Buffering against pH Changes in Biological Systems: -A Simple Expression Relates pH, pKa, and Buffer Concentration

The titration curves of acetic acid, H₂PO₄^-, and NH⁺₄ (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

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 product [H+] [OH-] transiently exceeds Kw (1x10 ^-14 M2). The equilibrium quickly adjusts so that this product equals 1 10 14 M2 (at 25 C), thus transiently reducing the concentration of H+. But now the quotient [H+][Ac-] / [HAc] is less than K_a, so HAc dissociates further to restore equilibrium. Similarly, the conjugate base, Ac-, can react with H ions added to the system; again, the two ionization reactions simultaneously come to equilibrium. Thus a conjugate acid-base pair, such as acetic acid and acetate ion, tends to re sist a change in pH when small amounts of acid or base are added. Buffering action is simply the consequence of two reversible reactions taking place simultaneously and reaching their points of equilibrium as governed by their equilibrium constants, K_W and K_a.

Hasselbalch equation, which is important for understanding buffer action and acid-base balance in the blood and tissues of vertebrates. This equation is sim ply a useful way of restating the expression for the dissociation constant of an acid. For the dissociation of a weak acid HA into H+ and A-, the Henderson Hasselbalch equation can be derived as follows:

Now invert -log [HA]/[A-], which involves changing its sign, to obtain the Henderson-Hasselbalch equation:

Stated more generally,

This equation fits the titration curve of all weak acids and enables us to deduce a number of important quantitative relationships. For example, it shows why the pKa of a weak acid is equal to the pH of the solution at the midpoint of its titration. At that point, [HA] equals [A-], and pH=pKa + log 1= pKa+0=pKa As shown in Box 2–3, the Henderson-Hasselbalch equation also allows us to (1) calculate pK_a, given pH and the molar ratio of proton donor and acceptor; (2) calculate pH, given pK_a and the molar ratio of proton donor and acceptor; and (3) calculate the molar ratio of pro ton donor and acceptor, given pH and pK_a.

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مواضيع ذات صلة

Amino Acids: - Amino Acids Share Common Structural Features

Buffering against pH Changes in Biological Systems:- Buffers Are Mixtures of Weak Acids and Their Conjugate Bases

Ionization of Water, Weak Acids, and Weak Bases: -Titration Curves Reveal the pKa of Weak Acids

Ionization of Water, Weak Acids, and Weak Bases: -Weak Acids and Bases Have Characteristic Dissociation Constants

Ionization of Water, Weak Acids, and Weak Bases: -The pH Scale Designates the H+ and OH- Concentrations

Ionization of Water, Weak Acids, and Weak Bases: -The Ionization of Water Is Expressed by an Equilibrium Constant

Ionization of Water, Weak Acids, and Weak Bases: - Pure Water Is Slightly Ionized

Weak Interactions in Aqueous Systems: -Solutes Affect the Colligative Properties of Aqueous Solutions

Weak Interactions in Aqueous Systems: -Weak Interactions Are Crucial to Macromolecular Structure and Function

Weak Interactions in Aqueous Systems: - van der Waals Interactions Are Weak Interatomic Attractions

Weak Interactions in Aqueous Systems: -Nonpolar Compounds Force Energetically Unfavorable Changes in the Structure of Water

Weak Interactions in Aqueous Systems: -Entropy Increases as Crystalline Substances Dissolve