The Henderson-Hasselbalch equation describes the relationship between pH and the concentrations of an acid and its conjugate base at equilibrium. Specifically, it shows that pH equals the acid's pKa plus the logarithm of the concentration of the conjugate base divided by the concentration of the acid. This equation is useful for estimating pH in buffer solutions and calculating the isoelectric point of proteins. It is derived from acid/base equilibrium chemistry and the acid dissociation constant Ka.

1. HENDERSON HASSELBALCH
EQUATION
By
Mrs Sanchita Choubey
(M.Sc., PGDCR, Pursuing Ph. D)
Assistant Professor of Microbiology
Dr. D Y Patil Arts Commerce and Science College Pimpri, Pune

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:
A-
HA + H20 + H3O+
( Acid ) ( Conjugate

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. Histoy
- 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-] )
[HA]

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-]
[HA]
= pKa + log [conjugate
base]
[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

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),