By Sunil Bhardwaj

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As per Bronsteo-Lowry Concept the strength of an acid is determined by its tendency to give up a proton and the strength of a base is determined by its tendency to accept a proton.

A strong acid has a strong tendency to donate \({ H }^{ + }\) and a strong base has a strong tendency to accept \({ H }^{ + }\).

Aqueous solutions of acids like \(HCI{ O }_{ 4 }, HCl, { H }_{ 2 }S{ O }_{ 4 }, HN{ O }_{ 3 },\) etc., are essentially completely ionized, i.e., the reactions in which they donate protons to water proceed to almost completion. These acids are, therefore, relatively strong.

However, Aqueous solutions of acids like acetic acid, formic acid, hydrocyanic acid, etc., are partially ionized; they tend to donate protons to water only to a limited extent. Such acids are relatively weak.

Between these two extremes we have moderately strong acids.

The position of equilibrium in an acid-base reaction tells us the relative strengths of the acids and bases involved. Hydrogen chloride is a strong acid in water because the equilibrium lies far to the right. However, ionisation of a weak acid such as acetic acid occurs to a limited extent, i.e. equilibrium lies far to the left. It is therefore a weaker acid.

A Bronsted-Lowry reaction between both strong acid and strong base produce weak acid and weak base. $$ \underset { Strong Acid }{ HA } + \underset { Strong Base }{ { B }^{ - } } \rightleftharpoons \underset { Weak Base }{ { A }^{ - } } + \underset { Weak Acid }{ HB } $$