Acids and Bases: Definitions and Theories

Arrhenius, Brønsted-Lowry, and Lewis Acids and Bases

Understanding acids and bases is fundamental in chemistry. Three main theories define acids and bases, each broadening the scope of what substances can be considered acids or bases.

• Arrhenius Theory:

  • Acid: Produces H+ ions in aqueous solution.

  • Base: Produces OH- ions in aqueous solution.

  • Example: HCl (acid) dissociates in water to give H+ and Cl-; NaOH (base) dissociates to give Na+ and OH-.

• Brønsted-Lowry Theory:

  • Acid: Proton (H+) donor.

  • Base: Proton (H+) acceptor.

  • Example: NH3 + H2O → NH4+ + OH-. Here, NH3 is the base (accepts H+), H2O is the acid (donates H+).

• Lewis Theory:

  • Acid: Electron pair acceptor.

  • Base: Electron pair donor.

  • Example: BF3 (acid) + NH3 (base) → F3B:NH3.

Additional info: The Lewis definition is the most general, encompassing reactions that do not involve protons.

Acid and Base Strength

Predicting and Quantifying Acid/Base Strength (Ka, Kb)

The strength of an acid or base is determined by its tendency to donate or accept protons, quantified by equilibrium constants.

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

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

• Acid Dissociation Constant (Ka): Measures the strength of a weak acid.

• Base Dissociation Constant (Kb): Measures the strength of a weak base.

• Relationship between Ka and Kb: where is the ion-product constant for water.

Example: Acetic acid (CH3COOH) has a Ka of 1.8 × 10-5, indicating it is a weak acid.

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

Concept and Calculations

Water can act as both an acid and a base, undergoing autoionization:

• Reaction:

• Ion-Product Constant for Water: At 25°C, .

• Calculations: If [H3O+] is known, [OH-] can be found using and vice versa.

Example: If [H3O+] = 1.0 × 10-5 M, then [OH-] = M.

The pH Scale and Solution Acidity

Quantifying Acidity and Basicity

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

• pH Definition:

• pOH Definition:

• Relationship: (at 25°C)

• Acidic Solution: pH < 7

• Neutral Solution: pH = 7

• Basic Solution: pH > 7

Example: If [H3O+] = 1.0 × 10-3 M, then pH = 3 (acidic).

Percent Ionization of Weak Acids

Definition and Calculation

Percent ionization indicates the fraction of acid molecules that ionize in solution.

• Formula:

• Application: Used to compare the strengths of weak acids at different concentrations.

Example: If 0.10 M acetic acid produces 1.3 × 10-3 M H3O+, percent ionization = .

Calculating pH of Acid and Base Solutions

Strong and Weak Acids/Bases, Mixtures

• Strong Acids/Bases: Assume complete dissociation.

  • pH of Strong Acid: (where [H3O+] = initial acid concentration)

  • pOH of Strong Base: (where [OH-] = initial base concentration)

• Weak Acids/Bases: Use ICE tables and Ka or Kb to solve for equilibrium concentrations.

  • Example: For HA, set up and solve for [H+].

• Mixtures of Acids: For mixtures of strong acids, add concentrations before calculating pH. For mixtures of strong and weak acids, the strong acid usually dominates.

Acidity and Basicity of Salt Solutions

Predicting the pH of Salt Solutions

Salts can affect the pH of their solutions depending on the acid-base properties of their constituent ions.

• Salts from Strong Acid and Strong Base: Neutral solution (e.g., NaCl).

• Salts from Strong Base and Weak Acid: Basic solution (e.g., NaCH3COO).

• Salts from Strong Acid and Weak Base: Acidic solution (e.g., NH4Cl).

• Salts from Weak Acid and Weak Base: pH depends on relative strengths (Ka vs. Kb).

Example: NH4Cl solution is acidic because NH4+ hydrolyzes to produce H3O+.

Predicting Acidity Based on Molecular Structure

Structural Factors Affecting Acid Strength

• Bond Polarity: The more polar the H–A bond, the stronger the acid.

• Bond Strength: Weaker H–A bonds are easier to ionize, increasing acid strength.

• Stability of Conjugate Base: Greater stability of A- increases acid strength.

• Electronegativity: For binary acids (e.g., HX), acid strength increases with electronegativity (across a period) and with bond length (down a group).

• Oxoacids: Acid strength increases with more electronegative atoms attached to the central atom and with more oxygen atoms.

Example: HClO4 > HClO3 > HClO2 > HClO in acid strength due to increasing number of oxygen atoms.

Summary Table: Acid/Base Theories and Examples