2.6 Reactions of ions in aqueous solution

• 3+ ions are much more acidic as they are smaller and more highly charged, have a higher charge density, and attract the electrons from the oxygen of the ligands much more strongly, weakening the O-H bonds and causing the complex to release a H+ ion into the solution, making it acidic.
• The equation for the hydrolysis of a 3+ ion to release a proton

  [Fe(H2O)6]3+ (aq) ⇌ [Fe(H2O)5(OH)]2+ (aq) + H+ (aq) REVERSIBLE REACTION
• The complex ion acts as a Bronsted-Lowry acid by donating a proton.
• A Lewis acid is an electron pair acceptor.
• A Lewis base is an electron pair donor.
• During the formation of complex ions, the metal ion acts as the Lewis acid and the ligands as Lewis bases.
• Fe2+ reactions:

  A) green

  B) green precipitate

  C) brown

  D) air

  E) green precipitate

  F) brown

  G) air

  H) green precipitate

  
• The green ppt go brown on standing in air in the reaction of Fe2+ with NH3 because oxygen in the air oxidises Fe2+ to Fe3+, and [Fe(OH)3 (H2O)3 ] is brown, so colour changes
• reactions of Cu2+:

  A) blue

  B) blue precipitate

  C) blue precipitate

  D) deep blue solution

  E) blue-green precipitate

  F) pale green / yellow solution

  
• reactions of Fe3:

  A) purple solution

  B) yellow-brown

  C) brown precipitate

  D) orange-brown

  E) brown precipitate

  F) orange-brown

  G) brown precipitate

  H) orange-brown

  I) effervescence

  
• reactions of Al3+:

  A) colourless solution

  B) white precipitate '

  C) colourless solution

  D) white precipitate

  E) white precipitate

  F) effervescence

  
• reactions of Co2+:

  A) pink

  B) blue precipitate

  C) pale brown solution

  D) purple precipitate

  E) blue solution

  
• The following equilibria happen in aqueous solutions of metal ions.

  [M(H2O)6]2+ + H2O --> [M(H2O)5 (OH)]+ + H3O+

  [M(H2O)6 ]3+ + H2O --> [M(H2O)5 (OH)]2+ + H3O+
• The bases OH- and ammonia when in limited amounts form the same hydroxide precipitates. They form in deprotonation acid base reactions.

  Here the NH3 and OH- ions are acting as Bronsted-Lowry bases accepting a proton
• With excess NaOH the aluminium hydroxide dissolves. Al becomes [Al(OH)4 ] - (aq) colourless solution.

  Al(H2O)3(OH)3(s) + OH- (aq) --> [Al(OH)4 - (aq) + 3H2O (l)

  This hydroxide is classed as amphoteric because it reacts and dissolves in both acids and bases.

  Al(H2O)3(OH)3 (s) + 3H+ (aq) --> [Al(H2O)6 ]3+ (aq)
• With excess NH3 a ligand substitution reaction occurs with Cu and its precipitate dissolves to form a deep blue solution. This substitution is incomplete with Cu

  
  

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

  
  

  In this reactions NH3 is acting as a Lewis base donating an electron pair.
• The 2+ ions with carbonate solution results in MCO3 ppt being formed (Cu blue/green, Fe(II) green )

  Cu2+ (aq) + CO3 2- (aq) --> CuCO3 (s)

  Fe2+ (aq) + CO3 2- (aq) -> FeCO3 (s)

  These are precipitation reactions
• The 3+ ions with carbonate solution form a M(OH)3(H2O)3 ppt and CO2 gas is evolved as 3+ has greater polarising power than 2+ due to its higher charge density

  
  

  2[Fe(H2O)6]3+ (aq) + 3CO3 2- (aq) --> 2Fe(OH)3(H2O)3(s) +3CO2 + 3H2O(l)

  2[Al (H2O)6]3+ (aq) + 3CO3 2- (aq) --> 2Al(OH)3 (H2O)3(s) +3CO2 + 3H2O(l)
• Acting as a Bronsted-Lowry acid:

  [Al(H2O)6]3+ + H2O -> [Al(H2O)5(OH)]2+ + H3O+