Acidity and Basicity of Functional Groups & The Henderson-Hasselbalch Equation

Acidity and Basicity of Functional Groups

Introduction to Acidity and Basicity

Acidity and basicity are fundamental properties of organic molecules, determined by the ability of functional groups to donate or accept protons (H+). These properties influence molecular reactivity, solubility, and biological activity.

• Acids: Compounds that can donate a proton (H+); characterized by lower pKa values.

• Bases: Compounds that can accept a proton; characterized by higher pKa values (of their conjugate acids).

Strengths of Acids and Bases

The strength of an acid or base is quantitatively described by its pKa value.

• Strong acids: Lower pKa (typically < 1)

• Weak acids: Higher pKa (typically > 7)

• Strong bases: Conjugate acids with high pKa values

• Weak bases: Conjugate acids with low pKa values

pKa is the negative logarithm of the acid dissociation constant (Ka):

$pK_a = -\log K_a$

Common Acidic Functional Groups

Acidic functional groups are those that can lose a proton to form a conjugate base. Their typical pKa ranges are as follows:

Functional Group	pKa
Sulfonic acids, RSO3H	−1 to 1
Carboxylic acids, RCOOH	2 to 6
Phenols, ArOH	9 to 11
Thiols, RSH	9 to 11
Imides, RCONHCOR	9 to 10
Sulfonamides, RNHSO2R	9 to 10

Common Basic Functional Groups

Basic functional groups are those that can accept a proton to form a conjugate acid. Their typical pKa ranges are as follows:

Functional Group	pKa
Aliphatic amines	8 to 11
Saturated nitrogen heterocycles	9 to 11
Imidazoles, pyridines	4 to 6
Anilines	3 to 5

Note: Most weak bases are nitrogen-containing groups, but not all nitrogen-containing groups are basic.

Effect of Substitution (-R Groups)

Substituents on a molecule can influence acidity and basicity through electronic effects:

• Electron-withdrawing groups (e.g., nitro, halogen, carbonyl, aryl):

  • Increase acidity of HA (stabilize conjugate base)

  • Decrease basicity of B (increase acidity of BH+)

• Electron-donating groups (e.g., alkyl):

  • Decrease acidity of HA

  • Increase basicity of B (decrease acidity of BH+)

• Some substituents can stabilize ionized species via resonance.

Example: Benzocaine

Benzocaine contains an aromatic amine and an ester functional group.

• Is benzocaine an acid or base? It is a weak base (aromatic amine).

• Estimated pKa value: Aniline-type amines have pKa ≈ 5.

• Ionization at pH 7.4: Mostly un-ionized (since pH > pKa).

• Percent ionization at pH 7.4: Approximately 1% ionized.

• Water solubility: More water-soluble at pH 2 (where it is more ionized).

The Henderson-Hasselbalch Equation

Introduction

The Henderson-Hasselbalch equation relates the pH of a solution to the pKa and the ratio of concentrations of the ionized and un-ionized forms of an acid or base. It is essential for understanding buffer systems and drug ionization in biological systems.

Equations for Acids and Bases

• For acids (HA):

  • $HA \rightleftharpoons H^+ + A^-$

  • $K_a = \frac{[H^+][A^-]}{[HA]}$

  • $pH = pK_a + \log \frac{[A^-]}{[HA]}$

• For bases (B):

  • $BH^+ \rightleftharpoons B + H^+$

  • $K_a = \frac{[B][H^+]}{[BH^+]}$

  • $pH = pK_a + \log \frac{[B]}{[BH^+]}$

• General form: $pH = pK_a + \log \frac{[Base]}{[Acid]}$

Pharmaceutical Relevance

• pKa is an intrinsic property of a drug, buffer, or excipient.

• The equation helps predict the degree of ionization at a given pH, which affects solubility, absorption, and distribution of drugs.

pH Ionization of Acids

Calculating Ionization

For a weak acid (HA) with a given pKa, the ratio of ionized (A−) to un-ionized (HA) forms at different pH values can be calculated:

• $pH = pK_a + \log \frac{[A^-]}{[HA]}$

• $\frac{[A^-]}{[HA]} = 10^{(pH - pK_a)}$

Table: Ionization of a Weak Acid (pKa = 4)

pH	[A−]/[HA]	%A−	%HA
1	0.001	0.1	99.9
2	0.01	1	99
3	0.1	9	91
4	1	50	50
5	10	91	9
6	100	99	1
7	1,000	99.9	0.1

Interpretation: When pH < pKa, the un-ionized form (HA) dominates. When pH > pKa, the ionized form (A−) dominates.

Example Calculation

• At pH = pKa, $[A^-] = [HA]$ and both forms are present at 50% each.

• At pH one unit above pKa, $[A^-]/[HA] = 10$, so ~91% is ionized.

• At pH one unit below pKa, $[A^-]/[HA] = 0.1$, so ~9% is ionized.

Applications

• Understanding the degree of ionization is crucial for predicting drug absorption, solubility, and distribution in biological systems.

• Pharmacists use these calculations to optimize drug formulations and delivery.

Example: Aspirin (Acetylsalicylic Acid)

• pKa: ~3.5

• Stomach (pH 1.8): Un-ionized form dominates (better absorption).

• Small intestine (pH 6.8): Ionized form dominates (less absorption).

• Blood (pH 7.4): Ionized form dominates.

Summary

• Acidity and basicity are determined by functional groups and their substituents.

• pKa values provide a quantitative measure of acid/base strength.

• The Henderson-Hasselbalch equation allows calculation of ionization at different pH values, which is essential for understanding drug behavior in the body.