Chemistry Reference
In-Depth Information
It is important to notice that water appears in these equations as both a
proton acceptor and a proton donor. This is an example of the amphoteric
(sometimes termed the amphiprotic) nature of water. Although the ionisa-
tion of acids and bases in water is best described using the equations above,
it is convenient to disregard the water when deriving useful expressions and
relationships.
Consider any weak acid, HA, which dissociates as shown below:
HA
3
H
A
The equilibrium constant for this reaction is given, as before, by
[H
]
[A
]
K
—------—
[HA]
In the case of an acid dissociation, the equilibrium constant for the reaction
is termed
K
a
, and is called the ionisation constant, the dissociation constant
or, sometimes, the acidity constant. The above equation can now be
rewritten as
[H
]
[A
]
K
a
----------
[HA]
For exact work, the concentration term must be replaced by the thermo-
dynamic activity of the ion, but for dilute solutions concentration may be
used.
K
a
is a constant for a given compound at a given temperature. Clearly,
the farther the above equilibrium lies to the right-hand side, the more
completely the acid will ionise and the greater will be the value of
K
a
.
To put it more simply, the greater the value of
K
a
, the stronger is the
acid. Using the equation above, it is possible to derive an expression for the
strength of acid solutions. If the acid, HA, ionises to
a
moles of H
ions
and
a
moles of OH
ions, where
a
is the fraction of the acid that is ionised,
then the number of moles of undissociated acid is given by (1
a
). This
acid solution can now be prepared with
c
moles of acid in 1 litre (or 1 dm
3
),
which will yield
ac
moles of H
and
ac
moles of A
. Hence,
HA
3
H
A
(1
a
)
c ac
ac
ac
ac
K
a
——---
(1
a
)
c
