqua ions

transition metal complex colours
transition metal complex colours
Formation
  • Formation soluble transition metal salts dissolve to give metal-aqua complex ions, with water ligands bonded by oxygen lone pairs
Copper(II) sulfate
  • CuSO4(s) + aq → [Cu(H2O)6]2+(aq) + SO42-(aq).
  • The hexaaqua copper(II) ion makes the solution blue.
IonColour of solution
[Cu(H2O)6]2+Blue
[Fe(H2O)6]2+Green
[Fe(H2O)6]3+Yellow-orange
[Al(H2O)6]3+Colourless
  • Structure
    • all have six water ligands and a 2+ or 3+ charge

Section 2 of 14Acidity of metal-aqua ions

  • Observation
    • solutions of metal salts are not neutral, and the 3+ ions are more acidic than 2+ ions
0.1 mol dm-3 solutionpH
Iron(III) chloride2.0
Aluminium chloride3.0
Copper(II) nitrate4.0
  • Explanation
    • 3+ ions are smaller than 2+ ions, so they have a higher charge density.
    • They attract the lone pair of a water ligand more strongly, polarising it.
    • The O–H bonds weaken and dissociate, releasing H+ and leaving a hydroxide ligand.

Section 3 of 14Deprotonation reactions

1Equation
  • [Fe(H2O)6]3+(aq) → [Fe(H2O)5(OH)]2+(aq) + H+(aq)
2With water
  • [Fe(H2O)6]3+ + H2O → [Fe(H2O)5(OH)]2+ + H3O+
3Charge
  • Charge the remaining OH- cancels one charge of the 3+ ion
4Steps
  • Steps a second deprotonation gives [Fe(H2O)4(OH)2]+, and the third needs a base, giving red-brown Fe(H2O)3(OH)3(s)
5Name
  • Name splitting water molecules can be called hydrolysis

Section 4 of 14Hexaaqua iron(II) with bases

Hexaaqua iron(II) with bases
With hydroxide
  • Two deprotonation steps give Fe(H2O)4(OH)2(s), a dark green precipitate.
  • The green precipitate slowly turns orange-brown as it is oxidised to hydrated iron(III) hydroxide.
  • Excess hydroxide causes no further reaction.
With ammonia
  • [Fe(H2O)6]2+ + 2NH3 → Fe(H2O)4(OH)2(s) + 2NH4+.
  • Ammonia acts as a base and removes protons, giving the same dark green precipitate that oxidises to orange-brown.
  • Excess ammonia causes no further reaction.
With carbonate
  • [Fe(H2O)6]2+ + CO32- → FeCO3(s) + 6H2O.
  • A green precipitate of iron(II) carbonate forms.

Section 5 of 14Hexaaqua copper(II) with bases

copper hydroxide ammonia complex changes
copper hydroxide ammonia complex changes
Hexaaqua copper(II) with bases
With hydroxide
  • Two deprotonation steps give a blue precipitate of Cu(H2O)4(OH)2.
  • Excess hydroxide causes no further reaction.
With ammonia
  • Dropwise ammonia first acts as a base: [Cu(H2O)6]2+ + 2NH3 → Cu(H2O)4(OH)2(s) + 2NH4+.
  • Excess ammonia partly substitutes water, as ammonia is a stronger ligand: Cu(H2O)4(OH)2(s) + 4NH3 → [Cu(NH3)4(H2O)2]2+ + 2OH- + 2H2O.
  • The precipitate dissolves to a deep blue solution.
With carbonate
  • [Cu(H2O)6]2+ + CO32- → CuCO3(s) + 6H2O.
  • A green precipitate of copper(II) carbonate forms.

Section 6 of 14Hexaaqua iron(III) with bases

With hydroxide
  • Three deprotonation steps give Fe(H2O)3(OH)3(s), a red-brown precipitate.
  • Excess hydroxide causes no further reaction.
With ammonia
  • [Fe(H2O)6]3+ + 3NH3 → Fe(H2O)3(OH)3(s) + 3NH4+.
  • The same red-brown precipitate forms, and excess ammonia causes no further change.
With carbonate
  • With carbonate neutralisation of acidic hexaaqua 3+ ions
  • The acid equilibrium [Fe(H2O)6]3+ + 3H2O ⇌ Fe(H2O)3(OH)3(s) + 3H3O+ exists.
  • H3O+ reacts with carbonate: 2H3O+ + CO32- → CO2 + 3H2O.
  • Removing H3O+ shifts the equilibrium right, so the hydroxide precipitates.
  • Overall2[Fe(H2O)6]3+ + 3CO32- → 2Fe(H2O)3(OH)3(s) + 3CO2 + 3H2O.

Section 7 of 14Hexaaqua aluminium with bases

Hexaaqua aluminium with bases
With hydroxide
  • Three deprotonation steps give a white precipitate of Al(H2O)3(OH)3.
  • Excess hydroxide redissolves it by further deprotonation, via [Al(H2O)2(OH)4]- and [Al(H2O)(OH)5]2- to [Al(OH)6]3-, a colourless solution.
With ammonia
  • [Al(H2O)6]3+ + 3NH3 → Al(H2O)3(OH)3(s) + 3NH4+.
  • A white precipitate forms and excess ammonia does not dissolve it.
With carbonate
  • It follows the same acid-equilibrium route as iron(III), forming white Al(H2O)3(OH)3(s) and carbon dioxide.

Section 8 of 14Amphoteric hydroxides

Amphoteric = reacting with both acids and bases
Aluminium hydroxide
  • Aluminium hydroxide it is insoluble in water but dissolves in both dilute acid and excess strong base
With acid
  • Al(OH)3(H2O)3(s) + 3HCl(aq) → [Al(H2O)6]3+(aq) + 3Cl-(aq)
With base
  • Al(OH)3(H2O)3(s) + NaOH(aq) → Na[Al(OH)4](aq) + 3H2O(l), giving sodium tetrahydroxoaluminate
Conditions
  • Conditions a strong base is needed, usually hot concentrated NaOH
Equivalent forms
  • Al(OH)3(H2O)3 + OH- → [Al(OH)4]- + 3H2O, or Al(OH)3 + OH- → [Al(OH)4]-

Section 9 of 14Required Practical 9: aim and apparatus

Aim
  • Aim to identify transition metal ions in aqueous solution by test-tube reactions
Apparatus
  • Apparatus solutions Q, R and S, 1.0 mol dm-3 sodium hydroxide, 1.0 mol dm-3 sodium carbonate, 0.05 mol dm-3 silver nitrate, 12 test tubes with rack, 7 dropping pipettes, 250 cm3 beaker, hot water and deionised water

Section 10 of 14Required Practical 9: methods

1
Test 1 part a
  • Note the initial colour of solution Q.
  • Put 10 drops of Q in a test tube and add NaOH dropwise with gentle shaking until in excess.
  • Keep the tube for part b and record observations.
  • Repeat with R and S.
2
Test 1 part b
  • Half fill a 250 cm3 beaker with freshly boiled water.
  • Stand the three tubes in it for about 10 minutes and record observations.
3
Test 2
  • Put 10 drops of sodium carbonate in a test tube and add 10 drops of Q.
  • Shake gently and record observations.
  • Repeat with R and S.
4
Test 3
  • Put about 10 drops of Q in a tube and add about 10 drops of silver nitrate.
  • Shake gently, repeat with R and S, and let the tubes stand for about 10 minutes.
  • Record observations.

Section 11 of 14Required Practical 9: practical tips and results

Tips
  • Use clean tubes and pipettes, because the tests are sensitive to contamination.
  • Label the tubes, especially in the water bath.
  • Use a white background to see colour changes clearly.
Results
  • Results record the observations for Q, R and S in tests 1a , 2 (sodium carbonate) and 3 (silver nitrate)
    • (initial colour, NaOH, water bath)
Evaluation
  • Evaluation use all three tests together to identify the ions

Section 12 of 14Required Practical 9: interpreting results

TestQRS
Initial colourYellowLight bluePale green
NaOHOrange-brown precipitateBlue precipitateGrey-green precipitate
Excess NaOHNo changeNo changeNo change
Na2CO3Orange-brown precipitate and effervescenceBlue-green precipitateGrey-green precipitate
AgNO3No changeWhite precipitateNo change
Q
Q yellow solution, orange-brown precipitate unchanged in excess, and effervescence with carbonate give Fe3+, with no halide,
  • so the anion cannot be determined (iron(III) nitrate)
R
R light blue solution, blue precipitate unchanged in excess and blue-green carbonate precipitate give Cu2+,
  • and the white precipitate with silver nitrate shows chloride (copper(II) chloride)
S
S pale green solution, grey-green precipitate unchanged in excess and a grey-green carbonate precipitate give Fe2+, with no halide,
  • so the anion cannot be determined

Section 13 of 14Key terms

TermMeaning
Metal-aqua ionA metal ion surrounded by water ligands.
DeprotonationThe loss of H+ from a ligand.
Charge densityThe charge of an ion relative to its size.
Amphoteric hydroxideA hydroxide that dissolves in both acids and excess alkali.
HydrolysisSplitting of a water molecule, here releasing H+.
TermMeaning
AmphotericAble to react with both acids and bases.
TetrahydroxoaluminateThe [Al(OH)4]- ion.
PrecipitateAn insoluble solid formed in a solution.
EffervescenceFizzing caused by gas release.
ExcessMore reagent than is needed for complete reaction.

Section 14 of 14Exam technique

1
Saying 2+ and 3+ ions are equally acidic — the 3+ ions are more acidic because of higher charge density.
2
Saying iron(III) carbonate precipitates — the hydroxide precipitates and CO2 is released.
3
Saying aluminium hydroxide stays in excess ammonia — it does not dissolve in excess ammonia, but does in excess hydroxide.
4
Forgetting that green iron(II) hydroxide oxidises in air — it turns orange-brown.
1
Saying iron(III) forms a carbonate precipitate — it forms the hydroxide with CO2.
2
Saying aluminium hydroxide is insoluble in excess NaOH — it redissolves.
3
Saying Cu(OH)2 dissolves in excess NaOH — it dissolves only in excess ammonia.
4
Using dirty glassware — contamination gives wrong observations.
Chemistry 0620 & Edexcel OL Product Prediction Rules

Aqueous Electrolysis Rules (Inert Electrodes)

  • At Cathode (-): Least reactive cation discharged. If metal is more reactive than hydrogen (Na+, K+, Mg2+), H2 gas is produced: 2H+ + 2e- → H2. If Cu2+ or Ag+, the metal coats the cathode.
  • At Anode (+): If concentrated halide present (Cl-, Br-, I-), halogen is formed (e.g. 2Cl- → Cl2 + 2e-). If sulfate (SO42-) or nitrate (NO3-), O2 gas is discharged: 4OH- → O2 + 2H2O + 4e-.