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There are numerous possible ways to define and acid and a base. One common
way is the Lewis acid base model:
- An acid is a electron pair acceptor.
- A base is a electron pair donor.
This definition is more broad
than the Bronsted-Lowery
definition. Obviously, H+ is an acid and OH- is a base
under either definition, since for an proton to bind to a base, it must accept a
pair of electrons
- H+(aq) + :OH-(aq) -> H2O(l)
However, the Lewis definition extends beyond just the proton. For
example, many metal ions can act as Lewis acids when they form complex ions
- Fe+3(aq) + 6CN-(aq) ->
Fe(CN)6-3(aq)
Here, the electrons which form the
bond between the iron and the cyanide ion start as lone pairs on the cyanide.
The iron accepts the electrons and is thus an electron pair acceptor and a Lewis
acid, the cyanide ions donate a pair of electrons are are thus Lewis bases. Note
that there are no protons in this reaction, but it is still an acid/base
reaction.
Example: For the reaction below, identify the acid and the base
- Cu+2(aq) + 4NH3(aq) ->
Cu(NH3)4+2(aq)
Solution: Cu+2 accepts four pairs of electrons from the
ammonia molecules, so it is an acid. The ammonia molecules are donating their
lone pairs, so the ammonia molecules are the bases. |