Explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model

Acid/Base Reactions > Properties of Acids and Bases >

Arrhenius’ Theory

• According to Arrhenius’ theory, acids are compounds that have hydrogen atoms and can release an H⁺ ion in aqueous solution. Such compounds are termed as Arrhenius Acids.

Some examples of Arrhenius Acids are HCl (hydrochloric acid), H₂SO₄ (sulphuric acid), CH₃COOH (acetic acid), HNO₃ (nitric acid), H₃PO₄ (phosphoric acid) etc.

• According to Arrhenius’ theory, bases are compounds which can produce hydroxyl ion in their aqueous solution. Such compounds are termed as Arrhenius Bases.

Some examples of Arrhenius Bases are NaOH (sodium hydroxide), NH₄OH (ammonium hydroxide) etc.

• Limitations of Arrhenius’ Theory:
  • Arrhenius’ theory defines acids and bases merely upon their dissociation in aqueous solutions rather than the nature of the compound. Thus, this theory is applicable for aqueous solutions only and not for non-aqueous or gaseous reactions.
  • It is only applicable for compounds having the general formula HA for acids and BOH for bases. Thus, Arrhenius’ theory cannot explain acidic properties of CuSO_4, AlCl₃, CO₂, SO₂ and the basicity of NH₃, Na₂CO₃, amines, pyridines etc.
  • This theory is unable to explain neutralization reactions that might not involve formation of water by the combination of H⁺ and OH^–. For example:

Brønsted–Lowry theory

• According to the Brønsted–Lowry theory, acids are compounds that have the ability to donate protons in aqueous solutions. Such compounds are termed as Brønsted–Lowry Acids.

Since water accepts the proton from nitric acid to form H₃​O⁺, water acts as a Brønsted-Lowry base.

• Brønsted–Lowry Bases on the other hand are compounds that have the ability to accept protons in aqueous solutions.

In this reaction, water is donating one of its protons to ammonia. After losing a proton, water becomes hydroxide, OH^–. Since water is a proton donor in this reaction, it is acting as a Brønsted-Lowry acid. Ammonia accepts a proton from water to form an ammonium ion,   ​. Therefore, ammonia is acting as a Brønsted-Lowry base.

• Strong acids and bases dissociate completely in aqueous solution. Weak acids and bases however, dissociate sparingly. For this reason, dissociation of strong acids and bases in aqueous solution is irreversible where as that of weak acids and bases is reversible.
• Two important terms are introduced in the Brønsted–Lowry concept of acids and bases. These are Conjugate Acid and Conjugate Base. Conjugate acid is a chemical entity formed after a base accepts protons. On the other hand, a conjugate base is formed when an acid donates proton.

Here, Cl^– is the conjugate base of the acid HCl because it was formed after HCl donated a proton. Again, H₃O⁺ is considered as the conjugate acid of H₂O since it was formed on water’s acceptation of a proton from hydrochloric acid.

• Limitations:
  • The Brønsted–Lowry concept cannot explain the reactions occurring in non-protonic solvents such as COCl₂, SO₂, N₂O₄, etc.
  • It cannot explain reactions of acid and basic oxides which can take place even in the absence of solvents.
  • Brønsted–Lowry theory fails to explain acidic properties of non-hydrogen containing compounds such as BF₃, AlCl₃