The Henderson-Hasselbalch equation describes the pH of a buffer solution in terms of the acid dissociation constant (pKa) and the concentrations of the conjugate acid and base. It relates pH to the log of the concentration ratio of the conjugate base to the conjugate acid. The equation was derived from acid-base equilibrium chemistry and is useful for estimating pH and studying metabolic processes involving buffers. It has limitations for very strong or dilute solutions where dissociation reactions are not negligible.

1. HENDERSON HASSELBALCH
EQUATION
Prepared by :-
Rajina Shakya
Bachelor In Pharamcy
4th year, 7th semester

2. The Henderson–Hasselbalch
Equation
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 (CH 3COOH)
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-] = 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),