CELL BIOLOGY PH AND BUFFERS
pH and Buffers Laboratory
Introduction
The control of pH is important in organisms and
their cells because chemical reactions and processes
are affected by the hydrogen ion concentration. For
example, enzyme function is affected by the
ambient pH. Some side chains of amino acids (R-
groups) have ionizable groups like carboxyl or
amine groups. Changes in pH can alter the number
of positively and negatively charged groups. The
net charge on the protein effects its three
dimensional structure and thus the enzymatic
activity.
When working with living systems in a
laboratory it is necessary to pay close attention to
the pH of solutions for these same reasons. Buffers
are chemicals or combinations of chemicals that
tend to prevent changes in the concentration of
hydrogen ions. In this laboratory we will titrate
some buffered solutions to gain an understanding of
how buffers work and to observe the range of
buffering capacity. In analyzing the data, we will
make use of graphs and see the relationship of pH to
pKa using the Henderson-Hasselbalch equation.
Buffers are composed of mixtures of weak acids
and their corresponding salts. Using the Lowry-
Bronstead definition, an acid is a compound that
can donate a hydrogen ion. A weak acid is one that
does not completely ionize, or dissociate, in
solution. The extent of dissociation is given by the
equilibrium constant Ka for the reaction:
HA ↔ H+ + A–
The equilibrium constant for the ionization of
this acid is then:
Ka = [H+][A–]/[HA]
This is the measure of the ease with which the acid
donates its hydrogen ion. Higher Ka’s indicate that
the acid will dissociate more completely into ions.
The equilibrium constant equation may also be
used to determine the hydrogen ion concentration:
[H+] = (Ka)([HA]/[A–].
And, since pH = –log[H+] and
pKa = –log Ka, then,
pH = pKa + log([A–]/[HA])
This last equation is the Henderson-Hasselbalch
equation. It reveals the relationship of pH and pKa.
Remember that pH is a measure of the acidity of a
solution while pKa is a measure of the affinity of a
molecule for its proton. The former can be altered
by various means, but the latter is an inherent
property of a molecule. See the following table for
pKa’s of some biologically important substances.
pKa Values of selected acids at 25°C
Acetic acid 4.76
Carbonic acid 6.36 10.24
Citric acid 3.13 4.76 6.40
Glutamic acid 2.16 4.27 9.36
Glycine 2.35 9.77
Lysine 2.18 8.95 10.53
Phosphoric acid 2.18 7.20 12.40
Succinic acid 3.46 5.10
The Henderson-Hasselbalch equation permits the
calculation of the pH from the ratio of the salt to the
acid. Notice that the pH is equal to the pKa when
the salt and the acid are equal in concentration. If
H+ is added to such a solution, the acid concentra-
tion increases while the salt concentration decreases
and the pH is lowered. Conversely, if OH– is
added, the salt concentration increases while the
acid concentration decreases and the pH of the
buffer mixture is raised.
The assignment of the appropriate pKa to the
correct step in the ionizaton of a polybasic acid can
be thought of in the following terms. The first H+ to
be donated by an acid is the one which is most
weakly bound and is the one which will be released
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