2.6 Reactions of ions in aqueous solution

Fe(H2O)62+ Co(H2O)62+ Cu(H2O)62+ Al(H2O)63+ Cr(H2O)63+ Fe(H2O)63+
Colour of
aqua ion
Green Pink Blue Colourless Ruby/Violet Yellow/brown
Acidity [M(H2O)6]2+ + H2O \rightleftharpoons [M(H2O)5(OH)]+ + H3O+ [M(H2O)6]3+ + H2O \rightleftharpoons [M(H2O)5(OH)]2+ + H3O+
NaOH **[Fe(H2O)4(OH)2]
green(s)
[Co(H2O)4(OH)2]
blue(s)
[Cu(H2O)4(OH)2]
blue(s)
[Al(H2O)3(OH)3]
white(s)
[Cr(H2O)3(OH)3]
green(s)
[Fe(H2O)3(OH)3]
brown(s)
Excess NaOH No change
**[Fe(H2O)4(OH)2]
green(s)
No change
[Co(H2O)4(OH)2]
blue(s)
No change
[Cu(H2O)4(OH)2]
blue(s)
[Al(OH)4]
colourless solution
(i.e. amphoteric
metal hydroxide)
[Cr(OH)6]3−
green solution
(i.e. amphoteric
metal hydroxide)
No change
[Fe(H2O)3(OH)3]
brown(s)
NH3 **[Fe(H2O)4(OH)2]]
green(s)
[Co(H2O)4(OH)2]
blue(s)
[Cu(H2O)4(OH)2]
blue(s)
[Al(H2O)3(OH)3]
white(s)
[Cr(H2O)3(OH)3]
green(s)
[Fe(H2O)3(OH)3]
brown(s)
Excess NH3 No change
**[Fe(H2O)4(OH)2]
green(s)
[Co(NH3)6]2+
yellow solution
oxidation by Oor H2O2 to
[Co(NH3)6]3+
dark brown
[Cu(NH3)4(H2O)2]2+
deep blue solution
No change
[Al(H2O)3(OH)3]
white(s)
[Cr(NH3)6]3+
purple solution
No change
[Fe(H2O)3(OH)3]
brown(s)
HCl No reaction [CoCl4]2−
blue solution
[CuCl4]2−
yellow-green solution
No reaction No reaction [FeCl4]
yellow solution
Na2CO3 FeCO3
green(s)
CoCO3
pink/purple(s)
CuCO3
blue-green(s)
[Al(H2O)3(OH)3]
white(s) + CO2
[Cr(H2O)3(OH)3]
green(s) + CO2
[Fe(H2O)3(OH)3]
brown(s) + CO2
Other reactions **oxidation by O(air) to
[Fe(H2O)3(OH)3]
brown(s)
oxidation by O2 (air) or alkaline H2O2 to
[Co(OH)6]4−  to [Co(OH)6]3− blue → brown precipitate
H2Oin alkaline solution
oxidises
[Cr(OH)6]3−  to CrO42−
green → yellow solution

Definitions of terms

Brønsted-Lowry: B-L acid a proton donor; B-L base is a proton acceptor.

Lewis: Lewis acid is an electron pair acceptor; Lewis base is an electron pair donor.

All Brønsted-Lowry acids are Lewis acids.

Ligand: An atom, ion or molecule that donates a lone pair of electrons (Lewis base action) to a central metal ion to form a co-ordinate bond.

Bidentate ligand: A bidentate ligand donates two electron pairs to a transition metal ion, each pair arising from a different atom located on the same ligand molecule.


The acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ 

A hydrolysis reaction occurs between metal aqua ions and free water molecules present outside the complex and in the solution. In these reactions, the metal aqua ion acts as a Bronsted-Lowry acid and donates a proton from one of its water ligands to a free water molecule to form H3O+ and a hydrated metal hydroxide complex.

M3+, compared to M2+ ions have higher charge density (higher nuclear charge relative to smaller ionic size).
The higher charge density means greater polarisation of the O-H bond of the co-ordinated water ligand molecules that surround M3+, compared to M2+ aqua ion. This weakens the water ligand’s O-H bond and increases its tendency to donate protons to approaching free water molecules, thus making more acidifying hydronium H3O+ ions.

Hence, the hydrolysis equilibrium reaction is further to the right for M3+, compared to M2+ ions, and more acidifying H3O+ ions are produced by M3+ than M2+ hydrolysis.

M3+ hydrolysis equilibrium is further to the right:
[M(H2O)6]3+ + H2\rightleftharpoons  [M(H2O)5(OH)]2++ H3O+
than M2+ hydrolysis equilibrium:                                 
[M(H2O)6]2+ + H2\rightleftharpoons [M(H2O)5(OH)]+ + H3O
+

Importantly, these are NOT ligand substitution reactions: the original coordinate bond between the metal ion and oxygen atom of the water ligand remains intact, only that the water ligand molecule is now deprotonated to form a hydroxide ion.


Reactions with dilute NaOH or dilute NH3

Addition of dilute OH (dilute NaOH or dilute ammonia solution), acts as a base, and shifts the position of the hydrolysis equilibrium to the right through a series of successive water ligand deprotonation reactions until an insoluble uncharged hydrated metal hydroxide complex is formed.

In the case of M3+
[M(H2O)6]3++ 3OH ⇌ [M(H2O)3(OH)3] + 3H2O
In the case of M2+
[M(H2O)6]2++ 2OH⇌ [M(H2O)4(OH)2] + 2H2O

 

 

Examples

[Fe(H2O)6]3++ 3OH ⇌ [Fe(H2O)3(OH)3] + 3H2O    (brown solution → brown ppt.)
[Cu(H2O)6]2++ 2OH⇌ [Cu(H2O)4(OH)2] + 2H2O      (blue solution → blue ppt.)

[Al(H2O)6]3++  3NH ⇌ [Al(H2O)3(OH)3] +  3NH4+   (colourless solution → white ppt.)
[Fe(H2O)6]2++ 2NH⇌ [Fe(H2O)4(OH)2+ 2NH4+   (green solution → green ppt.)

 

Reactions with excess NaOH and excess NH3

Apart from Al3+ and Cr3+ (which form amphoteric metal hydroxides), further addition of NaOH to hydrated metal hydroxides of M2+ and M3+ ions results in no further reaction.

However, further addition of NH3 may induce complete or partial  ligand substitution of the hydrated metal hydroxide complex by NH3.

[Cr(H2O)3(OH)3] + 6NH3 ⇌ [Cr(NH3)6]3+ + 3H2O + 3OH     (green ppt → purple solution)
[Co(H2O)4(OH)2] + 6NH3 ⇌ [Co(NH3)6]2+ + 4H2O + 2OH     (blue ppt → yellow/straw solution)
[Cu(H2O)4(OH)2] + 4NH3 ⇌ [Cu(NH3)4(H2O)2]2+ + 2H2O + 2OH     (blue ppt → deep blue solution)


Al3+ and Cr3+ form amphoteric hydrated metal hydroxides

A metal hydroxide capable of dissolving in both acids and bases is termed an amphoteric metal hydroxide.

Acting as a base, a hydrated metal hydroxide will react with an acid to form its original metal aqua ion: in effect, addition of acid shifts the hydrolysis reaction equilibrium to the left.

However, acting as acids[Al(H2O)3(OH)3] and [Cr(H2O)3(OH)3] are specific examples of hydrated metal hydroxides that are capable of reacting with excess strong bases (e.g.NaOH).

Acting as base (reaction with acid)
[Al(H2O)3(OH)3] + 3H3O⇌ [Al(H2O)6]3+ + 3H2O
[Cr(H2O)3(OH)3] + 3H3O⇌ [Cr(H2O)6]3+ + 3H2O
Acting as acid (reaction with strong base)
[Al(H2O)3(OH)3] +OH ⇌ [Al(OH)4] + 3H2O
[Cr(H2O)3(OH)3] +3OH ⇌ [Cr(OH)6]3− + 3H2O

 

 

 


Reaction with CO32− 

The carbonate ion, CO32−, is able to deprotonate [M(H2O)6]3+ to form a hydrated metal hydroxide complex and HCO3 because of the greater acidity and ligand water molecule polarisation associated with M3+ aqua ions.
In contrast, the less acidic [M(H2O)6]2+complex reacts with CO32− to form a metal carbonate.
i.e. MCO3 is formed but M2(CO3)3 is not formed.

In the case of M3+
2[M(H2O)6]3++3CO32− ⇌ 2[M(H2O)3(OH)3] + 3H2O+ 3CO2
In the case of M2+
[M(H2O)6]2+ + CO32−  ⇌ MCO3 + 6H2O

 

 

 

[Fe(H2O)6]2++ CO32− ⇌  FeCO3 + 6H2O      (green solution → green ppt)
[Co(H2O)6]2++ CO32− ⇌  CoCO3 + 6H2O      (pink solution → pink/purple ppt)
[Cu(H2O)6]2++ CO32− ⇌  CuCO3 + 6H2O      (blue solution → blue-green ppt)

2[Al(H2O)6]3++3CO32− 2[Al(H2O)3(OH)3] + 3H2O+ 3CO2  (colourless solution → white ppt.and effervescence)
2[Cr(H2O)6]3++3CO32− ⇌ 2[Cr(H2O)3(OH)3] + 3H2O+ 3CO2     (ruby/violet solution → green ppt.and effervescence)
2[Fe(H2O)6]3++3CO32− ⇌ 2[Fe(H2O)3(OH)3] + 3H2O+ 3CO2  (brown solution → brown ppt.and effervescence)


Iron and copper chemistry - Iron and copper were first encountered during AS Chemistry’s ‘Extraction of Metals‘ topic. The link between both metals being exemplified by the use of scrap iron to extract copper from copper (II) ions: Fe(s) + Cu2+(aq) → Cu(s) + Fe2+(aq) (blue)                        (green) This method of copper extraction, usually from low-grade
Chromium, cobalt, vanadium and manganese chemistry - At lower oxidation states, transition metals form ionic bonds e.g. Mn2+ and Cr3+  At higher oxidation states, transition metals cannot form monatomic ions. Instead they bond covalently with electronegative elements (like oxygen) to form compounds or molecular ions e.g. oxoanions such as MnO4−and Cr2O72-. In general,  Reduction, from high to low metal oxidation state, is more feasible in acidic transition

Reactions with ethane-1,2-diamine

a)    H2NCH2CH2NH2 acting as a ligand:
[Co(NH3)6]2+ + 3H2NCH2CH2NH→ [Co(H2NCH2CH2NH2)3]2+ + 6NH3

The enthalpy change for the above reaction is close to zero, as the same number, and same type of bonds, are broken (energy required) and made (energy released). However, the entropy change for the reaction is positive, as the reaction converts 4 to 7 particles. Consequently, the overall reaction is associated with a negative Gibbs free energy (ΔG = ΔH-TΔS) and is therefore energetically favoured.

b)   H2NCH2CH2NH2 acting as a base:
2[Al(H2O)6]3+ + 3H2NCH2CH2NH→2[Al(H2O)3(OH)3] + 3[H3NCH2CH2NH3]2+

Acting as a base, ethane-1,2-diamine reacts with [Al(H2O)6]3+ to form a white precipitate of [Al(H2O)3(OH)3] and [H3NCH2CH2NH3]2+. The +2 charge arises due to each nitrogen atom acquiring a relative +1 charge by donating a lone pair to the proton that originates from the water ligand of the aluminium complex. 

Similarly, acting as a base, ethane-1,2-diamine reacts with excess HCl to yield [H3NCH2CH2NH3]2+ 2Cl


Required practical
Carry out simple test-tube reactions to identify transition metal ions in aqueous solution.