The Henderson-Hasselbalch equation describes the relationship between pH and the concentrations of an acid and its conjugate base at chemical equilibrium in aqueous solutions. It allows calculation of pH from the acid dissociation constant (pKa) and the logarithm of the concentration ratio between the conjugate base and acid. The equation is widely used to determine pH values in biological and chemical systems like proteins and buffer solutions.

1. HENDERSON HASSELBALCH
EQUATION
Prepared by :-
 Rabita Maharjan
 Rajina Shakya

2. The Henderson–Hasselbalch
Equation
Describes the derivation of pH as a measure of
acidity in
biological and chemical systems.
 The equation is also useful for estimating the pH of
a buffer solution.
it is widely used to calculate the isoelectric point of
proteins( point at which protein neither accept nor
yield
proton) .

3. The Henderson hasselbalch equation for acid is :-
pH = pKa + log [ Aˉ ]
[HA]
Here, pKa= -log(Ka)
where Ka is the acid dissociation constant, that is
pKa= -log [H3O+][A-]
[HA]
for the non-specific Brønsted acid-base reaction:
HA + H20 A- + H3O+
( Acid ) ( Conjugate base )

4. The Henderson Hasselbalch Equation for base is
:
pOH = pKb + log [ BH+ ]
[B]
where BH+ denotes the conjugate acid of the
corresponding base B.
B + H2O+ BH + OH-
(Base ) (Conjugate acid)

5. History
- Lawrence Joseph Henderson wrote an
equation, in
1908, describing the use of carbonic acid as a
buffer
solution.
- Karl Albert Hasselbalch later re-expressed that
formula in logarithmic terms, resulting in the
Henderson–Hasselbalch equation.
- Hasselbalch was using the formula to
study metabolic acidosis.

6. Henderson-Hasselbalch Equation
Derivation:
-According to the Brønsted-Lowry theory of acids
and
bases, an acid (HA) is capable of donating a
proton
(H+) and a base (B) is capable of accepting a
proton.
-After the acid (HA) has lost its proton, it is said to
exist as the conjugate base (A-).
-Similarly, a protonated base is said to exist as the
conjugate acid (BH+).

7. The dissociation of an acid can be
described by an equilibrium expression:
HA + H20 H3O+ + A-
Consider the case of acetic acid
(CH3COOH) and acetate anion (CH3COO-):
CH3COOH + H2O CH3COO- + H3O+

8. Acetate is the conjugate base of acetic acid.
Acetic acid and acetate are a conjugate acid/base
pair. We can describe this relationship with an
equilibrium constant:
Ka = [H3O+][A-]
[HA]
Taking the negative log of both sides of the
equation
gives
-logKa = -log [H3O+][A-]
[HA]
or, -logKa = -log [H3O+] + (-log [A-] )

9. By definition,
pKa = -logKa and pH = -log[H3O+], so
pka=pH – log [A-]
[HA]
This equation can then be rearranged to
give the Henderson-Hasselbalch equation:
pH = pKa + log [A-] = pKa + log [conjugate
base]
[HA] [acid]

10. Estimating blood pH
A modified version of the Henderson–Hasselbalch
equation can be used to relate the pH of blood to
constituents of the bicarbonate buffering system.
pH = pKaH2CO3 + log [HCO3-]
[H2CO3]
, where:
-pKa H2CO3 is the acid dissociation
constant of carbonic acid. It is equal to 6.1.
[HCO3-] is the concentration of bicarbonate in the
blood
[H2CO3] is the concentration of carbonic acid in the
blood

11. Limitation :
-The most significant is the assumption that the
concentration of the acid and its conjugate base
at
equilibrium will remain the same.
-This neglects the dissociation of the acid and the
hydrolysis of the base.
-The dissociation of water itself is neglected as
well.
-These approximations will fail when dealing with:
relatively strong acids or bases
dilute or very concentrated solutions (less than
1mM or greater than1M),