1. Lewis Acids “An acid is an electron pair acceptor,
A base is an electron pair donor”

2. Comparison of definitions
All Bronsted Lowry acids are also Lewis acids.
Bronsted-Lowry – Proton donor
Eg HCl → H+ + Cl-
Lewis – electron pair acceptor
H+ has no electrons, so can accept a pair to form a dative bond.
H+ + (:OH)- → H←:OH

3. Cu2+ + 6H2O → [Cu(H2O)6]2+ Cu2+←:OH2
But other Lewis acid/bases would not qualify under the Bronsted-Lowry definition.
Eg; When a water molecule acts as a ligand in a complex ion, forming a dative bond, it acts as a Lewis base.
By accepting an electron pair the transition metal is acting as a Lewis acid.
Cu2+ + 6H2O → [Cu(H2O)6]2+
Cu2+←:OH2
Acid Base

4. Acidity of complex ions
[Fe(H2O)6]2+ solutions are not noticeably acidic.
But [Fe(H2O)6]3+ is more acidic than ethanoic acid. (pKa = 2.2)
This is because Fe3+ has a higher charge density than Fe2+ and is more polarising.
Fe3+ more strongly attracts the electrons on oxygen, weakening the OH bonds, releasing protons.
[Fe(H2O)6]3+ ⇌ [Fe(H2O)5OH]2+ + H+

5. This type of reaction is often referred to as hydrolysis;
[Fe(H2O)6]3+ + H2O ⇌
[Fe(H2O)5OH]2+ + H3O+
The complex is thus acting as a Bronsted Lowry acid.
Generally the M3+ aqua ions are more acidic than the M2+.
(The same is also true for Al3+.)