Metal-Aqua Ions (AQA A Level Chemistry): Revision Note

Metal Aqua-Ions

• Transition metal salts that are readily soluble in water are forming metal-aqua complex ions when they dissolve in water

• When we show copper sulfate dissolving in water it is usually written as:

CuSO₄ (s) + aq  → Cu²⁺ (aq)   + SO₄^2- (aq)

• However, what we really mean by this is that copper(II) ions are forming the hexaaqucopper(II) ion

CuSO₄ (s) + aq  → [Cu(H₂O)₆] ²⁺ (aq)   + SO₄^2- (aq)

• The water molecules are ligands attached to the central transition metal cation by dative covalent bonding from the lone pairs on the oxygen molecules

• Iron(II) salts also form the hexaaqua complex ions:

Fe(NO₃)₂ (s) + aq  → [Fe(H₂O)₆] ²⁺ (aq)   + 2NO₃^- (aq)

• Two common +3 aqua ions you should know are iron(III) and aluminium:

Fe(NO₃)₃ (s) + aq  → [Fe(H₂O)₆] ³⁺ (aq)   + 3NO₃^- (aq)

Al₂(SO₄)₃ (s) + aq  → 2[Al(H₂O)₆] ³⁺ (aq)   + 3SO₄^2- (aq)

Hexaaqua complex ions of +2 and +3 metal ions

Lewis Acids and Bases

• In the section of acids and bases we saw the Brønsted-Lowry definition of acids and bases

An acid is a substance which can behave as a proton donor

A base is a substance which can behave as a proton acceptor

• Gilbert Lewis is better known in the topic of bonding theory as the chemist whose name is given to electron dot structures of atoms, ions and molecules

• Lewis applied these structures to Brønsted-Lowry theory and realised that acid-base reactions can be interpreted in terms of electron pairs movements rather than in terms of proton transfer

• For example when ammonia reacts with a proton an electron pair moves from ammonia to the proton:

H⁺     +     NH₃   →   NH₄⁺

The Lewis diagram shows that the base donates an electron pair to forming a dative covalent bond with the proton

• This lead to a new definition of acids and bases

A Lewis acid is a species which can bond by accepting a lone pair of electrons

A Lewis base is a species which can bond by donating a lone pair of electrons

• The bond formed is an example of a co-ordinate or dative covalent bond

• Following these definitions you should be able to see that in a metal-aqua complex ion:

  • The metal is a Lewis acid

  • The water is a Lewis base

• Although this definition of acids and bases is not required to be learned for the exam, you can see that it is a very useful concept for chemists when explaining movements of electrons pairs in a variety of situations

Acidity in Metal-Aqua Ions

• You might imagine that salts of transition metal ions would be neutral in water

• However, ions of +3 aqua complexes are noticeably acidic compared to the +2 ions

• For example the pK_a of [Fe(H₂O)₆] ³⁺ (aq) is 2.2 which is significantly more acidic than ethanoic acid which has a pK_a of 4.8 (remember the lower the pK_a the more acidic the species is)

• The reason for this is that +3 ions are smaller and therefore have a higher charge density than +2 ions

• The higher charge density pulls the water molecules more strongly, which weakens the O-H bond and results in more dissociation, producing a more acidic solution

• We say that the metal ion polarises the water molecules

Metal(III) ions have a high charge density and polarise water molecules in the hexaaqua complexes

• This results in hydrogen ions (protons) splitting from the complexes creating acidic solutions

 [Fe(H₂O)₆] ³⁺ (aq)   →    [Fe(H₂O)₅(OH)] ²⁺ (aq)  + H⁺ (aq)

• Notice that the loss of the proton results in the new pentaaqua complex having a +2 charge

  • Another way to think of this is the left over hydroxide ion from splitting a water molecule is negatively charged and cancels one of the charges on the complex ion

• The reaction process can also be shown with the hydroxonium ion as the product:

 [Fe(H₂O)₆] ³⁺ (aq) + H₂O (l)   →    [Fe(H₂O)₅(OH)] ²⁺ (aq)  + H₃O⁺ (aq)

• These reactions are sometime called deprotonation reactions and occur in several steps

[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)

• However, the third deprotonation does not usually occur without the presence of a base

• The base facilitates the removal of the third proton and results in the precipitation of the insoluble hydrated iron(III)hydroxide which appears as a foxy red precipitate:

[Fe(H₂O)₄(OH)₂] ⁺ (aq)  + OH^- →    Fe(H₂O)₃(OH)₃ (s)  + H₂O (l)