Acids, Bases, and Aqueous Equilibria: Brønsted–Lowry Theory, pH, and Buffer Solutions

Acids and Bases: Brønsted–Lowry Theory and Aqueous Equilibria

Introduction to Acids and Bases

Acids and bases are fundamental chemical species that play a crucial role in both natural and laboratory environments. Their behavior in aqueous solutions is central to understanding chemical reactivity, environmental chemistry, and biological processes.

• Acid: A substance that donates a proton (H+) to another substance.

• Base: A substance that accepts a proton (H+) from another substance.

• Brønsted–Lowry Theory: Defines acids as proton donors and bases as proton acceptors, expanding the concept beyond aqueous solutions.

Brønsted–Lowry Acid–Base Reactions

Acid–base reactions involve the transfer of protons between reactants. These reactions can be represented as:

• General Equation:

$\text{Acid} + \text{Base} \rightleftharpoons \text{Conjugate Base} + \text{Conjugate Acid}$

• Example: The reaction between hydrochloric acid and water:

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

• Conjugate Acid–Base Pairs: Each acid has a conjugate base (formed by loss of H+), and each base has a conjugate acid (formed by gain of H+).

Common Acids and Bases

Acids and bases are found in both laboratory and natural settings. The following tables summarize common examples:

Formula	Name	Structure
HCl	Hydrochloric acid	Cl–H
H2SO4	Sulfuric acid	O=S(=O)(O)O
NaOH	Sodium hydroxide	Na–O–H
NH3	Ammonia	H–N(H)–H

Additional info: See textbook for more examples and structures.

Acid–Base Reactions in Water

When acids and bases dissolve in water, they react to form hydronium (H3O+) and hydroxide (OH–) ions:

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

• This process is called the autoprotolysis of water.

• The equilibrium constant for this reaction is the ion-product constant ($K_w$):

$K_w = [\mathrm{H_3O^+}][\mathrm{OH^-}]$

• At 25°C, $K_w = 1.0 \times 10^{-14}$

The Concept of pH

The pH scale is a logarithmic measure of the concentration of hydronium ions in solution:

$pH = -\log_{10}[\mathrm{H_3O^+}]$

• pOH: $pOH = -\log_{10}[\mathrm{OH^-}]$

• Relationship: $pH + pOH = 14$ at 25°C

• Acidic solution: pH < 7; Neutral solution: pH = 7; Basic solution: pH > 7

Strength of Acids and Bases

Acids and bases are classified as strong or weak based on their degree of ionization in water:

• Strong acids/bases: Completely ionize in solution (e.g., HCl, NaOH).

• Weak acids/bases: Partially ionize in solution (e.g., CH3COOH, NH3).

The equilibrium constant for a weak acid is called the acid dissociation constant ($K_a$):

$K_a = \frac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$

• pKa: $pK_a = -\log_{10} K_a$

Calculating pH in Solutions

For strong acids and bases, pH can be calculated directly from concentration. For weak acids and bases, equilibrium calculations are required, often using the ICE table (Initial, Change, Equilibrium).

• Example (Strong Acid): 0.01 M HCl solution:

$[\mathrm{H_3O^+}] = 0.01$ M, so $pH = -\log_{10}(0.01) = 2$

• Example (Weak Acid): 0.10 M acetic acid ($K_a = 1.8 \times 10^{-5}$):

$CH_3COOH + H_2O \rightleftharpoons CH_3COO^- + H_3O^+$

Set up ICE table and solve for $[H_3O^+]$ using $K_a$.

Buffer Solutions

Buffer solutions are mixtures of a weak acid and its conjugate base (or a weak base and its conjugate acid) that resist changes in pH upon addition of small amounts of acid or base.

• The pH of a buffer can be calculated using the Henderson–Hasselbalch equation:

$pH = pK_a + \log_{10}\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$

• Example: A buffer made from acetic acid and sodium acetate.

The Molecular Basis of Acid Strength

The strength of an acid depends on the stability of its conjugate base and the molecular structure:

• Binary acids: Acid strength increases down a group and across a period (e.g., HI > HBr > HCl > HF).

• Oxoacids: Acid strength increases with more electronegative atoms and more oxygen atoms.

• Inductive effects: Electronegative atoms near the acidic proton increase acid strength by stabilizing the conjugate base.

Tables: Common Acids and Bases

Some common acids and bases used in the laboratory and found in nature are summarized below:

Formula	Name	Source/Use
HCl	Hydrochloric acid	Stomach acid, laboratory reagent
H2SO4	Sulfuric acid	Battery acid, acid rain
CH3COOH	Acetic acid	Vinegar
NaOH	Sodium hydroxide	Drain cleaner
NH3	Ammonia	Cleaning agent, biological waste

Summary

• Acids and bases are defined by their ability to donate or accept protons (Brønsted–Lowry theory).

• pH is a measure of hydronium ion concentration calculated using $pH = -\log_{10}[H_3O^+]$.

• Strong acids/bases fully ionize; weak acids/bases partially ionize and require equilibrium calculations.

• Buffer solutions resist changes in pH and are calculated using the Henderson–Hasselbalch equation.

• The molecular structure determines acid strength, influenced by periodic trends and inductive effects.

Additional info: For more detailed worked examples and advanced calculations, refer to the full textbook chapter.