Biological Acids and the Henderson-Hasselbalch Equation

Assume you have a mixture of naphthalene and benzoic acid that you want to separate. How might you take advantage of the acidity of one component in the mixture to effect a separation  

The for dichloroacetic acid is 3.32 x 10 . Approximately what percentage of the acid is dissociated in a 0.10 M aqueous solution  

In acidic solution, at low pH, a carboxylic acid is completely undissociated and exists entirely as RCO2H. In basic solution, at high pH, a carboxylic acid is completely dissociated and exists entirely as RC02. Inside living cells, however, the pH is neither acidic nor basic but is instead buffered to nearly neutral pH—in humans, to pH = 7.3, a value often referred to as physiological pH. In what form, then, do carboxylic acids exist inside cells The question is an important one for understanding the acid catalysts so often found in biological reactions. 

If the pKa value of a given acid and the pH of the medium are known, the percentages of dissociated and undissociated forms can be calculated using the Henderson-Hasselbalch equation. 

This equation says that the logarithm of the concentration of dissociated acid [A ] divided by the concentration of undissociated acid [HA] is equal to the pH of the solution minus the pKa of the acid. Thus, if we know both the pH of the solution and the pl a of the acid, we can calculate the ratio of [A ] to [HA]. Furthermore, when pH = pK, the two forms HA and A are present in equal amounts because log 1 = 0. 

As an example of how to use the Henderson-Hasselbalch equation, let s find out what species are present in a 0.0010 M solution of acetic acid at pH = 7.3. According to Table 15.2, the pK of acetic acid is 4.76. From the Henderson-Hasselbalch equation, we have 

What is true for acetic acid is also true for other carboxylic acids at the physiological pH that exists inside cells, carboxylic acids are almost entirely dissociated. To reflect this fact, we always refer to cellnlar carboxylic acids by 

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Chapter 20, Carboxylic Aciils and Nitriles—A new Section 20.3 discusses biological carboxylic acids and the Henderson-Hasselbalch equation. 

The Henderson-Hasselbalch equation provides a general solution to the quantitative treatment of acid-base equilibria in biological systems. Table 2.4 gives the acid dissociation constants and values for some weak electrolytes of biochemical interest. 

When a weakly acidic or basic drug is administered to the body, the drug will ionise to a greater or lesser extent depending on its piCa and the pH of the body fluid in which it is dissolved. The pH of the body varies widely, but the most important biological solution is the blood, which, as stated above, normally has a pH of 7.4. An equation can be derived that will predict the extent to which the drug ionises, and, as is often the case, the starting point for the derivation is the Henderson-Hasselbalch equation (1.7). 

The Henderson-Hasselbalch equation was developed independently by the Ameriean biological chemist L. J. Henderson and the Swedish physiologist K. A. Hasselbaleh, for relating the pH to the bicarbonate buffer system of the blood (see below). In its general form, the Henderson-Hasselbalch equation is a useful expression for buffer caleulations. It can be derived from the equilibrium constant expression for a dissociation reaction of the general weak acid (HA) in Equation (1.3) ... 

The degree of drug ionization depends upon both the pH of the solution in which it is presented to the biological membrane and on the pKa (dissociation constant) of the drug (whether it is an acid or base). The entire concept of pKa is derived from the Henderson-Hasselbalch equation for both acids and bases as follows ... 

This equation is known as the Henderson—Hasselbalch equation. Workers in the biological sciences use it frequendy. Because the amount of weak acid dissociated is very small, the values for [conjugate base] and [acid] are essentially their initial concentrations after mixing, ignoring the negligible reaction that they then undergo. 

June 3, 1978, Lynn, Massachusetts, USA - Feb. 10, 1942 Boston, USA) Henderson studied medicine at Harvard and was Professor of Biological Chemistry at Harvard University, Cambridge, Massachusetts, from 1904 to 1942 [i]. Henderson published on the physiological role of -> buffers [ii-vii] and the relation of medicine to fundamental science. Because he and also - Hasselbalch made use of the law of mass action to calculate the - pH of solutions containing corresponding acid-base pairs, the buffer equation is frequently (esp. in the biological sciences) referred to as -> Henderson-Hasselbalch equation. 

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