Theories of acids and bases
HF + NH3 ↔ NH4+ + F-
The HF transfers a H+ to the NH3, so it acts as an
acid; the NH3 accepts the H+ so it acts as a base.
In Bronsted-Lowry theory, an acid can only behave as a proton
donor if there is also a base present to act as a proton acceptor.
Consider the acid-base reaction between a generic acid HA and
HA + B ↔ A- + BH+
We can see that HA acts as an acid, while B acts as a base. But, if we
look at the reverse reaction, BH+ acts as an acid (donates a proton)
and A- acts as a base (accepts a proton).
So, acids react to form bases and bases react to form acids. Acid-
base pairs that are related to each other in this way are called
conjugate acid-base pairs.
Conjugate acid-base pairs differ by one proton.
One example of a conjugate pair is H2O and H3O+, which
is found in all acid-base reactions in aqueous solution.
H2O + H+ ↔ H3O+ -- this reaction takes place when an
acid dissolves in water. H3O+ is called the hydroxonium
ion, the oxonium ion, or the hydronium ion.
For most reactions it is convenient to simply write H+
So, H3O+ is the conjugate acid and H2O is the conjugate
Water can act as a base or as an acid, depending what it
is reacting with.
Substances that can act as acids or bases
are said to be amphoteric or amphiprotic.
To act as a Bronsted-Lowry acid, substances must be
able to dissociate and release H+.
To act as a Bronsted-Lowry base, substances must be
able to accept H+, which means, they must have a lone
pair of electrons.
Example: HCO3- acts as acid and base
HCO3-(aq) + H2O(l) ↔ CO32-(aq) + H3O+(aq)
HCO3-(aq) + H2O(l) ↔ H2CO3(aq) + OH-(aq)
Lewis acids and Bases
The Lewis definition of acids is broader than the Bronsted-Lowry
By the Lewis definition, an acid is any species capable of accepting a
lone pair of electrons (no longer restricted to just H+).
Lewis acid-base reactions result in the formation of a covalent bond
which will always be a dative bond, because both electrons come from
BF3 has an incomplete octet, so it is able to act as a Lewis acid by
accepting a pair of electrons; NH3 acts as a Lewis base by donating
its lone pair of electrons. The arrow shows the bond is a coordinate
covalent bond (dative bond)
Other good examples of Lewis acid-base reactions are found
in the chemistry of transition elements.
Transition metals often form ions with vacant orbitals. Thus
they are able to act as Lewis acids and accept lone pairs of
The Lewis base, which donates the lone pair, is called a ligand,
and usually surrounds the ion in a fixed number ratio.
Dative bonds form between the metal ion and the ligands,
resulting in a complex ion that has a characteristic color.
Typical ligands found in complex ions include H2O, CN- and
NH3. Note that these all possess lone pairs of electrons.
Theory Definition of acid Definition of base
Bronsted-Lowry Proton donor Proton acceptor
Lewis Electron pair acceptor Electron pair donor
Although all Bronsted-Lowry acids are Lewis acids, not all Lewis
acids are Bronsted-Lowry acids, so the term Lewis acids are usually
reserved for those species which can only be described by Lewis
theory (that is those that do not release H+).