Section 5.1     Drug molecules often have one or more acidic and/or basic functional groups. The charge state of these groups varies with pH and affects a range of properties, including drug binding to its target, reactivity, and solubility.

Section 5.2     Three common theories have been used to describe the behavior of acids and bases. The Brønsted–Lowry theory defines an acid as a proton donor and a base as a proton acceptor in a proton transfer reaction. The Lewis theory defines an acid as an electron pair acceptor and a base as an electron pair donor.

Section 5.3     Water molecules are both weak Brønsted– Lowry acids and weak Brønsted–Lowry bases that react with each other with an equilibrium constant K_w = [H₃O⁺][OH⁻] = 10⁻¹⁴. The relative acidity of aqueous solutions is measured using the pH scale where pH = −log [H₃O⁺]. In neutral water, [H₃O⁺] = [OH⁻] = 10⁻⁷ M and pH = 7, in acidic solutions [H₃O⁺] > [OH⁻] so pH < 7, and in basic solutions [H₃O⁺] < [OH⁻] so pH > 7.

Section 5.4     Just as the terms “acid” and “base” can mean different things depending on which theory of acids and bases is being used, the terms “acidic” and “basic” and “acidity” and “basicity” can be used in different contexts. Acidic and basic are used to refer to the properties of solutions, properties of compounds, properties of functional groups, or properties of specific X–H bonds. Acidity and basicity typically refer either to how far the pH deviates from pH 7 or to the relative strength of an acid or base.

Section 5.5     Acid dissociation constants measured in water (K_a) for H–A acids and for the conjugate acids (BH⁺) of bases are used to describe the relative strengths of acids and bases. As the range of values is very large (10⁻⁵⁰–10¹²), acid dissociation constants are expressed as pK_a = –log K_a in analogy to the pH scale. Acids (X–H bonds) with pK_a < 0 are very strong and the acid strength decreases as pK_a increases above 0. Base strength changes in the opposite direction: strong acids have weak conjugate bases and weak acids have strong conjugate bases.

Section 5.6     The acidity of X–H bonds in the same molecular context is related to the intrinsic properties of the element X. For X atoms in the same row (period), the acidity increases as the electronegativity of X increases (C–H < N–H < O–H < F–H). For X atoms in the same column (group), the acidity increases with size (H–F < H–Cl < H–Br < H–I).

Section 5.7     The acidity of X–H bonds varies with the hybridization state of the X atom. As the “s” character increases, the acidity of X–H increases (sp³ < sp² < sp).

Section 5.8     Resonance electron withdrawal or delocalization of electrons from a basic X atom into a conjugated π system decreases the basicity of X and increases the acidity of its corresponding X–H acid. Conversely, resonance electron donation of electrons from another basic X atom toward an sp² hybridized X atom of a +X–H acid increases the basicity of X and decreases the acidity of +X–H.

Section 5.9     Inductive electronic effects arise from induced dipoles in σ bonds connecting polarized substituents with acidic X–H bonds. Alkyl groups and negatively charged groups are electron donating, which increases the basicity of X and decreases the acidity of the corresponding X–H. Substituents with highly polar bonds to more electronegative atoms are electron withdrawing, which decreases the basicity of X and increases the acidity of its corresponding X–H acid. Inductive effects are highly distance dependent, that is, they become much weaker as the number of connecting σ bonds increases.

Section 5.10   Combined inductive and resonance electronic effects arise when substituents are conjugated with acidic functional groups via aromatic rings or other extended π systems. For substituents meta to the acidic group in aromatic rings, the inductive effects are dominant. For para- and ortho-substituents, the resonance effects are dominant. Resonance electron donors have lone pairs of electrons conjugated with the π system. Resonance electron withdrawers have polar π bonds (e.g., C=O or N=O) conjugated with the π system.

Section 5.11   The close proximity of hydrophobic substituents to acidic functional groups decreases the ability of water to solvate the ionic forms (anionic or cationic). Thus, hydrophobic substituents decrease the acidity of H–A acids and increase the acidity of BH⁺ acids. The close proximity of substituents that form hydrogen bonds with acidic groups can alter their acidity. Hydrogen bonds that stabilize the base increase the acidity, while hydrogen bonds that stabilize the acid decrease the acidity.

Section 5.12   The equilibrium between an acid and its conjugate base shifts as the pH of a solution changes by an amount defined by its acid dissociation constant (K_a). The concentrations of acid and base can be calculated using the Henderson–Hasselbalch equation: pK_a = pH – log([base]/[acid]). In all cases, [acid] > [base] when pH < pK_a and [base] > [acid] when pH > pK_a. Thus, qualitatively the dominant form of an H–A acid is uncharged at pH < pK_a and becomes charged (anionic) at pH > pK_a. Conversely, the dominant form of a BH⁺ acid is charged (cationic) at pH < pK_a and becomes uncharged at pH > pK_a.

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