Acid-Base Chemistry, pH, and Buffer Systems in Physiology and Pharmacy

Henderson-Hasselbalch Equation and Acid-Base Chemistry

Introduction

This study guide covers the fundamental principles of acid-base chemistry, the calculation and significance of pH, and the role of buffer systems in physiological and pharmaceutical contexts. Understanding these concepts is essential for interpreting drug stability, absorption, and compatibility in biological systems.

Objectives

• Calculate pH with and without changing volumes (without being given the equation).

• List and explain factors that can cause a pH drift.

• State or estimate the pH ranges for the stomach, intestines, and blood.

• Identify common buffer systems and classify them as acidic or basic.

Acids and Bases

Arrhenius Classification

The Arrhenius definition is a classical approach to acids and bases in aqueous solutions:

• Acid: A substance that yields hydrogen ions (H+) in aqueous solution. $\mathrm{HA} \rightarrow \mathrm{H}^+ + \mathrm{A}^-$

• Base: A substance that yields hydroxyl ions (OH-) in aqueous solution. $\mathrm{BOH} \rightarrow \mathrm{B}^+ + \mathrm{OH}^-$

Limitation: The Arrhenius definition does not account for acid-base reactions outside of water.

Bronsted-Lowry and Lewis Theories

• Bronsted-Lowry Acid: Proton (hydrogen ion) donor.

• Bronsted-Lowry Base: Proton (hydrogen ion) acceptor.

• Lewis Acid: Electron pair acceptor.

• Lewis Base: Electron pair donor.

Amphoteric substances can act as either acids or bases. Water is a classic example.

Weak Acids and Bases

• Weak acids and weak bases dissociate less than 50% in water and are called weak electrolytes.

• Most pharmaceutical agents are weak acids or weak bases.

• Drugs are often formulated as salts to improve solubility and stability.

pH and Its Importance

Definition and Scale

• pH is a measure of hydrogen ion concentration in solution.

• The pH scale ranges from 0 (strong acid) to 14 (strong base), with 7 being neutral.

• At 25°C, neutral pH is 7, where $[\mathrm{H}^+] = [\mathrm{OH}^-]$.

Physiological pH Ranges

• Blood: pH ≈ 7.4

• Stomach: pH 1.5 – 4.0 (varies with digestion and disease)

• Intestine: pH ≈ 6.45 (increases along the tract)

Significance of pH in Pharmacy and Physiology

• pH affects drug stability, solubility, and efficacy.

• Enzyme and transporter function is sensitive to pH changes.

• Drug dissolution and absorption are pH-dependent.

• Each route of administration (ROA) has an acceptable pH range (e.g., ophthalmic: 4.5–11.5, ideally 6.5–8.5).

Buffer Systems

Definition and Function

• Buffers are substances that resist changes in pH upon addition of small amounts of acid or base.

• Buffer efficiency is measured by buffering capacity.

• A buffered solution maintains a relatively constant pH.

Common Buffer Systems in Pharmacy

Buffer System	pH Range	Notes
Sodium hydroxide	11–13	Strong base, used diluted (0.1–0.2 N)
Hydrochloric acid	1–3	Strong acid, used diluted (0.1–0.2 N)
Sodium citrate–citric acid	Varies	Used for penicillin, cefuroxime
Sodium borate–boric acid	Varies	Not for injection
Phosphate buffers	Varies	Monobasic sodium/potassium phosphate
Sodium acetate–acetic acid	Varies	Commonly used

Factors Causing pH Drift

• Drug, adjuvant, or excipient instability

• Interactions with containers/closures (e.g., glass surfaces can increase pH)

• Atmospheric CO2 absorption (acidic shift)

• Addition of excipients, adjuvants, or pH adjusters

• Buffering capacity of the solvent

Acid and Base Dissociation Constants

Definitions

• Ka: Acid dissociation constant; measures acid strength.

• Kb: Base dissociation constant; measures base strength.

• pKa: $\mathrm{p}K_a = -\log K_a$

• pKb: $\mathrm{p}K_b = -\log K_b$

• Kw: Ionic product of water; $K_w = 1.0 \times 10^{-14}$ at 25°C

• pKw: $\mathrm{p}K_w = 14$

Relationship Between pKa and pKb

• $\mathrm{p}K_a + \mathrm{p}K_b = \mathrm{p}K_w$

• For water at 25°C: $\mathrm{p}K_w = 14$

Calculating pH Without the Henderson-Hasselbalch Equation

General Approach

• For weak acids/bases, use the ratio of salt to acid/base and the pKa or pKb value.

• When volumes and concentrations are given, calculate the number of moles (mmol) of acid and salt.

• Use the formula: $\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right)$ Additional info: This is the Henderson-Hasselbalch equation, but students are expected to know how to apply the concept without being given the formula.

Example Calculation

• Given: 25 mL of 0.2 M benzoic acid and 45 mL of 0.1 M sodium benzoate, pKa = 4.22

• Calculate mmol: 25 mL × 0.2 mmol/mL = 5 mmol benzoic acid 45 mL × 0.1 mmol/mL = 4.5 mmol sodium benzoate

• Apply the ratio: $\mathrm{pH} = 4.22 + \log \left( \frac{4.5}{5} \right)$

Clinical and Pharmaceutical Applications

Drug Formulation and Stability

• Pharmacists must determine the pH of maximum stability before preparing solutions.

• Maximal stability is often near neutral pH, but not always.

• Temperature, pH, and the nature of the medium influence drug stability.

Metabolic Disorders

• Metabolic acidosis: Excess acid in the body.

• Metabolic alkalosis: Increased bicarbonate or base.

• Polyacidic and polybasic compounds (e.g., citric acid) have multiple pKa values.

Summary Table: Physiological pH Ranges

Body Fluid/Region	Typical pH Range	Notes
Blood	7.35–7.45	Maintained by respiratory and renal systems
Stomach	1.5–4.0	Varies with digestion, disease, and medication
Intestine	~6.45	Increases along the tract

Key Equations

• $\mathrm{p}K_a = -\log K_a$

• $\mathrm{p}K_b = -\log K_b$

• $\mathrm{p}K_a + \mathrm{p}K_b = \mathrm{p}K_w$

• $\mathrm{p}K_w = 14$ (at 25°C)

• $\mathrm{pH} = \mathrm{p}K_a + \log \left( \frac{[\text{salt}]}{[\text{acid}]} \right)$

Conclusion

Understanding acid-base chemistry, pH calculation, and buffer systems is essential for both physiology and pharmaceutical sciences. These concepts underpin drug formulation, stability, and absorption, as well as the maintenance of homeostasis in the human body.