6. 6
Salt hydrolysis may be defined as the reaction of the
cation or anion of the salt with water to produce
acidic or basic solution.
Thus depending upon the relative strengths of the
acid or base produced, the resulting solution is
acidic, basic or neutral. There are four distinct types
of hydrolytic behaviour of various salts. These are :
1.Salts of strong acids and strong bases.
2.Salts of strong acid and weak bases.
3.Salts of weak acids and strong bases.
4.Salts of weak acids and weak bases.
7. 7
1. SALTS OF STRONG ACID AND STRONG BASE.
Examples are NaCl, NaNO3
, Na2
SO4
, KCl, KNO3
,
K2
SO4
, etc.
As an illustration, let us discuss the hydrolysis of
NaCl. We may write
Thus it involves only ionisation and no hydrolysis.
Further in the resulting solution [ H+
] = [ OH−
] So the
solution is neutral. Hence it may be generalized that
the salts of strong acids and strong bases do not
undergo hydrolysis and the resulting solution is
8. 8
2. SALTS OF WEAK ACIDS AND STRONG BASES.
Examples are : CH3
COONa, Na2
CO3
, K2
CO3
, Na3
PO4
etc.
As an illustration , the hydrolysis of sodium acetate
(CH3
COONa) may be represented as follows:
CH3
COONa + H2
O ↔ CH3
COOH + NaOH
or CH3
COO−
+ Na+
+ H2
O ↔ CH3
COOH + Na+
+ OH−
or CH3
COO−
+ H2
O ↔ CH3
COOH + OH−
As it produces OH−
ions , the solution of such a salt is
alkaline in nature (pH>7), litmus-blue.
9. 9
3. SALTS OF STRONG ACIDS AND WEAK BASES.
Examples are :
NH4
Cl, CuSO4
, NH4
NO3
, AlCl3
,etc.
As an illustration , the hydrolysis of NH4
Cl may
be represented as follows:
NH4
Cl + H2
O ↔ NH4
OH + HCl
or NH4
+
+ Cl−
+ H2
O ↔ NH4
OH + H+
+ Cl−
or NH4
+
+ H2
O ↔ NH4
OH + H+
As it produces H+
ions , the solution of such a
salt is acidic in character (pH<7), litmus-red.
10. 10
4. SALTS OF WEAK ACIDS AND WEAK BASES
Examples are CH3
COONH4
, (NH4
)2
CO3
, AlPO4
etc.
As an illustration, the hydrolysis of ammonium
acetate may be represented as follows:
CH3
COONH4
+ H2
O↔ CH3
COOH + NH4
OH
Or CH3
COO−
+ NH4
+
+ H2
O ↔ CH3
COOH + NH4
OH
Thus it involves both anionic and cationic
hydrolysis. The resulting solution may be neutral or
slightly acidic or basic depending upon the relative
degrees of ionisation of weak acid and weak base
produced. Hence the resulting solution is almost
neutral , (pH=7), litmus-violet.
12. 12
salt
HYDROLYTIC CONSTANT (Kh
)
The general equation for the hydrolysis of a salt (BA)
may be written as :
BA + H2
O HA + BOH
Applying the law of chemical equilibrium, we get :
13. 13
K = the equilibrium constant.Since water is present
in large excess in aqueous solution, its
concentration [H2O] may be regarded as constant so
that we have:
where K h is called hydrolysis constant.
14. 14
DEGREE OF HYDROLYSIS (h)
The degree of hydrolysis of a salt is defined as the
fraction (or percentage) of the total salt hydrolysed.
i.e.,
18. 18
Such solutions which oppose the change in
their pH on the addition of small amounts of
an acid or a base are called buffer solutions
or simply buffers. The buffer solutions have
thus reserved acidity & reserved alkalinity.
19. 19
Characteristics of Buffer solutions:
•Its pH remains unsalted either on
keeping the solution for long or on
dilution
•Its pH is very slightly changed by
the addition of small amount of
strong base or an acid.
●It has a definite pH values
20. 20
Acidic Buffer Solution.
Mixture of weak acid + its conjugate base
Eg. : aqueous mixture of ethanoic acid
(CH3COOH )+ sodium ethanoate (CH3COONa)
Sodium ethanoate dissociates fully in water.
Ethanoic acid dissociates partially in water.
An aqueous mixture of ethanoic acid and
sodium ethanoate contains a large quantity
of,
Undissociated CH3COOH (the acid)
CH3COO-
ions (the base conjugate)
21. 21
Mixture of weak base + its conjugate acid
Eg. : aqueous mixture of ammonia NH4OH
(the base) and ammonium chloride NH4Cl (the
conjugate acid).
Ammonium chloride dissociates fully in
water.
Ammonia dissociates partially in water.
The aqueous mixture contains a large
quantity of,
Undissociates NH4OH (the base)
NH +
(the conjugate acid)
Basic Buffer Solution.
22. 22
Acidic Buffer
Action.
weak a. c.b. strong
The presence of a common ions from highly
ionized
a further suppresses the ionization of weak acid .
23. 23
Thus even an adding HCl, a strong
electrolyte does not produce an
appreciable change in pH of the solution.
However, when a strong base is added,
OH-
ions are furnished by the base and
are neutralized by
and no change in pH is observed.
24. 24
Basic Buffer
Action.
Weak base. strong
conjug.acid
Similary when a few drops of strong acid (HCl) is
added, the H+
NH4OHions combines with excess of
to form feebly ionized water molecules.
25. 25
• In chemistry, the Henderson–Hasselbalch (often
misspelled as Henderson–Hasselbach) equation
describes the derivation of pH
• measure of acidity (using pKa, the acid dissociation
constant) in biological and chemical systems.
• The equation is also useful for estimating the pH of
a buffer solution and finding the equilibrium pH in
acid-base reactions (it is widely used to calculate
the isoelectric point of proteins).
• ( is the pH at which a particular molecule or surface
carries no net electrical charge).
pH value of Buffer Solution
(Henderson – Hasselbah’s
Equation).
26. 26
Consider a buffer solution containing a weak
acid
(CH3COOH)
& its highly ionized salt (CH3COONa).
The dissociation of weak acid occurs as’
Applying Law Mass Action, we
have
