Acid-Base Equilibrium

Brønsted–Lowry Acids and Bases

The Brønsted–Lowry theory is the most widely used definition for acids and bases in chemistry. It focuses on the transfer of protons (H+) between chemical species.

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

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

• Acid-base reaction: Involves the transfer of a single proton from the acid to the base.

Example: The reaction of hydrochloric acid (HCl) with water:

• HCl (acid) donates a proton to H2O (base), forming Cl– and H3O+.

Equation:

Conjugate Acid-Base Pairs

In every Brønsted–Lowry acid-base reaction, there are two pairs of species that differ by one proton. These are called conjugate acid-base pairs.

• Conjugate acid: The species formed when a base gains a proton.

• Conjugate base: The species formed when an acid loses a proton.

Example: For the reaction above:

• HCl and Cl– are a conjugate acid-base pair.

• H2O and H3O+ are a conjugate acid-base pair.

The Concept of pH

pH is a logarithmic scale used to express the concentration of hydronium ions ([H3O+]) in a solution, making it easier to handle very small numbers.

• pH:

• pOH:

• At 25°C, pure water has pH = 7 and pOH = 7.

• At 25°C, .

• Acidic solutions: pH < 7; Basic solutions: pH > 7.

Strength of Acids and Bases

The strength of an acid or base refers to its degree of ionization in water.

• Strong acid: Completely ionizes in water (reaction goes to completion).

• Weak acid: Partially ionizes in water (establishes equilibrium).

• Strong base: Completely reacts with water to form OH–.

• Weak base: Partially reacts with water to form OH–.

• Inverse relationship: The stronger the acid, the weaker its conjugate base, and vice versa.

Weak Acids and Bases: Equilibrium and Constants

Weak acids and bases do not fully react with water, resulting in an equilibrium between reactants and products. The extent of this reaction is quantified by equilibrium constants.

• For a weak acid:

• For a weak base:

• Acidity constant (Ka):

• Basicity constant (Kb):

• Typically, Ka and Kb are much less than 1 for weak acids and bases.

To simplify calculations, logarithmic forms are used:

• at 25°C

• at 25°C

Strength relationships:

• Larger Ka (smaller pKa) means a stronger acid.

• Larger Kb (smaller pKb) means a stronger base.

pH Calculations for Weak Acids and Bases

To determine the pH of a solution containing a weak acid or base, use the equilibrium constant and an ICE (Initial, Change, Equilibrium) table.

Example: Calculate the pH of a 1.0 M acetic acid (CH3COOH) solution.

• Set up the equilibrium:

• Assume x is small compared to 1.0, so 1.0 – x ≈ 1.0.

• Solve for x using Ka, then calculate pH: .

Note: The 'p' in pH, pKa, etc., means '–log' of the variable.

Buffer Solutions

A buffer solution contains significant amounts of both a weak acid and its conjugate base (or a weak base and its conjugate acid). Buffers resist changes in pH when small amounts of acid or base are added, or upon moderate dilution.

• The effective pH range of a buffer is determined by the pKa (or pKb) and the ratio of conjugate pairs.

• Buffers are most effective when the ratio [A–]/[HA] is close to 1.

• Buffer capacity is best when pH ≈ pKa (or pOH ≈ pKb).

pH Calculations in Buffer Solutions: The Henderson–Hasselbalch Equation

The Henderson–Hasselbalch equation relates the pH of a buffer to the concentration ratio of the conjugate base and acid:

• When [A–] = [HA], pH = pKa.

• Buffers are effective when the ratio [A–]/[HA] is between 1/10 and 10/1, i.e., pH = pKa ± 1.

Example: Calculate the pH of a buffer solution with [NaOOCCH3] = 0.11 M and [CH3COOH] = 0.090 M. Given pKa for acetic acid is 4.74:

Summary Table: Key Acid-Base Equilibrium Relationships

Additional info: For more details on specific weak acids and bases, refer to tables in your textbook (e.g., Tables 11.6 and 11.7). The above notes provide a comprehensive overview of acid-base equilibrium, weak acids and bases, and buffer solutions, suitable for exam preparation in General Chemistry.