(PROTEIN) WATER MOLECULE AMINO GROUP

K a = [H + ] [A – ] / [HA]
Water and Ocular Fluids • 5
where K a is the dissociation constant; [H + ] is the molar hydrogen ion
concentration; [A – ] is the molar anion concentration; and [HA] is the
molar concentration of the undissociated (nonionized) acid. Using the
example of acetic acid, it has been found that the dissociation constant
(K a) for this acid is: 1.74 × 10 –5 . The negative logarithm or pK a of that
value = 4.76. Using the negative logarithm of the K a is convenient to
determine the pH of a solution of this acid.
Now, if you are getting a little bored with this discussion, go open a
bottle of vinegar and sniff it. The characteristic odor is acetate anion
from the acetic acid and the sour taste (take a little taste) is from the
hydrogen ions dissociated in the vinegar. Although the amount dissociated
is only 0.000009 M from the total of 0.5 M present in the vinegar,
it is potent to one’s senses! It is easily realized then that not much is
needed to produce a pH effect in biological systems including the eye.
Let us return to the pK a value for acetic acid (4.76). This value is
equivalent to the pH value of a solution of acetic acid when it is 50%
ionized. Its pH can be changed, but only with difficulty, by adding acid
or base to the weak electrolyte (i.e., the acetic acid/acetate mixture). This
is a desirable property since, by resisting pH changes, the weak electrolyte
acts as a buffer to pH change and does so by ionizing to a greater
extent whenever base enters the solution and by ionizing to a lesser
extent whenever acid enters the solution (Figure 1–5). The capacity of
the buffer is the extent to which it will absorb acid or base and depends
on its concentration in solution. The range of the buffer is the pH limit
(increased or decreased) that will be buffered and this depends on the
pK a of the buffer to act as the mid-range value. The range will extend
approximately 1 pH unit above and below the pK a of the buffer as
shown in Figure 1–6. Buffering is a very useful and necessary property of
all biological fluids both ocular and nonocular.
THE HENDERSON-HASSELBALCH EQUATION
This equation is useful in the determination of pH, pK a, and the relative
concentrations of ionized and nonionized components of a buffer. It is
important in the preparation of buffers to be used in clinical and research
laboratories and in the formulation of drugs and drug vehicles that
require buffers on topical installation of a drug (such as ocular drugs,
which are placed on the corneal surface). The equation is:
pH = pK a + log [salt] / [acid]
where pH is the negative logarithm of the hydrogen ion concentration,
pK a is the negative logarithm of the dissociation constant of the weak
electrolyte, and log [A – ]/[HA] is the logarithm of the ratio of the ionized
anion to the nonionized acid of the weak electrolyte. The derivation of
the equation may be found in a number of contemporary biochemical
texts (Lehninger, Nelson, Cox, 1993). Examples of problems involving
the Henderson-Hasselbalch equation are found at the end of this
chapter.