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Summary Chapter 2.3

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Course
Biochemistry I (BIOL2080U)

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University Of Ontario Institute Of Technology (UOIT )

Summary study book Lehninger Principles of Biochemistry of Nelson David L., Albert L. Lehninger, David L. Nelson, Michael M. Cox, University Michael M Cox (2.3) - ISBN: 9780716743392 (Chapter 2.3)

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2.3 Buffering against pH Changes in Biological Systems

SUMMARY 2.3 Buffering against pH Changes in Biological Systems

■ A mixture of a weak acid (or base) and its salt resists changes in pH caused by the
+ −
addition of H or OH . The mixture thus functions as a buffer.

■ The pH of a solution of a weak acid (or base) and its salt is given by the Henderson-
Hasselbalch equation:

■ In cells and tissues, phosphate and bicarbonate buffer systems maintain intracellular and
extracellular fluids at their optimum (physiological) pH, which is usually close to 7. Enzymes
generally work optimally at this pH.

■ Medical conditions that lower the pH of blood, causing acidosis, or raise it, causing
alkalosis, can be life threatening.

Almost every biological process is pH-dependent; a small change in pH produces a large
change in the rate of the process. This is true not only for the many reactions in which the
+
H ion is a direct participant, but also for those reactions in which there is no apparent role
+
for H ions. The enzymes that catalyze cellular reactions, and many of the molecules on
which they act, contain ionizable groups with characteristic pKa values. The protonated
amino and carboxyl groups of amino acids and the phosphate groups of nucleotides, for
example, function as weak acids; their ionic state is determined by the pH of the
surrounding medium. (When an ionizable group is sequestered in the middle of a protein,
away from the aqueous solvent, its pKa, or apparent pKa, can be significantly different from
its pKa in water.) As we noted above, ionic interactions are among the forces that stabilize a
protein molecule and allow an enzyme to recognize and bind its substrate.

Cells and organisms maintain a specific and constant cytosolic pH, usually near pH 7,
keeping biomolecules in their optimal ionic state. In multicellular organisms, the pH of
extracellular fluids is also tightly regulated. Constancy of pH is achieved primarily by
biological buffers: mixtures of weak acids and their conjugate bases.

Buffers Are Mixtures of Weak Acids and Their

Conjugate Bases

Buffers are aqueous systems that tend to resist changes in pH when small amounts of acid
+ –
(H ) or base (OH ) are added. A buffer system consists of a

weak acid (the proton donor) and its conjugate base (the proton acceptor). As an example, a
mixture of equal concentrations of acetic acid and acetate ion, found at the midpoint of the
titration curve in Figure 2-17, is a buffer system. Notice that the titration curve of acetic acid
has a relatively flat zone extending about 1 pH unit on either side of its midpoint pH of 4.76.

, + –
In this zone, a given amount of H or OH added to the system has much less effect on pH
than the same amount added outside the zone. This relatively flat zone is the buffering
region of the acetic acid–acetate buffer pair. At the midpoint

of the buffering region, where the concentration of the proton donor (acetic acid) exactly
equals that of the proton acceptor (acetate), the buffering power of the system is maximal;
+ –
that is, its pH changes least on addition of H or OH . The pH at this point in the titration
curve of acetic acid is equal to its pKa. The pH of the acetate buffer system does change
slightly when a small

+ –
amount of H or OH is added, but this change is very small compared with the pH change
+ –
that would result if the same amount of H or OH were added to pure water or to a solution
of the salt of a strong acid and strong base, such as NaCl, which has no buffering power.

Buffering results from two reversible reaction equilibria occurring in a solution of nearly
equal concentrations of a proton donor and its conjugate proton acceptor. Figure 2-19
+ −
explains how a buffer system works. Whenever H or OH is added to a buffer, the result is
a small change in the ratio of the relative concentrations of the weak acid and its anion and
thus a small change in pH. The decrease in concentration of one component of the system is
balanced exactly by an increase in the other. The sum of the buffer components does not
change, only their ratio changes.

Each conjugate acid-base pair has a characteristic pH zone in which it is an effective buffer
(Fig. 2-18). The pair has a pKa of 6.86

and thus can serve as an effective buffer system between approximately pH 5.9 and pH 7.9;
the pair, with a pKa of 9.25, can act as a buffer

between approximately pH 8.3 and pH 10.3.

The Henderson-Hasselbalch Equation Relates pH,

pK , and Buffer Concentration
a

The titration curves of acetic acid, , and (Fig. 2-18) have nearly identical shapes, suggesting
that these curves reflect a fundamental law or relationship. This is indeed the case. The
shape of the titration curve of any weak acid is described by the Henderson-Hasselbalch
equation, which is important for understanding buffer action and acid-base balance in the
blood and tissues of vertebrates. This equation is simply a useful way of restating the
expression for the ionization constant of an acid. For the ionization of a weak acid HA, the
Henderson-Hasselbalch equation can be derived as follows:

FIGURE 2-19 The acetic acid–acetate pair as a buffer system. The system is capable
+ −
of absorbing either H or OH through the reversibility of the dissociation of acetic acid.
+
The proton donor, acetic acid (HAc), contains a reserve of bound H , which can be released
−
to neutralize an addition of OH to the system, forming H2O. This happens because the

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