Bronsted Lowry Theory

The Meaning of Bronsted Lowry’s theory involves the concept of acids and bases beyond aqueous solutions and provides a broader understanding of acid-base reactions. The Bronsted Lowry theory of acids and bases definition is widely used in chemistry to describe a variety of reactions beyond the traditional acid-base scenarios.

1.0Introduction

Johannes Nicolaus Bronsted, a Danish chemist, and Thomas M. Lowry, an English chemist, introduced a broader perspective on acids and bases known as the Brönsted-Lowry theory. According to this theory, an acid is defined as a substance capable of donating a hydrogen ion (H⁺), while a base is a substance capable of accepting a hydrogen ion. In essence, acids act as proton donors, and bases act as proton acceptors. This theory provides a more comprehensive understanding of acid-base reactions beyond the traditional definitions.

According to this concept, acid and base definition will be -

• Bronsted Lowry Acid: Substances which donate H⁺ are Bronsted Lowry acids (H    donor)
• Bronsted Lowry Base: Substances which accept H⁺ are Bronsted Lowry bases (H acceptor)

2.0Conjugate Acid-Base Pairs

In a typical acid-base reaction, 

Forward reaction - Here, "HX" being a proton donor is an acid

"B" being a proton acceptor is a base.

Backward reaction -   Here, HB^⊕  being a proton donor is an acid

X^⊝ being a proton acceptor is a base.

$\begin{array}{ccccc}\text{ Acid }& \text{ Base }& & \text{ Conjugate Acid }& \text{ Conjugate Base }\\ \text{ HCl }& {\mathrm{H}}_2\mathrm{O}& \rightleftharpoons & {\mathrm{H}}_3{\mathrm{O}}^{\oplus }& +{\mathrm{Cl}}^{\ominus }\\ {\mathrm{HSO}}_4{}^{\ominus }& {\mathrm{NH}}_3& \rightleftharpoons & {\mathrm{NH}}_4{}^{\oplus }& +{\mathrm{SO}}_4{}^{2\ominus }\end{array}$

Note-  Conjugate acid-base pairs differ by only one proton.

Strong acids will have weak conjugate bases and vice versa.

3.0Bronsted Lowry Theory of Acids and Bases Examples

Consider the dissolution of NH₃ (ammonia) in H₂O (water), represented by the equation:

NH₃   +   H₂O         ⇌           NH₄⁺   +   OH ⁻

This equation represents the dissolution of ammonia (NH₃) in water (H₂O). In this reaction, ammonia acts as a base, and water acts as an acid. The ammonia molecule accepts a proton (H⁺) from water, forming the ammonium ion (NH₄⁺) and leaving behind the hydroxide ion (OH⁻ ). 

This reaction exemplifies the Bronsted-Lowry theory, where a base (NH₃) accepts a proton and forms its conjugate acid (NH₄⁺), while water acts as an acid and forms its conjugate base (OH⁻).

Acid and Conjugate base-

Base and Conjugate acid-

Amphoteric (amphiprotic): Substances which can act as acid as well as base are known as amphoteric.

$\begin{array}{rl}& \mathrm{HCl}+\underset{\text{ Base }}{{\mathrm{H}}_2\mathrm{O}}\rightleftharpoons {\mathrm{H}}_3{\mathrm{O}}^{\oplus }+{\mathrm{Cl}}^{\ominus }\\ & {\mathrm{NH}}_3+\underset{\text{ Acid }}{{\mathrm{H}}_2\mathrm{O}}\rightleftharpoons {\mathrm{NH}}_4^{\oplus }+{\mathrm{OH}}^{\ominus }\end{array}$

4.0Classification of Bronsted Lowry Theory

Bronsted - Lowery acids and bases can be -

(i)   Molecular

(ii) Cationic and

(iii) Anionic