Acids and Bases

Properties of Acids and Bases

Acids and bases are fundamental chemical species with distinct properties that are easily observed in laboratory and everyday contexts.

• Acids: Sour taste, ability to dissolve metals, turn litmus paper red, neutralize bases.

• Bases: Bitter taste, slippery feel, turn litmus paper blue, neutralize acids.

Definitions of Acids and Bases

There are several definitions for acids and bases, each with its own scope and application:

• Arrhenius Definition: An acid produces H+ ions in aqueous solution; a base produces OH– ions in aqueous solution.

• Brønsted–Lowry Definition: An acid donates H+ in a chemical reaction; a base accepts H+. This definition is more broadly applicable.

• Lewis Definition: (To be covered later) Focuses on electron pair donation and acceptance.

Acid Dissociation and Conjugate Pairs

Acids dissociate in water, forming conjugate acid-base pairs. For example:

• CH3COOH(aq) + H2O(l) → CH3COO–(aq) + H3O+(aq)

• Acid (donates H+): CH3COOH

• Base (accepts H+): H2O

• Conjugate base: CH3COO–

• Conjugate acid: H3O+

The weaker the acid, the stronger its conjugate base.

Acid and Base Strength

Acid Ionization Constant (Ka)

The equilibrium constant for acid dissociation is called the acid ionization constant (Ka):

• General reaction: HA(aq) + H2O(l) → A–(aq) + H3O+(aq)

• Ka expression:

• The stronger the acid, the larger the Ka value, and the equilibrium favors products.

Strong vs. Weak Acids

• Strong acids: Dissociate completely in water; equilibrium lies far to the right. Example: HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)

• Weak acids: Partially dissociate; equilibrium exists and Ka is used to quantify strength.

Base Ionization Constant (Kb)

For bases, the equilibrium constant is called Kb:

• General reaction: B(aq) + H2O(l) → BH+(aq) + OH–(aq)

• Kb expression:

• The larger the Kb, the stronger the base.

Relationship Between Ka and Kb

For a conjugate acid-base pair:

• At 25°C,

• If Ka is known, Kb can be calculated for the conjugate base, and vice versa.

Autoionization of Water and the Ion-Product Constant (Kw)

Autoionization Reaction

Water can act as both an acid and a base (amphoteric). The autoionization reaction is:

• 2H2O(l) → H3O+(aq) + OH–(aq)

• Equilibrium constant:

• At 25°C,

Acidic, Basic, and Neutral Solutions

• If [H3O+] > [OH–], the solution is acidic.

• If [H3O+] < [OH–], the solution is basic.

• If [H3O+] = [OH–], the solution is neutral.

Concentrations in Pure Water

• In pure water at 25°C: M

The pH and pOH Scales

Definition and Calculation

The pH scale is a logarithmic scale used to quantify acidity:

• At 25°C: pH < 7.0 is acidic, pH = 7.0 is neutral, pH > 7.0 is basic.

pOH Scale

pOH is defined similarly:

Relationship Between pH and pOH

• At 25°C:

• General relationship:

pKa and pKb

Percent Ionization of Weak Acids

Definition and Calculation

Percent ionization quantifies the fraction of a weak acid that dissociates in water:

• Example: For Ka = , [HA]in = 0.600 M, [H3O+]eq = M:

Effect of Concentration

• As the concentration of a weak acid increases, percent ionization typically decreases.

Calculating pH from Ka

Procedure

To find the pH of a weak acid solution:

1. Write the balanced dissociation reaction.

2. Write the Ka expression.

3. Create an equilibrium table (ICE table).

4. Solve for [H3O+] using the Ka expression.

5. Check approximation validity or use the quadratic formula if necessary.

6. Calculate pH:

Example: For HNO2 with Ka = and [HNO2]initial = 0.200 M, pH = 2.02.

Summary Table: Acid/Base Strength and Equilibrium Constants

Additional info: The notes above expand on the original lecture slides, providing full academic context, definitions, and stepwise procedures for calculations. All equations are formatted in LaTeX as required.