Ions in Aqueous Solution (Edexcel International A Level Chemistry)

• When transition metal ions in an aqueous solution react with aqueous sodium hydroxide and aqueous ammonia they form precipitates
• However, some of these precipitates will dissolve in an excess of sodium hydroxide or ammonia to form complex ions in solution

The Reactions of Aqueous Transition Metal Ions with Aqueous Sodium Hydroxide

Transition Metal Ion	Metal-aqua ion	With NaOH (aq)	With excess NaOH (aq)
Cr3+	Green solution [Cr(H2O)6]3+ (aq)	Green precipitate Cr(OH)3(H2O)3 (s)	Green solution [Cr(OH)6]3– (aq)
Fe2+	Green solution [Fe(H2O)6]2+ (aq)	Green precipitate Fe(OH)2(H2O)4 (s)	No further change
Fe3+	Yellow-brown solution [Fe(H2O)6]3+ (aq)	Brown precipitate Fe(OH)3(H2O)3 (s)	No further change
Co2+	Pink solution [Co(H2O)6]2+ (aq)	Blue precipitate Co(OH)2(H2O)4 (s)	No further change
Mn2+	Pale pink solution [Mn(H2O)6]2+ (aq)	Pale brown precipitate Mn(OH)2(H2O)4 (s)	No further change
Ni2+	Green solution [Ni(H2O)6]2+ (aq)	Green precipitate Ni(OH)2(H2O)4 (s)	No further change
Cu2+	Blue solution [Cu(H2O)6]2+ (aq)	Blue precipitate Cu(OH)2(H2O)4 (s)	No further change
Zn2+	Colourless solution [Zn(H2O)6]2+ (aq)	White precipitate Zn(OH)2(H2O)4 (s)	Colourless solution [Zn(OH)4(H2O)2]2– (aq)

 

• Examples of ionic equations for the reactions in the table above
  • [Fe(H₂O)₆]²⁺ (aq) + 2OH^- (aq) → [Fe(H₂O)₄(OH)₂] (s) + 2H₂O (l) 
  • [Cu(H₂O)₆]²⁺ (aq) + 2OH^- (aq) → [Cu(H₂O)₄(OH)₂] (s) + 2H₂O (l) 
  • [Fe(H₂O)₆]³⁺ (aq) + 3OH^- (aq) → [Fe(H₂O)₃(OH)₃] (s) + 3H₂O (l) 

The Reactions of Aqueous Transition Metal Ions with Ammonia

Transition Metal Ion	Metal-aqua ion	With NH3 (aq)	With excess NH3 (aq)
Cr3+	Green solution [Cr(H2O)6]3+ (aq)	Green precipitate Cr(OH)3(H2O)3 (s)	Purple solution [Cr(NH3)6]3+ (aq)
Fe2+	Green solution [Fe(H2O)6]2+ (aq)	Green precipitate Fe(OH)2(H2O)4 (s)	No further change
Fe3+	Yellow-brown solution [Fe(H2O)6]3+ (aq)	Brown precipitate Fe(OH)3(H2O)3 (s)	No further change
Co2+	Pink solution [Co(H2O)6]2+ (aq)	Blue precipitate Co(OH)2(H2O)4 (s)	Yellow solution [Co(NH3)6]2+ (aq)
Mn2+	Pale pink solution [Mn(H2O)6]2+ (aq)	Pale brown precipitate Mn(OH)2(H2O)4 (s)	No further change
Ni2+	Green solution [Ni(H2O)6]2+ (aq)	Green precipitate Ni(OH)2(H2O)4 (s)	Dark blue solution [Ni(NH3)6]2+ (aq)
Cu2+	Blue solution [Cu(H2O)6]2+ (aq)	Blue precipitate Cu(OH)2(H2O)4 (s)	Dark blue solution [Cu(NH3)4(H2O)2]2+ (aq)
Zn2+	Colourless solution [Zn(H2O)6]2+ (aq)	White precipitate Zn(OH)2(H2O)4 (s)	Colourless solution [Zn(NH3)4]2+ (aq)

 

• Solutions of metal aqua ions react as​ acids ​with aqueous ammonia
• This means that the ammonia solution will initially act as a​ base​ to remove one H⁺ ion per ammonia molecule used
  • Examples of ionic equations for this type of reaction from the table above:
    • [Fe(H₂O)₆]²⁺ (aq) + 2NH₃ (aq) → Fe(H₂O)₄(OH)₂ (s) + 2NH₄⁺ (aq)
    • [Cu(H₂O)₆]²⁺ (aq) + 2NH₃ (aq) → Cu(H₂O)₄(OH)₂ (s) +2NH₄⁺ (aq)
    • [Fe(H₂O)₆]³⁺ (aq) + 3NH₃ (aq) → Fe(H₂O)₃(OH)₃ (s) +3NH₄⁺ (aq)
• Some metal aqua ions will react further and undergo ligand substitution with excess ammonia
  • Two specific examples to be aware of are Co²⁺ and Cu²⁺ with excess ammonia:
  • Starting with the hexa aqua ions:
    • [Cu(H₂O)₆]²⁺ (aq) + 4NH₃ (aq) → [Cu(H₂O)₂(NH₃)₄]²⁺ (aq) + 4H₂O (l) 
    • [Co(H₂O)₆]²⁺ (aq) + 6NH₃ (aq) → [Co(NH₃)₆]²⁺ (aq) + 6H₂O (l) 
  • Starting with the precipitates: 
    • [Cu(H₂O)₄(OH)₂] (s) + 4NH₃ (aq) → [Cu(H₂O)₂(NH₃)₄]²⁺ (aq) + 2H₂O (l) + 2OH^- (aq)
    • [Co(H₂O)₄(OH)₂] (s) + 6NH₃ (aq) → [Co(NH₃)₆]²⁺ (aq) + 4H₂O (l) + 2OH^- (aq)​

Reaction with limited OH^- and limited NH₃

• The bases OH^- and ammonia when in limited amounts form the same hydroxide precipitates.
• They form in deprotonation acid base reactions
• For example, consider the reaction that occurs when aqeuous sodium hydroxide is added to copper(II) sulfate solution

[Cu(H₂O)₆]²⁺ (aq) + 2OH^- (aq) → [Cu(H₂O)₄(OH)₂] (s) + 2H₂O (l) 

• This seems like a ligand substitution reaction - two hydroxide ions replacing two water molecules
• However this is actually a deprotonation reaction - two hydroxide ions removing hydrogen ions from two of the water ligands converting them into water molecules
  • The two ligands that have lost hydrogen ions are now hydroxide ligand

Reaction with excess OH^-

• From above, we have seen how hydrated transition metal ions can be deprotonated by adding a base such as aqueous sodium hydroxide to form a metal hydroxide precipitate
• For example 

[Cr(H₂O)₆]³⁺ (aq) + 3OH^- (aq) → [Cr(H₂O)₃(OH)₃] (s) +3H₂O (l) 

• When an excess of sodium hydroxide is added further deprotonation takes place

[Cr(H₂O)₃(OH)₃] (s) + 3OH^- (aq) → [Cr(OH)₆]^3- (aq) + 3H₂O (l) 

• In this reaction, chromium(III) hydroxide acts as an acid, as it is reacting with a base
• Chromium(III) hydroxide can also act as a base because it can react with acids as follows

[Cr(H₂O)₃(OH)₃] (s) + 3H⁺ (aq) → [Cr(H₂O)₆]³⁺ (aq) 

• A metal hydroxide that can act as both an acid and a base is called an amphoteric hydroxide
  • This is an example of amphoteric behaviour

Reaction with excess NH₃

• With excess NH₃ ligand substitution reactions occur with Cu, Co and Cr and their precipitates dissolve
• The ligands NH₃ and H₂O are similar in size and are uncharged
• Ligand exchange occurs without a change of co-ordination number for Co and Cr
• For example when excess aqueous ammonia is added to a copper(II) hydroxide precipitate it dissolves forming a deep blue solution

[Cu(H₂O)₄(OH)₂] (s) + 4NH₃ (aq) → [Cu(NH₃)₄(H₂O)₂]²⁺ (aq) + 2H₂O (l) + 2OH^- (aq)

• This is a ligand substitution - four ammonia molecules replace two water molecules and two hydroxide ions
• In these reactions NH₃ is acting as a Lewis base donating an electron pair