Arrhenius Theory

The Arrhenius theory, introduced by Swedish scientist Svante Arrhenius in 1887, defines acids as substances that dissociate in water to produce ions, including hydrogen ions (H⁺), and bases as substances that ionize in water to yield hydroxide ions (OH⁻). However, it's now understood that hydrogen ions in water exist in combination with water molecules as hydronium ions (H₃O⁺). Despite this clarification, the term "hydrogen ion" is still commonly used to refer to the hydronium ion.

1.0Definition of Arrhenius Theory

The Definition of Arrhenius theory involves the explanation of acidic properties of well-known acids (e.g., sulfuric, hydrochloric, nitric, and acetic acids) based on their ability to yield hydrogen ions and the basic properties of hydroxides (e.g., sodium, potassium and calcium hydroxides) based on their ability to produce hydroxide ions in solution. Acids and bases are further classified as strong or weak depending on the concentration of hydrogen or hydroxide ions they generate.

In practical terms, the reaction between an acid and a base results in the formation of a salt and water, with water arising from the combination of a hydrogen ion and a hydroxide ion. While the Arrhenius theory laid the foundation for understanding acid-base behavior, later theories, such as the Brønsted-Lowry and Lewis theories, provided more comprehensive views of acid-base interactions by considering proton transfer and electron pair donation.

2.0Arrhenius Concept

Arrhenius Theory of  Acid: 

Substance which gives H^⊕ ion on dissolving in water (H^⊕ donor)

For Example-  HNO₃, HClO₄, HCl, HI, HBr, H₂SO₄, H₃PO₄ etc.

Types of Acid - 

Arrhenius Theory of Base:

Any substance which releases OH^⊝ (hydroxyl) ion in water (OH^⊝ ion donor)

Types of Base

Strength of Acid or Base

 (a) The strength of acids or bases is determined by the extent of their ionization, which is reflected in the equilibrium constant (Ka for acids, Kb for bases) associated with their ionization reactions. The equilibrium constant provides a quantitative measure of the degree to which an acid or base dissociates in water, indicating its strength.

 (b) For acids, the ionization equilibrium is represented as:

$\mathrm{HA}\rightleftharpoons {\mathrm{H}}^{\oplus }+{\mathrm{A}}^{\ominus }$;

${\mathrm{K}}_{\mathrm{a}}=\frac{\left[{\mathrm{H}}^{\oplus }\right]\left[{\mathrm{A}}^{\ominus }\right]}{[\mathrm{HA}]}$ = dissociation or ionization constant of acid.   

(c)  Similarly, for bases, the ionization equilibrium is represented as:

$\mathrm{BOH}\rightleftharpoons {\mathrm{B}}^{\oplus }+{\mathrm{OH}}^{\ominus };{\mathrm{K}}_{\mathrm{b}}=\frac{\left[{\mathrm{B}}^{\oplus }\right]\left[{\mathrm{OH}}^{\ominus }\right]}{[\mathrm{BOH}]}$=   dissociation or ionization constant of base

(d) The Larger the value of K_a or K_b, the stronger the acid or base respectively.

3.0Limitations of Arrhenius Theory

The Arrhenius theory, though foundational, has some limitations that became apparent as the understanding of acid-base chemistry evolved. Here are some of  Arrhenius theory limitations-

Aqueous Limitation:

• The Arrhenius theory is restricted to describing acid-base behavior solely in aqueous solutions, limiting its applicability to non-aqueous systems.

Hydrogen Ion Misconception:

• Originally proposed as releasing free hydrogen ions (H⁺), the theory did not consider the formation of hydronium ions (H₃O⁺) when hydrogen ions combine with water molecules.

Exclusion of Lewis Concept:

• The theory does not encompass the broader Lewis acid-base concept, which involves electron-pair transfer, focusing solely on proton transfer.

Inapplicability to Coordination Compounds:

• The theory does not extend to the behavior of acids and bases in coordination compounds, where metal ions exhibit acid-base characteristics.