Acids, Bases, and Buffers in the Body: Study Guide

Acids, Bases, and Buffers in the Body

Acids: Definitions and Properties

Acids are fundamental chemical substances that play a crucial role in biological and chemical systems. Their definitions have evolved over time:

• Arrhenius Definition: Acids are substances that dissociate in water to produce hydrogen ions (H+).

• Brønsted-Lowry Definition: Acids are compounds that donate a proton (hydrogen ion) to another substance.

• Properties: Acids have a sour taste and can corrode metals.

• In Aqueous Solution: Protons associate with water molecules, forming hydronium ions (H3O+).

Example: Hydrochloric acid dissociates in water: $\mathrm{HCl(g) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)}$

Bases: Definitions and Properties

Bases are substances that counteract acids and are essential in many biological processes.

• Arrhenius Definition: Bases are ionic compounds that dissociate in water to form a metal ion and a hydroxide ion (OH-).

• Brønsted-Lowry Definition: Bases are proton acceptors.

• Properties: Bases have a bitter taste and a slippery feel.

• Common Bases: Most Arrhenius bases are formed from Group 1A and 2A metals.

Example: Sodium hydroxide dissociates in water: $\mathrm{NaOH \rightarrow Na^+ + OH^-}$

Acids and Bases in Aqueous Solutions

Water can act as both an acid and a base, donating or accepting a proton. This property is called amphoterism.

Strong vs. Weak Acids and Bases

Strong Acids

• Definition: Strong acids dissociate completely (100%) in water, producing hydronium ions and anions.

• Example: $\mathrm{HCl(g) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)}$

Weak Acids

• Definition: Weak acids dissociate only partially (about 5%) in water.

• Example: $\mathrm{CH_3COOH(l) + H_2O(l) \rightleftharpoons CH_3COO^-(aq) + H_3O^+(aq)}$

Strong Bases

• Definition: Strong bases dissociate completely in water to give a metal ion and hydroxide ion.

• Example: $\mathrm{B(OH) \rightarrow B^+ + OH^-}$

Weak Bases

• Definition: Weak bases dissociate only partially in water.

• Example: $\mathrm{B + H_2O \rightleftharpoons BH^+ + OH^-}$

Naming Acids

• Anion ending in -ide: Hydro[anion name minus ide]-ic acid (e.g., chloride → hydrochloric acid)

• Anion ending in -ate: [anion name minus ate]-ic acid (e.g., sulfate → sulfuric acid)

• Anion ending in -ite: [anion name minus ite]-ous acid (e.g., nitrite → nitrous acid)

Neutralization

• Mixing strong acids and strong bases results in the production of salt and water, neutralizing each other.

Example: $\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}$

Antacids

• Antacids are used to neutralize excess stomach acid (HCl).

• They are generally weak bases.

Example: $\mathrm{CaCO_3(s) + 2 HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)}$

Common Antacid	Active Ingredient
Tums	CaCO3
Rolaids	CaCO3, Mg(OH)2
Maalox	Al(OH)3, Mg(OH)2
Mylanta	Al(OH)3, Mg(OH)2
Additional info:	Other antacids may contain sodium bicarbonate (NaHCO3).

Chemical Equilibrium

Reversible Reactions and Equilibrium

Some chemical reactions are reversible, meaning products can revert to reactants. Over time, the rates of forward and reverse reactions become equal, establishing chemical equilibrium.

The Equilibrium Constant, K

• The equilibrium constant (K) expresses the ratio of product concentrations to reactant concentrations at equilibrium.

• For a generic reaction: $\mathrm{aA + bB \rightleftharpoons cC + dD}$

• Equilibrium constant expression: $\mathrm{K = \frac{[C]^c [D]^d}{[A]^a [B]^b}}$

Interpreting K

• K = 1: Equal amounts of products and reactants.

• K > 1: Products predominate.

• K < 1: Reactants predominate.

Le Châtelier’s Principle

• Applying stress (change in concentration, temperature, etc.) to an equilibrium causes the system to shift to offset the stress and restore equilibrium.

• If a substance is added, the reaction shifts away from that side; if removed, it shifts toward that side.

Weak Acids, Bases, and Equilibrium

Acid Dissociation Constant, Ka

• Weak acids and bases only partially dissociate, so equilibrium applies.

• The equilibrium constant for acid dissociation is called Ka.

• Ka values closer to 1 indicate stronger acids.

Acid	Ka (25°C)	pKa
Phosphoric acid (H3PO4)	7.5 × 10−3	2.12
Acetic acid (CH3COOH)	1.75 × 10−5	4.76
Bicarbonate ion (HCO3−)	4.8 × 10−11	10.32

Weak Acids and Oxygen Transport

• Weak acid-base equilibria modulate oxygen binding and release in hemoglobin.

• Protons (H+) and O2 bind to hemoglobin with opposite affinity; only one binds at a time.

• Shifts in H+ or O2 concentration affect the equilibrium of this reaction.

Autoionization of Water and pH

Autoionization of Water, Kw

• Water can act as a weak acid or base and autoionizes.

• Equilibrium constant for water: $\mathrm{K_w = [H_3O^+][OH^-]}$

• In pure water: $\mathrm{[H_3O^+] = [OH^-]}$

pH and the pH Scale

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

• Calculation: $\mathrm{pH = -\log[H^+]}$ $\mathrm{pH = -\log[H_3O^+]}$

• To find [H3O+] from pH: $\mathrm{[H_3O^+] = 10^{-pH}}$

• Acidic solutions: pH < 7, [H3O+] > 1.0 × 10−7 M

• Neutral solutions: pH = 7, [H3O+] = 1.0 × 10−7 M

• Basic solutions: pH > 7, [H3O+] < 1.0 × 10−7 M

Normal blood pH: 7.35–7.45 (strictly regulated).

Measurement: pH meter (electronic) or pH paper (color indicators).

pKa and Acid Strength

• pKa is the negative logarithm of Ka: $\mathrm{pK_a = -\log K_a}$

• The smaller the pKa, the stronger the acid.

Drug Solubility, pH, and pKa

Drug Ionization and Solubility

• Drug molecules can be charged or uncharged depending on pH and their pKa.

• Example: Propranolol (pKa = 9.42) is charged in acidic environments (stomach, pH 2) and uncharged in basic environments (intestine, pH 8).

• Charged form is soluble in stomach; uncharged form diffuses across intestinal lining.

• If pH < pKa, more acid form is present; if pH > pKa, more conjugate base is present.

The Henderson-Hasselbalch Equation

• Relates pH, pKa, and the ratio of conjugate base to acid.

• Equation: $\mathrm{pH = pK_a + \log \left(\frac{[A^-]}{[HA]}\right)}$

Tissue pH and Drug Delivery

• Local anesthetics contain an amine group; the acid form is charged, the conjugate base is uncharged.

• Uncharged form crosses nerve-cell membrane; charged form relieves pain.

The Bicarbonate Buffer System

Buffers and pH Regulation

• A buffer is a solution containing a weak acid and its conjugate base (or weak base and its conjugate acid).

• Buffers resist changes in pH.

• Blood is buffered by the bicarbonate system, maintaining homeostasis.

• At equilibrium, pH = pKa.

Homeostatic Imbalance

• Imbalances in pH can lead to health issues.

Respiratory Imbalance

• Hypoventilation: CO2 builds up, blood becomes acidic (respiratory acidosis). Treatment: IV bicarbonate.

• Hyperventilation: Excess CO2 is exhaled, blood becomes basic (respiratory alkalosis). Treatment: Breathe into a paper bag.

Regulating CO2 and Bicarbonate in the Body

• The body regulates CO2 and bicarbonate to maintain blood pH within a narrow range.

Additional info: The study notes include inferred details about common antacids and the Henderson-Hasselbalch equation for completeness.