9.3 Strength of Acids and Bases

This section explores the concept of acid and base strength according to the Brønsted-Lowry theory, focusing on the relative ability of substances to donate or accept protons.

Acid Strength

The strength of an acid is a measure of its ability to donate a proton ($H^{+}$).

• A strong acid is one that donates protons to a high degree.
• A weak acid is one that donates protons to a lesser degree.

Different Brønsted acids donate protons to different extents. The strength is relative; an acid is considered "strong" or "weak" in comparison to another.

Example: Hydrochloric acid ($\mathrm{HCl}$) is a stronger acid than acetic acid ($\mathrm{C}{\mathrm{H}}_3\mathrm{COOH}$). Acetic acid, in turn, is a stronger acid than water (${\mathrm{H}}_2\mathrm{O}$).

This can be represented as: $\mathrm{HCl}>\mathrm{C}{\mathrm{H}}_3\mathrm{COOH}>{\mathrm{H}}_2\mathrm{O}$

$\mathrm{HCl}$ fully ionizes in water (complete proton donation), whereas $\mathrm{C}{\mathrm{H}}_3\mathrm{COOH}$ only partially ionizes (partial proton donation).

Base Strength

The strength of a base is a measure of its ability to accept a proton ($H^{+}$).

• A strong base is one that accepts protons to a high degree.
• A weak base is one that accepts protons to a lesser degree.

Similar to acids, the strength of bases is relative.

Example: Ammonia ($\mathrm{N}{\mathrm{H}}_3$) is a relatively stronger base than water (${\mathrm{H}}_2\mathrm{O}$) because it has a greater ability to accept a proton.

$\mathrm{N}{\mathrm{H}}_3>{\mathrm{H}}_2\mathrm{O}$

Inverse Relationship: Acid/Base Strength and Conjugate Pairs

There is an inverse relationship between the strength of an acid and the strength of its conjugate base:

• A strong acid has a weak conjugate base.
• A weak acid has a strong conjugate base.

For example, $\mathrm{C}{\mathrm{H}}_3\mathrm{CO}{\mathrm{O}}^{-}$ (conjugate base of the weak acid $\mathrm{C}{\mathrm{H}}_3\mathrm{COOH}$) is a stronger base than $\mathrm{C}{\mathrm{l}}^{-}$ (conjugate base of the strong acid $\mathrm{HCl}$).

See for more on conjugate pairs.

The Levelling Effect

In aqueous solutions, the strength of strong acids is limited by the basicity of water. This is known as the Levelling Effect.

Water acts as a base and reacts with any acid stronger than the hydronium ion ($H_3{O}^{+}$) to form $H_3{O}^{+}$:

$\mathrm{HA}+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{H}}_3{\mathrm{O}}^{+}+{\mathrm{A}}^{-}$

Consequently, all strong acids like $\mathrm{HCl}$, $\mathrm{HN}{\mathrm{O}}_3$, and $\mathrm{HCl}{\mathrm{O}}_4$ appear to have the same strength in water — they are all levelled to the strength of $H_3{O}^{+}$, which is the strongest acid that can exist in aqueous solution.

To differentiate the strengths of strong acids, a non-aqueous solvent (e.g., glacial acetic acid) with weaker basicity than water must be used. This is called the differentiating effect.

Summary

• Acid order: $\mathrm{HCl}>\mathrm{C}{\mathrm{H}}_3\mathrm{COOH}>{\mathrm{H}}_2\mathrm{O}$
• Base order: $\mathrm{N}{\mathrm{H}}_3>{\mathrm{H}}_2\mathrm{O}$

See Strong and Weak Acids→ and Strong and Weak Bases→ for detailed discussion of ionization differences.