Acids, Bases, and Buffers: Principles and Biological Applications

Acids and Bases: Definitions and Concepts

Arrhenius and Brønsted-Lowry Definitions

Acids and bases are fundamental chemical species with several definitions. Understanding these definitions is crucial for predicting their behavior in aqueous solutions.

• Arrhenius Acid: Dissolves in water to release H+ ions.

• Arrhenius Base: Dissolves in water to release OH- ions.

• Brønsted-Lowry Acid: A substance that donates a proton (H+).

• Brønsted-Lowry Base: A substance that accepts a proton (H+).

• Conjugate Acid-Base Pair: Two species that differ by a single proton (H+).

Example: In the reaction HCl + H2O → Cl- + H3O+, HCl is the acid (proton donor), H2O is the base (proton acceptor), Cl- is the conjugate base, and H3O+ is the conjugate acid.

Amphoteric Nature of Water

Water can act as both an acid and a base, depending on the reaction partner.

• As an acid: H2O + NH3 → OH- + NH4+

• As a base: H2O + HCl → H3O+ + Cl-

Example: Water donates a proton to ammonia, acting as an acid, and accepts a proton from HCl, acting as a base.

Charges on Conjugate Acids and Bases

Adding or removing a proton (H+) changes the charge of a species:

• Adding H+ increases the charge by +1.

• Removing H+ decreases the charge by -1.

Example: NH3 (neutral) + H+ → NH4+ (charge increases by +1).

Polyprotic Acids

Polyprotic acids can donate more than one proton per molecule. Each proton is lost in a stepwise fashion, with each step having its own dissociation constant.

• Example: H3PO4 (phosphoric acid) is triprotic.

Acid	Number of Protons Donated
H2SO4	2 (diprotic)
H3PO4	3 (triprotic)

Acid-Base Equilibria and pH

Self-Ionization of Water

Water can self-ionize to form hydronium and hydroxide ions:

$\mathrm{2H_2O (l) \rightleftharpoons H_3O^+ (aq) + OH^- (aq)}$

• The equilibrium constant for this reaction is $K_w$.

• $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C.

pH and pOH

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

• $\mathrm{pH = -\log[H_3O^+]}$

• $\mathrm{pOH = -\log[OH^-]}$

• $\mathrm{pH + pOH = 14}$ (at 25°C)

Example: If $[H_3O^+] = 1 \times 10^{-7}$ M, then pH = 7 (neutral solution).

Acidity and Basicity of Solutions

• Acidic: pH < 7

• Neutral: pH = 7

• Basic: pH > 7

Strong and Weak Acids and Bases

Strong Acids and Bases

Strong acids and bases dissociate completely in water.

• Strong acids: HCl, HBr, HI, HNO3, H2SO4

• Strong bases: Group IA and IIA metal hydroxides (e.g., NaOH, KOH, Ca(OH)2)

Example: HCl (aq) + H2O (l) → Cl- (aq) + H3O+ (aq)

Weak Acids and Bases

Weak acids and bases do not fully dissociate in water. Their strength is described by the acid dissociation constant ($K_a$) or base dissociation constant ($K_b$).

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• p$K_a = -\log K_a$

• Smaller p$K_a$ = stronger acid; larger p$K_a$ = weaker acid

Acid	p$K_a$	Strength
HCl	<0	Strong
Acetic acid	4.76	Weak
H2O	15.7	Very weak

Acid-Base Reactions and Neutralization

General Acid-Base Reactions

When acids react with bases, they form a salt and water. This is called a neutralization reaction.

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

• Acids react with metals to form hydrogen gas and a salt.

• Acids react with metal carbonates to form a salt, water, and carbon dioxide.

• Acids react with ammonia or amines to form ammonium salts.

Buffer Solutions

Definition and Function

A buffer is a combination of substances that resists changes in pH when small amounts of acid or base are added. Buffers are usually made from a weak acid and its conjugate base.

• Buffer capacity is greatest when pH = p$K_a$.

• Buffer range: p$K_a$ ± 1

Henderson-Hasselbalch Equation:

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

Example: For acetic acid/acetate buffer, if [CH3COO-] = [CH3COOH], then pH = p$K_a$.

Biological Buffer Systems

Phosphate Buffer System

The phosphate buffer system is important in intracellular fluids.

• $\mathrm{H_2PO_4^- \rightleftharpoons HPO_4^{2-} + H^+}$

• p$K_a$ = 7.2

$\mathrm{pH = pK_a + \log \left( \frac{[HPO_4^{2-}]}{[H_2PO_4^-]} \right)}$

Carbonic Acid-Bicarbonate Buffer System

This buffer system is crucial for maintaining blood pH.

• $\mathrm{CO_2 (g) + H_2O (l) \rightleftharpoons H_2CO_3 (aq) \rightleftharpoons HCO_3^- (aq) + H^+ (aq)}$

• p$K_a$ = 6.4

• Blood pH is maintained at ~7.4 by the ratio of bicarbonate to carbonic acid.

Protein Buffering

Proteins contain functional groups (e.g., amino, carboxyl) that can act as buffers by accepting or donating protons.

Blood Buffering and Homeostasis

Blood pH is regulated by multiple systems, including the lungs (CO2 exhalation), kidneys (excretion of H+ and reabsorption of HCO3-), and buffer systems.

• Red blood cells convert CO2 to HCO3- for transport.

• Kidneys excrete excess acid or base to maintain pH balance.

Summary Table: Major Buffer Systems in the Body

Buffer System	Main Components	Location
Carbonic acid-bicarbonate	H2CO3 / HCO3-	Blood plasma
Phosphate	H2PO4- / HPO42-	Intracellular fluid
Protein	Various amino acid side chains	Cells, plasma

Acid-Base Disorders

Acidosis and Alkalosis

Disruptions in acid-base balance can lead to acidosis (low blood pH) or alkalosis (high blood pH). These conditions can be caused by respiratory or metabolic changes.

• Acidosis: pH < 7.35

• Alkalosis: pH > 7.45

• Symptoms include changes in breathing, heart rate, and nervous system function.

Key Equations and Concepts

• $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$

• $\mathrm{pH = -\log[H_3O^+]}$

• $\mathrm{pOH = -\log[OH^-]}$

• $\mathrm{pH + pOH = 14}$

• $K_a = \frac{[H_3O^+][A^-]}{[HA]}$

• $\mathrm{pK_a = -\log K_a}$

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