Acid and Bases
Arrhenius Acids and Bases:
The Arrhenius definition of acids and bases is one of the oldest. An Arrhenius acid is a
substance that when added to water increases the concentration of H1+ ions present.
The chemical formulas of Arrhenius acids are written with the acidic hydrogens first.
An Arrhenius base is a substance that when added to water increases the
concentration of OH1- ions present. HCl is an example of an Arrhenius acid and NaOH
is an example of an Arrhenius base.
The H1+ ion produced by an Arrhenius acid is always associated with a water molecule
to form the hydronium ion, H3O1+(aq). Arrhenius acids are frequently referred to as
proton donors, hydrogen ion donors, or hydronium ion donors. To represent the transfer
of the H1+ ion to water to form the hydronium ion, we must include H2O in the chemical
equation for acid ionization.
Brønsted-Lowry Acids and Bases:
A Bronsted-Lowry acid is defined as anything that releases H1+ ions; a BronstedLowry base is defined as anything that accepts H1+ ions. This definition includes all
Arrhenius acids and bases but, as we will soon see, it is a bit more general.
The Bronsted-Lowry concept is based on the transfer of a proton from one substance to
another. The ionization of HCl in water can be viewed as a Bronsted-Lowry acid-base
reaction with HCl behaving as the acid (H1+ ion donor) and water serving as the base
(H1+ ion acceptor). Just as we saw with Arrhenius acids, the acidic hydrogens are
usually written at the beginning of the chemical formula of a Bronsted-Lowry acid. So,
for instance, we can tell from the formulas that HCl, HNO3, and HCH3O2 are
monoprotic acids, each liberating one H1+ ion per acid molecule, whereas H2SO4 is
diprotic, and H3PO4 is triprotic.
The reaction of HCl with water can be called either an Arrhenius acid-base reaction or a
Bronsted-Lowry acid-base reaction. There are many acid-base reactions, however, for
which the Arrhenius definition is inappropriate. Let's consider the gas phase reaction of
HCl with ammonia, NH3, which occurs as shown in the chemical equation below. The
arrows represent the movement of electron pairs as bonds are formed and broken.
In this reaction the ammonia behaves as the proton acceptor, a Bronsted-Lowry base,
while the HCl serves as the proton donor, a Bronsted-Lowry acid. Notice that a pair of
nonbonding electrons on the base is used to form a covalent bond with the hydrogen of
the acid. The covalent bond between the hydrogen and the chlorine atom is broken, and
this pair of electrons becomes nonbonding on the chlorine atom. To function as a
Bronsted-Lowry base a substance must have a pair of nonbonding electrons that can be
used to form a bond to H1+. Be careful though, because the presence of nonbonding
electrons is not always apparent. When we write chemical formulas, the nonbonding
electrons are not shown explicitly so it may not be obvious that systems such as PH 3,
F1-, or amines are bases.
The Lewis Acid-Base Concept:
A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor. This
definition is more general than those we have seen to this point; any Arrhenius acid or
base, or any Bronsted-Lowry acid or base can also be viewed as a Lewis acid or base.
The reaction of H1+ with OH1-, for instance, involves donation and acceptance of a
proton, so it is certainly legitimate to call it a Bronsted-Lowry acid-base reaction. But if
we look at the Lewis structures for the reactants and products, we see that it is also
legitimate to call this a Lewis acid-base reaction.
The hydroxide ion donates a pair of electrons for covalent bond formation, thus OH 1- is
a Lewis base in this reaction. The hydrogen ion accepts the pair of electrons so it is
acting as a Lewis acid. Shown below is an example of a Lewis acid-base reaction that
cannot be viewed as a Bronsted-Lowry acid-base reaction.
The BF3 is the Lewis acid and the N(CH3)3 is the Lewis base. Both of the electrons in
the covalent bond formed by a Lewis acid-base reaction come from the same atom (in
the above example, the nitrogen donates both electrons). Such bonds are called
coordinate covalent bonds. If we want to emphasize the fact that a particular bond is a
coordinate covalent bond, we draw an arrow instead of a line to represent the bond in
the Lewis structure. The arrow points from the donor atom to the acceptor atom.