A guide for the preparation and use of buffers in biological systems

This equilibrium expression can now be rearranged to
[CH 3
COOH]
[H + ] = K a
_______________
[CH 3
COO¯]
where the hydrogen ion concentration is expressed in terms of the equilibrium
constant and the concentrations of undissociated acetic acid and acetate ion. The
equilibrium constant for ionization reactions is called the ionization constant or
dissociation constant.
Henderson-Hasselbach Equation: pH and pK a
The relationship between pH, pK a
, and the buffering action of any weak acid and
its conjugate base is best explained by the Henderson-Hasselbach equation. In
biological experiments, [H + ] varies from 10 -1 M to about 10 -10 M. S.P.L.
Sorenson, a Danish chemist, coined the “p” value of any quantity as the negative
logarithm of the hydrogen ion concentration. Hence, for [H + ] one can write the
following equation:
pH = – log [H + ]
Similarly pK a
can be defined as – log K a
. If the equilibrium expression is
converted to – log then
[CH 3
COOH]
– log [H + ] = – log K a
– log ______________
[CH 3
COO¯]
and pH and pK a
substituted:
[CH 3
COOH]
pH = pK a
– log ________________
[CH 3
COO¯]
pH =
or
[CH 3
COO¯]
pK a
+ log _______________
[CH 3
COOH]
When the concentration of acetate ions equals the concentration of acetic acid,
log [CH 3
COO¯]/[CH 3
COOH] approaches zero (the log of 1) and pH equals pK a
(the
pK a
of acetic acid is 4.745). Acetic acid and acetate ion form an effective
buffering system centered around pH 4.75. Generally, the pK a
of a weak acid or
base indicates the pH of the center of the buffering region.
The terms pK and pK a
are frequently used interchangeably in the literature. The
term pK a
(“a” refers to acid) is used in circumstances where the system is being
considered as an acid and in which hydrogen ion concentration or pH is of
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