This document discusses various concepts related to ionic equilibrium in solution including strong and weak electrolytes, acid-base theories of Arrhenius, Bronsted-Lowry, and Lewis. It defines strong electrolytes as completely dissociating in water and weak electrolytes as achieving an equilibrium between dissociated and undissociated molecules. Acids are defined as proton donors and bases as proton acceptors under the Bronsted-Lowry theory. The Lewis theory further defines acids as electron pair acceptors and bases as electron pair donors. Dissociation constants and factors affecting acid strength are also covered.

1. 1
IONIC EQUILIBRIUM IN SOLUTION

2. 2
Electrolytes when dissolved in water splits up
into charged particles called ions.. The
process is called ionisation or dissociation.
Certain electrolytes such as NaCl, KCl, HCl
are almost completely ionised in solutions.
The electrolytes which are almost
completely ionised in their solutions are
called strong electrolytes .

3. 3
Strong electrolytes are:
1.All water soluble salts
(KCl,Na2SO4,Ca(NO3)2 ,etc.
2.Alkalines (NaOH, KOH, Ca(OH)2, Ba(OH)2),
etc.
3.Mineral acids (H2SO4, HNO3, HCl, HBr,
HI),etc.
The equation for dissociation of strong electrolytes
are written with only a single arrow directed to the
right.
KCl(aq) → K+
(aq) + Cl −
(aq)

4. 4
On the other hand, electrolytes which are
weakly ionised in their solutions are
called weak electrolytes . In case of
solutions of weak electrolytes, the ions
produced by dissociation of electrolyte
are in equilibrium with undissociated
molecules of the electrolyte.
NH4OH(aq) NH4
+
(aq) + OH−
(aq)
Equations for the dissociation of weak electrolytes
are written with double arrows( ).
CH3
COOH(aq) CH3
COO−
(aq) + H+
(aq)

5. 5
VARIOUS CONCEPTS OF ACIDS AND BASES
1.ARRHENIUS CONCEPT OF ACIDS AND BASES.
According to Arrhenius concept , an acid is
a substance which can furnish hydrogen
ions in its aqueous solution . A base is a
substance which can furnish hydroxyl
ions in its aqueous solution .
For example, substances such as HNO3
, HCl,
CH3
COOH etc are acids, whereas substances
such as NaOH , KOH , NH4
OH etc. are bases,
according to this concept.

6. 6
HNO3 → H+
(aq) + NO3
−
(aq)
HCl(aq) → H+
(aq) + Cl−
(aq)
CH3COOH(aq) H+
(aq) + CH3COO−
(aq)
NaOH(aq) → Na+
(aq) + OH−
(aq)
KOH(aq) → K+
(aq) + OH−
(aq)
NH4OH(aq) NH4
+
(aq) + OH−
(aq)

7. . 7
According to Arrhenius theory , neutralisation
of acids and bases is basically a reaction
between H+
and OH−
ions in solutions.
H +
+ OH− H2
O

8. 8
2.BRONSTED-LOWRY CONCEPT OF ACIDS AND
BASES.
The Brønsted-Lowry definition, formulated in 1923,
independently by Johannes Nicolaus Brønsted in
Denmark and Martin Lowry in England
It is based upon the idea of protonation of bases
through the de-protonation of acids
Johannes Nicolaus Brønsted Martin Lowry

9. 9
They proposed that : An acid is a
substance that can donate a proton. A
base is a substance that can accept a
proton .
These definitions are more general
because according to these definitions
even ions can behave as acids or bases.
Moreover, these definitions are not
restricted to reactions taking place in
aqueous solutions only.

10. 10
It is a reversible reactions that involve proton
transfer from the acid to the base
HA + B  HB+
+ A−
Acid Base
Acid is known as Proton Donor.
Base is known as Proton Acceptor.
HCl → H+
+ Cl−
Acid ( Proton Donor, donate H+
)
Base ( Proton Acceptor, accept H+
)

11. 11
Some more conjugate acid - base pairs has
been given in the following equations :

12. 12
In both Arrhenius and Bronsted concepts, acids are
sources of protons. Hence all Arrhenius acids are
also Bronsted acids. However, there is a difference in
the definition of bases. Arrhenius theory requires
base to the source of OH−
ions in aqueous medium,
but Bronsted theory requires base to be a proton
acceptor. Hence Arrhenius bases may not be
Bronsted bases. For example, NaOH is a base
according to Arrhenius theory because it gives OH−
ions in aqueous solution, but NaOH does not accept
proton as such. Hence it may not be classified as a
base according to Bronsted theory.

13. 13
Strengths of acids and bases.
Strength of an acid is measured in terms of its
tendency to lose proton whereas strength of a base
is measured in terms of its tendency to accept
proton. The conjugate base of a strong acid is a
weak base.
HCl(aq) H +
+ Cl−
(aq)
strong acid weak base

14. 14
On the other hand, conjugate base of a weak
acid is a strong base.
CH3COOH(aq) H+
(aq) + CH3
COO−
(aq)
weak acid strong base
The strength of acids or bases is
experimentally measured by determining its
ionisation or dissociation constants.

15. 15
3. THE LEWIS ACIDS AND BASES.
Although Bronsted-Lowry theory was more
general than Arrhenius theory of acids and
bases , but failed to explain the acid base
reactions which do not involve transfer of
protons. For example it fails to explain how
acidic oxides such as anhydrous CO2
, SO2
,
SO3
etc. can neutralise basic oxides such as
CaO, BaO etc. even in absence of solvent.

16. 16
Lewis proposed a more general definition
for acids and bases, which do not require
the presence of protons to explain the
acid-base behaviour.
Accoding to Lewis concept :
An acid is a substance which can accept a
pair of electrons.
A base is a substance which can donate a
pair of electrons .

17. 17
Acid-base reactions according to this concept
involve the donation of electron pair by a
base to an acid to form a co-ordinate bond.
Lewis bases can be neutral molecules such
as :
having one or more unshared pairs of
electrons. , or anions such as : −CN−
, −OH−
,
−Cl−
, etc.

18. . 18
Lewis acids are the species having
vacant orbitals in the valence shell of
one of its atoms. The following species
can act as Lewis acids.
Molecules having an atom with
incomplete octet.

19. 19
For example , BF3
and AlCl3
.

20. 20
It may be noted that all Bronsted bases are
also Lewis bases but all Bronsted acids are not
Lewis acids. Lewis bases generally contain one
or more lone pairs of electrons and therefore ,
they can also accept a proton (Bronsted base).
Thus, all Lewis bases are also Bronsted bases.
On the other hand, Bronsted acids are those
which can give a proton, for example , HCl,
H2
SO4
. But they are not capable of accepting a
pair of electrons .
Hence , all Bronsted acids are not Lewis acids.

21. . 21
THE DISSOCIATION CONSTANTS OF ACIDS (Ka
)
Strong acids dissociate almost
completely in water and therefore the
molar concentrations of H+
ions in the
solution is same as that of acid itself.
But weak acids are not completely
dissociated and relative strengths of
weak acids can be compared in terms of
their dissociation constants. For
example, the dissociation equilibrium of
an acid HA may be represented as :

22. 22
HA(aq) H+
(aq) + A −
(aq)
Applying the law of Chemical equilibrium:
Here Ka
is called dissociation constant of
the acid.

23. 23
The value of dissociation constant gives an
idea about the relative strength of the acid.
Larger the value of K a ,greater is the
concentration of H+
ions and stronger is the
acid. If dissociation constants of two acids
are known, their relative strength can be
compared. For example, consider the
following examples:
CH3
COOH(aq) H+
(aq) +CH3
COO−
(aq)

24. 24
Factors affecting acid strength
The extent of dissociation of an acid depends on the
strength and polarity of the H−A bond. In general ,
when strength of H−A bond decreases , that is , the
energy required to break the bond decreases. HA
becomes a stronger acid. Also, when the H−A bond
becomes a stronger acid. Also, when the H−A bond
becomes more polar i.e., the electronegativity
difference between the atoms H and A increases and
there is marked charge separation, cleavage of bond
becomes thereby increasing the acidity. But it
should be noted that while comparing elements in
the same group of the periodic table, H−A bond
strength is a more important factor in determining
acidity than its polar nature.

25. . 25
As the size of A increases down the
group, H−A bond strength decreases
and so the acid strength increases. For
example,

26. 26
Degree of ionisation (α) = (Number of
ions (n)) ÷ (Total number of ions and
molecules (N)).
α =

27. 27
According to Arrhenius theory of electrolyte
dissociation, the molecules of an electrolyte
in solution are constantly splitting up into
ions and the ions are constantly reuniting to
form unionized molecules. Therefore, a
dynamic equilibrium exists between ions and
unionized molecules of the electrolyte in
solution. It was pointed out by Ostwald that
like chemical equilibrium, law of mass action
can be applied to such systems also.
Ostwald’s Delution Law.

28. 28
H3
CCOOH(aq) H+
(aq) + CH3
COO−
(aq)
where:
Ka: constant of dissociation
α: degree of dissociation
C(CH3COO-
): concentrations of anions
C(H+
): concentration of cations
C(CH3COOH): concentration of
associated electrolyte.
C(1-α) Cα Cα

29. 29
For very weak electrolytes, α <<< 1, (1 - α ) =
1
K = Cα2

30. . 30
Knowing the value of Ka , it is possible to calculate
the degree of ionisation of weak acid at any
particular concentration C.
Knowing the value of Ka , it is possible to calculate
the degree of ionisation of weak acid at any
particular concentration C.
Thus, degree of dissociation
of a weak electrolyte is
proportional to the square
root of dilution.

31. 31
SOLUBILITY PRODUCT CONSTANT
Certain electrolytes such as BaSO4
and AgCl are sparingly
soluble in water. Even in their saturated solutions, the
concentration of the electrolytes is very low. So , whatever
little of electrolyte goes into solution, undergoes complete
dissociation (due to low concentration). Therefore , in
saturated solutions of such electrolytes solid electrolyte is in
equilibrium with the ions as represented below :
Consider a saturated solution of a salt containing the solid
salt. There are two equilibria, one between solid salt and
dissolved salt and second between the dissolved salt and its
ions.
AB A+
+ B− AB
(solid salt) (dissolved salt) (ions)

32. 32
Applying the Law of mass action to the second
equilibrium,
where K is the equilibrium constant and [AB] is the
concentration of the dissolved salt. Cross
multiplying we get
K[AB] = [A+
] [B−
]
Since the solution is saturated , the concentration of
the dissolved salt remains constant at a fixed
temperature.

33. . 33
Hence . [A+
] [B−
]= K × Constant = KSp
where KSp
is another
constant. This constant K sp is known as the solubility
product of the electrolyte. It is the maximum value of product
of concentrations of the ions of the electrolyte.
In the case of silver chloride, we have :
AgCl Ag+
+ Cl−
KSp = [Ag+
] [Cl−
]
In general , for any sparingly soluble salt Ax By which
dissociates to set up the equilibrium :
Ax
By x Ay+
y Bx−

34. 34
where Ay+
and Bx−
denote the positive and
negative ions , x and y represent the number
of these ions in the formula of the electrolyte.
The solubility product constant may be
expressed as :
KSp = [Ay+
]x
[Bx−
]y
Thus solubility product of a sparingly soluble salt at a
given temperature may be defined as the product of
the concentrations of its ions in the saturated
solution, with each concentration term raised to the
power equal to the number of times the ion occurs
in the equation representing the dissociation of the
electrolyte.

35. 35
KSp = [A+
] [B−
] = S × S = S2
Suppose at a particular temperature its solubility is
S mol L−1
. S moles of salt on ionisation give S moles
of A+
and S moles of B−
ions.
AB A+
(aq) + B−
(aq)
In general , for any sparingly soluble salt A x B y
which dissociates to set up the equilibrium :
Ax By In general , for any sparingly
soluble salt A x B y which dissociates to set up the
equilibrium :
Ax By
x Ay+
y Bx−
[Ay+
] = x S and [Bx−
] = y S

36. 36
KSp =[x S]x
[y S ]y
= xx
yy
S(x+y)
The concept of solubility product principle helps us to predict
whether a salt will precipitate or not.
Precipitation occurs : if calculated ionic product > K sp
No precipitation : if calculated ionic product < KSp
.

37. Thank You!
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