Successes and limitations of Arrhenius acid and base theory

In 1887, Swedish scientist Svante Arrhenius proposed definitions for acids and bases, which is widely known as the Arrhenius Theory.

Neutralization Reaction According to the Theory

According to this theory, the reaction between an acid and a base produces salt and water. In essence, neutralization is the process where H⁺ ions from the acid combine with OH^– ions from the base to form a neutral molecule, water.

Complete Reaction: HCl(aq) + NaOH(aq) → Na⁺(aq) + Cl^–(aq) + H₂O(l)

1. Acid-Base Titration: The fundamental principle of acid-base titration and the mechanism of indicators can be easily explained with this theory.
2. Buffer Solution: The acid-base mechanism of buffer solutions can be seamlessly described using the concept of aqueous ionization inherent to this theory.
3. Measurement of Acid-Base Strength: The quantitative strength of an acid or a base can be expressed through their respective dissociation constants (K_a or K_b).
   For a weak acid HA: HA + aq ⇌ H⁺(aq) + A^–(aq)
   Dissociation constant, K_a = [H⁺][A^–] [HA]
   An increase in the value of the dissociation constant indicates a higher strength of the acid, while a decrease indicates a lower strength. The same principle applies equally to bases (K_b).
4. Reactions in Aqueous Solutions: The Arrhenius concept is entirely successful in explaining various chemical reactions and properties of acids and bases in aqueous mediums.
5. Explanation of Constant Heat of Neutralization: In the neutralization reaction between any strong acid and a strong base, the heat evolved per mole of water formed is always constant (-57.34 kJ mol^-1 or -13.7 kcal mol^-1). This theory easily justifies this constancy because, in all such cases, the reaction fundamentally involves the combination of the same H⁺ and OH^– ions to produce water.
6. Nature of Aqueous Solutions: The underlying reason why an aqueous solution of a substance behaves as acidic or basic is explicitly revealed through the presence of free ions as stated by this theory.
7. Acid Catalysis: Acids act as catalysts in several chemical reactions. In those cases, the catalytic activity can be accurately explained by the presence of H⁺ ions produced from acid dissociation.

1. Hypothetical Existence of Free H⁺ Ions: The Arrhenius theory assumes the independent existence of free hydrogen ions (H⁺) in aqueous solutions. However, free H⁺ ions cannot exist independently in water; they bind with water molecules and exist as hydronium ions (H₃O⁺).
2. Limitation to Aqueous Medium: This theory is strictly limited to aqueous solutions. It cannot explain the acid-base behavior of substances in the absence of water or in non-aqueous solvents (such as alcohol, liquid ammonia).
3. Inability to Explain Bases Lacking OH^– Ions: According to this framework, only compounds containing hydroxyl groups (OH^–) qualify as bases. It provides no explanation as to why substances like NH₃, metallic oxides (CaO, MgO), and alkylamines (R-NH₂) act as strong bases despite lacking OH^– groups.
4. Reactions in Gaseous State: This concept fails to account for the acid-base behavior of compounds in the gaseous phase. For instance, it cannot prove that the reaction NH₃(g) + HCl(g) → NH₄Cl(s) is an acid-base neutralization reaction.
5. Nature of Aqueous Salt Solutions: The theory fails to clarify why aqueous solutions of certain salts like FeSO₄, ZnSO₄, CuSO₄, AlCl₃, BF₃ are acidic, whereas an aqueous solution of Na₂CO₃ is basic.
6. Anomaly of Acidic and Basic Salts: The aqueous solution of monosodium phosphate (NaH₂PO₄) is acidic, yet the aqueous solution of its disodium salt (Na₂HPO₄) is basic. This anomaly can by no means be explained through the Arrhenius concept.
7. Reactions in Other Solvents: In liquid NH₃ solvent, dissolved NH₄Cl behaves as an acid and sodamide (NaNH₂) behaves as a base. Due to the non-aqueous medium, their acid-base properties cannot be evaluated using this theory.