6.1.6 – Acid/Base Theories

Dot-Point 6: explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to: Arrhenius’ theory, Brønsted–Lowry theory (ACSCH064, ACSCH067)

Syllabus-Listed Theories (Key Theories)

Arrhenius' Theory (1887)

• Definition: Svante Arrhenius defined acids as substances that ionise in water to produce H⁺ ions, while bases are substances that dissociate to form OH⁻ ions in aqueous solutions. 

  • Example: 

    • Acids: HA(aq) → H⁺(aq) + A⁻(aq)

    • Bases: BOH(aq) → B⁺(aq) + OH⁻(aq)

  • Neutralisation: H⁺(aq) + OH⁻(aq) → H₂O(l)

• Improvements: 

  • Arrhenius’ theory explained the ionising properties of acids, which accounted for the conductivity of solutions and laid the foundation for the pH concept.

  • It provided a simple and clear explanation of neutralisation reactions.

• Limitations: 

  • Limited Scope: The theory only applies to acids of the form HA and bases of the form BOH. It could not account for acidic oxides like CO₂, SO₂, or AlCl₃, nor for bases like NH₃ (ammonia) or Na₂CO₃ (sodium carbonate).

  • Solvent Limitation: Arrhenius' theory does not account for acid-base reactions in solvents other than water, such as in acetone or other non-aqueous solvents.

  • Inadequate for Certain Reactions: It cannot explain acid-base reactions that do not occur in solution. For example, the reaction between gaseous ammonia and hydrochloric acid: 

    • NH₃(g) + HCl(g) ⇋ NH₄Cl(s)

  • Unexpected Results: Equal amounts of HCl and ammonia result in a slightly acidic solution, which contradicts Arrhenius’ idea that acids and bases always neutralise each other to form neutral solutions.

Bronsted-Lowry Theory (1923)

• Definition: The Bronsted-Lowry theory, proposed independently by Johannes Bronsted and Martin Lowry, defines acids as proton (H⁺) donors and bases as proton (H⁺) acceptors. 

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

    • Example: CH₃COOH (acetic acid) + H₂O → CH₃COO⁻ + H₃O⁺

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

    • Example: NH₃ + H₂O → NH₄⁺ + OH⁻

  • Neutralisation: An acid-base reaction is a proton transfer, where the acid donates a proton to the base. 

    • Example: NH₃(g) + HCl(g) ⇋ NH₄Cl(s)

  • Conjugate Acid-Base Pairs: When an acid donates a proton, the remaining part of the acid becomes its conjugate base. Similarly, when a base accepts a proton, it forms its conjugate acid. 

    • Example: HCl → Cl⁻ (conjugate base); H₂O → H₃O⁺ (conjugate acid)

• Improvements: 

  • This theory expanded the definition of acids and bases beyond aqueous solutions. It explained why some compounds that do not contain OH⁻ or H⁺ could still act as acids or bases.

  • It provided a better understanding of amphiprotic substances (e.g., water), which can act as both acids and bases depending on the circumstances.

  • The theory is applicable in non-aqueous solvents, explaining acid-base behaviour beyond water.

• Limitations: 

  • Acidic and Basic Oxides: The Bronsted-Lowry theory does not explain the acidity of acidic oxides (e.g., SO₂, SO₃) or the basicity of basic oxides (e.g., MgO, CaO), as these do not involve proton transfer.

  • Non-Proton Transfer Reactions: It cannot explain acid-base reactions that do not involve proton transfer, such as reactions between acidic and basic oxides: 

    • SO₃(g) + CaO(s) → CaSO₄(s)

  • The theory also struggles with reactions where no proton transfer occurs.

Key Differences: Arrhenius vs. Bronsted-Lowry

• Acids: 

  • Arrhenius: Acids are substances that release H⁺ ions in solution.

  • Bronsted-Lowry: Acids are proton donors, which may or may not release H⁺ ions in solution, depending on the context.

• Bases: 

  • Arrhenius: Bases are substances that release OH⁻ ions in solution.

  • Bronsted-Lowry: Bases are proton acceptors and can form OH⁻ ions indirectly by accepting a proton from water (e.g., NH₃).