Definition, Ostwald’s dilution law, dissociation theory, ionization of water, common-ion effect, ionization constants of acid and bases, ionization of polyprotic electrolytes, solubility products and its application in pharmacy.

1. Written By
Department of Pharmacy
University of Rajshahi
Rajshahi-6205, Bangladesh

2. ,
a-
Kc=-
v
For weak electrolytes
For weak electrolytes. the value of a is very small as compared to I, so we can take I - a == I. Thus the Ostwald's
dilution law expression becomes
l..~ ·- -
'a:
Kc~ (ii)
(I - a)V
This expression which correlates the variation of the degree of dissociation of an electrolyte with dilution is known as
Ostwald's dilution law.
The equilibrium constant Kc is called the dissociation·constant or ionization constant.
Ifone mole of an electrolyte can be dissolved in V litre of the solution. then ·
I
C= -
v
V is known as the dilution for the .solution. Thus the expression (i) becomes:
a2C
==> Kc= mol litre (i)
( 1-(l')
At equilibrium:
The Ostwald's dilution law
Ostwald noted that the law of mass action.can be applied to the ionic equilibrium as in the case ofchemical equilibrium.
Let us consider a binary electrolyte AB which dissociates in solution to form the ions A' and B"", The ionic equilibrium
of this reaction can be representing a~:
AB ¢::::: A• + ff
Let C moles per litre be the concentration of the electrolyte and a is its degrre of'dissociation.The concentration terms at
equilibrium may be written as: ;'. -: ):-' ;, ·
IAlli = C (I - a) mo! litre·
IA'j = Ca mol litre·
IB-1 = Ca mol litre'
Applying the law of mass action:
. Rate ofdissociation c: k, x C (I -- a)
Rate of combination= k2 x Ca x Ca
k1 x C (I - a)= k2 x Ca x Ca
.co«c« k1
==> - = K<
C(l-a) k.
J
Non-electrolytes: Non-electrolytes arc those substances which do not give any ion in solvent. For example. glucose ~ocs
not give any ion in solution.
NaCl ¢::::: Na' + Cr
.···.-:,
Electroly~f~; Electrolytes arc those substances which give ion in solvent. For example NaCl gives Na' and CT ions in
aqueous so'lu;f(p1.
.'
Ionic equilibrium: When an electrolyte dissociates into ions in water solutions and when the ions arc in a state of
equilibrium with the undissociated molecules then this equilibrium is called the ionic equilibrium. ·
Let us consider a binary electrolyte AB which dissociates in solution to form the-ions A' and s-: The ionic equilibrium;,(
this reaction can be representing as:
AB ¢::::: A' + ff
Ionic equilibria
J ()I·...;-.
Page 1 of 2
Md.
Imran
Nur
Manik

3. .....
AB <= A' + B"
By the addition of the salt (AC). the concentration of k increases. Therefore according to Le Chatelier's principle. the
equi Iibrium wiII shift to the left. thereby decreasing the concentration of A' ions or that the degree ofdissociation of AB wil I
be reduced.
The reduction of the degree of dissociation of a salt by the addition of a common-ion is called the common-ion
effect.
Let us consider a few examples to illustrate the common-ion effect.
Example I: In a saturated solution ofsilver chloride, we have the equilibrium
AgCl1s1 ¢::::: Ag•<••i> + Cr1.q1
When NaCl is added to the solution. the concentration of Cr ions will increase. The equilibrium shown above will be
shifted to the left to form more of AgCI. Thus the solubility of AgCI will decrease.
Example 2: When solid NH4CI is added to NH40H solution. the equilibrium shifts to left. Thereby the equilibrium
concentration of Off decreases. This procedure of reducing the concentration of OH' ions.
NH~OH ¢::::: NH/ + Off
Example 3: The common-ion effect can also be applied to the ionic equilibrium of a weak acid as HF
HF ¢::::: H' + F.
The common-ion effect
When a soluble salt (A'C) is added to a solution of another salt (A,B') containing a common ion (A'). the dissociation
of AB is expressed as
Degree of dissociation
When a certain amount of electrolyte (;/ · B") isdissolve in water. a small fraction ofit dissociates to form ions (A' & ff).
When the equilibrium has been reached be~een the undissociated and the free ions. we have, '· "'< • '. ~.
AB <= A+ + B'
The fraction of the amount ofthe electrolyte in solution present as free ions is called the degree of dissociation.
If the degree of dissociation is represented by x, we can write
amount dissociated (mo/ IL)x ~ ~~~~~~~~~~~~-
initiaI concentration (1110/ IL)
Limitation
i. It based on Arrhenius theory that is applicable in strong and weak electrolytes. But normal electrolytes dissociate
only in a fraction. where the law is not applicable.
ii. If the concentration of ion is very high. the pressure of ions affects on the equilibrium. For that reason law of mas-..
action is not being applied.
iii. The ions obtained by diss~iation may get hydrated and may effects the concentrationterms,
•
Kc~-----
(I- a)V
::) a~ = KcV - aKcV
which gives a quadratic equation.
a~+ aK1Y- KcY-= 0
from this equation the value of a can be calculated.
::) n > ~ KcV
::) a= K1 JV
That is. the degree of dissociation of a weak electrolyte is proportional to the square root of the dilution.
For strong electrolytes
For strong electrolytes, the volume ofa is large and it cannot be neglected in comparison with I. Thus we have to use the
expression
s,
' ·.. '.
< • ,
"""'o' •
- ~- ---
Page 2 of 2
Md.
Imran
Nur
Manik