Acid-Base Theory Lecture Flashcards

Acid-Base Theory and Homeostasis

• Acid-base theory describes the mechanisms by which the body maintains a stable internal environment, known as homeostasis.
• Maintenance of homeostasis is essential for survival; the body must maintain the pH of its various fluids within a narrow range to ensure cells and organs function properly.
• Normal blood pH is strictly maintained between $7.35$ and $7.45$.
• The typical average blood pH is $7.4$.

Properties and Classification of Drugs

• The majority of drugs currently on the market are small organic molecules.
• In aqueous solutions, these molecules behave as either weak acids or weak bases.
• Market distribution of drugs based on acid-base properties:
  • Weak Bases: 60-70%
  • Weak Acids: 15-25%
• A vital principle in pharmaceutical science is that the $pK_a$ value alone does not determine whether a molecule is an acid or a base.

Brønsted-Lowry Definitions and Dissociation

• Acids: Substances capable of donating a proton ($H^+$).
• Bases: Substances capable of accepting a proton ($H^+$).
• Strong Acids and Bases: These dissociate completely in water and exist entirely in their ionized (charged) form. They are classified as strong electrolytes.
  • Examples of complete dissociation:
    • $HCl \rightarrow H^+ + Cl^-$
    • $H_2SO_4 \rightarrow 2H^+ + SO_4^{2-}$
    • $NaOH \rightarrow Na^+ + OH^-$
• Weak Acids and Bases: These only partially dissociate when added to water, establishing a chemical equilibrium between the ionized and un-ionized species.

Acid-Base Equilibrium and the Dissociation Constant

• For a weak acid ($HA$), the dissociation reaction in water is:
  • $HA \rightleftharpoons H^+ + A^-$
• The acid dissociation constant ($K_a$) is defined as:
  • $K_a = \frac{[H^+][A^-]}{[HA]}$
• The magnitude of $K_a$ indicates acid strength:
  • A large $K_a$ means the acid readily dissociates into charged species.
  • The larger the $K_a$, the stronger the acid.
• For a weak base ($B$), the reaction in water reacts with a proton to form a conjugate acid ($BH^+$):
  • $B + H^+ \rightleftharpoons BH^+$
• Because the species $BH^+$ is the form that can donate a proton, it is termed the conjugate acid. The equilibrium is rewritten as:
  • $BH^+ \rightleftharpoons B + H^+$
  • $K_a = \frac{[B][H^+]}{[BH^+]}$
• The relationship for bases is: the larger the $K_a$, the more the $BH^+$ dissociates to donate protons, meaning the larger the $K_a$, the weaker the base.

The Nature of Water and Conjugate Pairs

• Water is amphoteric, meaning it can function as either an acid or a base.
• Water as a proton donor (Acid): $H_2O \rightleftharpoons H^+ + OH^-$.
• Water as a proton acceptor (Base): $HA + H_2O \rightleftharpoons H_3O^+ + A^-$. The species $H_3O^+$ is the hydronium ion.
• A conjugate acid-base pair consists of an acid and a base that differ only by a single proton ($H^+$).
  • Reaction: $HA + H_2O \rightleftharpoons H_3O^+ + A^-$
  • Pairs: $HA$ (Acid) / $A^-$ (Conjugate Base); $H_2O$ (Base) / $H_3O^+$ (Conjugate Acid).
• In a lipid bilayer, the distribution of species depends on the un-ionized form ($HA$) vs. the ionized form ($A^-$).

Functional Group Classification

• Weak Acid Functional Groups:
  • Carboxylic Acids ($R-COOH$)
  • Phenols ($ArOH$)
  • Sulfonic Acids ($RSO_3H$)
  • Sulfonamides ($ArSO_2NHR$)
  • Thiols ($RSH$)
• Weak Base Functional Groups (Amines):
  • Aliphatic Amines (Primary, Secondary, Tertiary):
    • Primary ($1^\circ$): Nitrogen bound to 1 carbon ($R-NH_2$).
    • Secondary ($2^\circ$): Nitrogen bound to 2 carbons ($R_2-NH$).
    • Tertiary ($3^\circ$): Nitrogen bound to 3 carbons ($R_3-N$).
  • Aromatic Amines: The nitrogen is attached directly to an aromatic ring (e.g., Anilines).
  • Nitrogen-containing heterocycles (e.g., Pyridines, Saturated nitrogen heterocycles).
• Neutral Functional Groups (Neither acidic nor basic):
  • Alcohols ($R-OH$)
  • Ketones ($R^1-C(=O)-R^2$)
  • Aldehydes ($R-CHO$)
  • Ethers ($R-O-R'$)
  • Esters ($R-COOR'$)
  • Amides ($R-CONH_2$)
• Quaternary Ammonium Ion: A nitrogen with four substituents ($R_4N^+$). It is a fixed electrolyte that cannot accept or donate a proton; therefore, it is neither acidic nor basic.

Mathematical Calculations of pH and pKa

• pH Definitions:
  • $pH = -\log([H^+])$
  • $pOH = -\log([OH^-])$
  • $pH + pOH = 14$
• Relationship to concentration:
  • As hydrogen ion concentration increases, pH decreases.
  • When hydroxide concentration increases, it reacts with $H^+$ ions, raising the pH.
• pKa Definition:
  • $pK_a = -\log(K_a)$
• Practice Problem 1: pH of a solution with $[H^+] = 4.6 \times 10^{-3}\,M$
  • $pH = -\log(4.6 \times 10^{-3}) = 2.33$
• Practice Problem 2: Hydrogen ion concentration for $pH = 6.8$
  • $6.8 = -\log([H^+])$
  • $[H^+] = 10^{-6.8} = 1.58 \times 10^{-7}\,M$
• Practice Problem 3: pH of a solution with hydroxyl concentration $[OH^-] = 3.4 \times 10^{-11}\,M$
  • $pOH = -\log(3.4 \times 10^{-11}) = 10.47$
  • $pH = 14 - 10.47 = 3.53$
• Practice Problem 4: $pK_a$ of acetic acid with $K_a = 1.74 \times 10^{-5}$
  • $pK_a = -\log(1.74 \times 10^{-5}) = 4.76$
• Practice Problem 5: $K_a$ of propranolol with $pK_a = 9.5$
  • $9.5 = -\log(K_a)$
  • $K_a = 10^{-9.5} = 3.16 \times 10^{-10}$

pH Spectrum of Body Fluids

• Saliva: $6.0-6.4$
• Blood: $7.35-7.45$
• Gastric (GI) Juice: $1.0-3.5$
• Colon: $7.8-8.0$
• Duodenum: $5.0-6.0$
• Rectum: $7.8$
• Ileum: $8.0$
• Vagina: $3.4-4.2$
• Large Intestine: $8.0$
• Sweat: $4.0-6.8$
• Cerebrospinal Fluid (CSF): $7.3$
• Breast Milk: $6.6-7.0$
• Urine: $4.0-8.0$
• Semen: $7.2$
• Tears: $7.4$

Drug Ionization and the Henderson-Hasselbalch Relationship

• Formula: $pH = pK_a + \log\left(\frac{[base]}{[acid]}\right)$
• Weak Acids ($HA$):
  • At $pH = pK_a$, the drug is 50% ionized.
  • When $pH \lt pK_a$, the un-ionized form ($HA$) predominates.
  • When $pH \gt pK_a$, the ionized form ($A^-$) predominates.
• Weak Bases ($BH^+$):
  • At $pH = pK_a$, the drug is 50% ionized.
  • When $pH \lt pK_a$, the ionized form ($BH^+$) predominates.
  • When $pH \gt pK_a$, the un-ionized form ($B$) predominates.
• Example - Percent Nonionized Comparison:
  • Benzoic Acid ($pK_a = 4.0$): At pH 1 (99.9% nonionized), pH 4 (50% nonionized), pH 7 (0.1% nonionized).
  • Aniline ($pK_a = 5.0$): At pH 2 (0.1% nonionized), pH 5 (50% nonionized), pH 7 (99% nonionized).

Ion Trapping and Clinical Implications

• Mechanism: Most drugs in a lipid-soluble (un-ionized) form are reabsorbed by passive diffusion in the kidney. Charged (ionized) forms are trapped and excreted.
• Increasing Excretion:
  • Weak Acids: Excreted faster in alkaline pH (anion form favored). Sodium Bicarbonate is used to increase urinary pH.
  • Weak Bases: Excreted faster in acidic pH (cation form favored). Ammonium Chloride is used to decrease urinary pH.
• Clinical Modulation Matrix:
  • To clear Weak Bases (Amphetamine, Chloroquine, Imipramine, Levophanol, Mecamylamine, Quinine): Acidify the urine.
  • To clear Weak Acids (Acetazolamide, Nitrofurantoin, Phenobarbital, Probenecid, Salicylates/Aspirin, Sulfathiazole): Alkalinize the urine.
• Case Study (Aspirin Overdose): A 38-year-old female with high salicylate levels. Salicylate/Aspirin is a weak acid. Treatment involves alkalinizing the urine (often with sodium bicarbonate) to trap the salicylate in its ionized form, preventing reabsorption and facilitating clearing.

Salt Formation in Pharmaceuticals

• Salts are formed to alter the physical properties of drugs, primarily to increase water solubility.
• Weak Acid Salts: Formed by reacting a weak un-ionized acid with a strong base.
  • $RCOOH + NaOH \rightarrow RCOO^-Na^+ + H_2O$
  • Common additives: Sodium Hydroxide ($NaOH$), Potassium Hydroxide, Ammonium Hydroxide.
• Weak Base Salts: Formed by reacting a weak un-ionized base with a strong acid.
  • $R_3N + HCl \rightarrow R_3NH^+Cl^-$
  • Common additives: Hydrochloric Acid ($HCl$), Tartaric Acid, Succinic Acid, Maleic Acid, Sulfuric Acid, Nitric Acid, Citric Acid.
• Clinical Identifiers:
  • Bupropion Hydrochloride: A weak base (indicated by the use of hydrochloric acid to form a salt).
  • Zolpidem Tartrate: A weak base (indicated by the use of tartaric acid to form a salt).
  • Depakote Sodium (Divalproex sodium): A weak acid (indicated by the use of sodium to form the salt from valproic acid).