Acidity and Basicity - Fundamental Concepts

Chemistry 4790 - Dr. T.H. Black

    Few concepts are as important to the medicinal chemist as that of the acidity and basicity of molecules introduced into the body. The human body comprises an immensely complex set of chemical equilibria, all of which are occurring in body compartments of varying acidity, so the ability to predict the interaction of acidic or basic drugs with these environments is essential to the study of drug action and absorption.

    The most common acid-base theory used by medicinal chemists is the Bronsted-Lowry definition, which is concerned with a molecule's ability to donate or accept protons. The strength of an acid is defined as the position of the following equilibrium expression,

where HA refers to an undissociated acid molecule. A- is the conjugate base of HA, or, the basic species which results from the loss of a proton by an acid (it is basic because it can accept a proton and regenerate the parent acid). The farther to the right the equilibrium lies (indicating more complete dissociation), the stronger the acid.

The position of equilibrium is measured by the equilibrium constant Ka (as usual, the brackets indicate concentration):

Thus, we can take the negative log of both sides of eqn. 2 and obtain

which can be reduced to

or

Equation 6 is commonly referred to as the Henderson-Hasselbach equation, and describes the degree of dissociation of an acid when placed in an environment of known acidity.

    The treatment for bases is similar, except for a very important differentiation: base strength is commonly measured as the pKa of its conjugate acid. The equilibrium expression for basicity would be

in which BH⁺ is the conjugate acid of base B. Thus, to avoid confusion, the Henderson- Hasselbach equation is often stated as follows:

This form is useful for either acids or bases.

    From equation 2, it is obvious that the stronger the acid, the larger Ka will be, and thus the smaller its pKa will be. However, since base strength is measured as the strength of the conjugate acid, stronger bases have larger pKa values. This is easily conceptualized if it is realized that the equilibrium of a weaker base will lie farther to the left (eqn. 7), making BH⁺ a more strongly acidic species, which of course would then have a lower pKa.

Utilization of the Henderson-Hasselbach Equation

    Why is this concept so essential? Remember that drug molecules must cross lipid-like membranes in the body if they are to reach their sites of action. Since the passive transport across these lipid bilayers is much like dissolution, the drug molecules which are most lipophilic (literally, "lipid-loving"; nonpolar) will pass most easily. Obviously, charged species are not particularly lipophilic relative to uncharged ones. Thus, undissociated acids and unprotonated bases are most lipophilic and absorb most easily, while dissociated acids (conjugate bases) and protonated bases (conjugate acids), which possess charges, absorbed only very slowly.

    It is therefore very useful to determine the extent of ionization (for acids) or protonation (for bases) for a given in vivo environment, so that the degree of absorption can be assessed. For example, if an acidic drug with a pKa of 5 were ingested into the stomach (pH 2), its degree of ionization would be (using equation 8):

so that the number of undissociated, lipophilic molecules would be 1000 times the number of ionic, conjugate base ions. Obviously, thus molecule would absorb easily from the stomach.

    Alternatively, a basic molecule with the same pKa value would, in the same environment, be protonated 1000 times more than not, thus, since its protonated form is ionic, it would absorb only poorly from the stomach.

    Of course, the preceding discussion is useless if you cannot distinguish between an acidic or basic molecule. Typically, Bronsted-Lowry acids contain the functional groups COOH, OH, SH, and others wherein a proton is bound to an electronegative atom. You should already be aware of the relative acid strengths of these functionalities. Bases are almost always amines, and aromatic amines are generally weaker (lower pKa) than nonaromatic counterparts.

    I would suggest that you peruse the table of drug pKa values found at the back of the text by Delgado & Remers (pg. 915), look up the structures for a dozen or so, and try to reconcile the reported pKa values with the assigned functional groups. Then, try to determine qualitatively where the drug would be best absorbed - from the stomach (pH 1-3) or from the intestine (pH 5-8). In this way, you will obtain the intuitive feel for pKa values of both acids and bases which is so essential to this area of chemistry.