IB Chemistry: HL Notes on Acid-Base Reactions and Properties

Brønsted–Lowry Acids & Bases

Brønsted-Lowry Theory

• Definition: Acids and bases are defined based on their ability to donate or accept protons (H+).
• Brønsted-Lowry Acid: A species that donates a proton.
• Brønsted-Lowry Base: A species that accepts a proton using a lone pair of electrons.
• Reactions: The theory applies to aqueous and gas-phase reactions.

Example of Proton Transfer Reaction

• Reaction: H₂PO₄⁻ (aq) + H₂O (l) → HPO₄²⁻ (aq) + H₃O⁺ (aq)
  • Identifying Acid and Base: H₂PO₄⁻ is the acid (donates H⁺), H₂O is the base (accepts H⁺).
  • Answer Choice: A) H₂PO₄⁻ is the acid, H₂O is the base.

Conjugate Acids & Bases

Definition

• Conjugate Acid-Base Pair: Two species differing by one proton (H+).
• Example: CH₃COOH (acetic acid) ⇌ CH₃COO⁻ (acetate) + H⁺.

Worked Example for Conjugate Pairs

• Reaction: CH₃CH₂CH₂COOH + H₂O ⇌ CH₃CH₂CH₂COO⁻ + H₃O⁺.
  • Identifying Pairs: B) H₂O and H₃O⁺ are the conjugate pair.

Amphiprotic Species

Definition

• Amphiprotic Species: Can act as both proton donors and acceptors.
• Example: Water (H₂O) can donate protons (acting as an acid) or accept protons (acting as a base).

Importance of Amphiprotic Species

• In reactions such as:
  • H₂O + NH₃ ⇌ H₂O (acting as an acid) + NH₄⁺ (ammonium)
  • H₂O + HCl ⇌ H₃O⁺ (hydronium) + Cl⁻ (chloride)

The pH Scale

Definition

• pH: Measure of acidity or alkalinity of a solution.
  • Formula: pH = -log[H⁺]
  • [H⁺]: Concentration of hydrogen ions in mol/dm³.

Characteristics

• Scale: 0-14 (Acidic: pH < 7, Neutral: pH = 7, Basic: pH > 7).
• Change in pH indicates an exponential change in the acidity/basicity:
  • Example: pH 5 is 10 times more acidic than pH 6.

Calculation Example of pH

• Given a solution with a known H⁺ concentration, calculate pH using the formula.

Ion Product of Water

Definition

• Kw = [H⁺][OH⁻]
• Temperature influences the ion product, where higher temperatures generally increase ion concentration and thus affect pH.

Strong & Weak Acids

Definition

• Strong Acid: Completely dissociates in solution (e.g., HCl, H₂SO₄).
  • Example: HCl → H⁺ + Cl⁻ (dissociation diagram).
• Weak Acid: Partially dissociates in solution (e.g., CH₃COOH).
  • Example: CH₃COOH ⇌ H⁺ + CH₃COO⁻.

Comparing Acids

• Strength Related to Ionization:
  • Strong acids produce high H⁺ concentrations leading to low pH.
  • Weak acids maintain a higher pH due to lower H⁺ concentrations.

Neutralisation Reactions

Definition

• Acid reacts with base to form water plus a salt:
  • General Equation: Acid + Base → Salt + Water.

Example Reaction

• HCl + NaOH → NaCl + H₂O (neutralisation).
• Spectator ions do not affect pH but are necessary to form the salt.

pH Curves

Definition

• pH curves illustrate the relationship between the volume of titrant added and the pH of the solution.
• Equivalence Point: Point where stoichiometrically equivalent amounts of acid and base have reacted (typically indicates complete neutralization).

Strong Acid - Strong Base Titration Example

• Observations: pH starts low (acidic) and rises sharply near equivalence point (pH ≈ 7).

Buffer Solutions

Definition

• A buffer solution resists changes in pH upon the addition of small amounts of acid or base.

Components

• Typically consists of a weak acid and its conjugate base (e.g., CH₃COOH and CH₃COONa).

pH Calculation for Buffers

• pH can be calculated using the Henderson-Hasselbalch equation:
  • pH = pKa + log([A⁻]/[HA]).