64 Strength of Acids

Strong Acids

In water, strong acids completely dissociate into free protons and their conjugate base.

Definition of Strong Acids

The strength of an acid refers to the ease with which the acid loses a proton. A strong acid ionizes completely in an aqueous solution by losing one proton, according to the following equation:

[latex]\text{HA} (aq) \rightarrow \text{H}^+ (aq) + \text{A}^- (aq)[/latex]

where [latex]\text{HA}[/latex] is a protonated acid, [latex]\text{H}^+[/latex] is the free acidic proton, and [latex]\text{A}^-[/latex] is the conjugate base. Strong acids yield weak conjugate bases. For sulfuric acid, which is diprotic, the “strong acid” designation refers only to the dissociation of the first proton:

[latex]\text{H}_2\text{SO}_4 (aq) \rightarrow \text{H}^+ (aq) + \text{HSO}_4^- (aq)[/latex]

More precisely, the acid must be stronger in aqueous solution than a hydronium ion ([latex]\text{H}^+[/latex]), so strong acids have a [latex]\text{pKa}[/latex] < -1.74. An example is hydrochloric acid ([latex]\text{HCl}[/latex]), whose [latex]\text{pKa}[/latex] is -6.3. This generally means that in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution.

Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution:

[latex][ \text{H}^+ ] = [ \text{A}^- ] = [ \text{HA} ]_\text{total}[/latex] and [latex]\text{pH} = \text{-log}[ \text{H}^+ ][/latex].

Strong acids, like strong bases, can cause chemical burns when exposed to living tissue.

Examples of Strong Acids

Some common strong acids (acids with [latex]\text{pKa}[/latex] < -1) include:

• Hydroiodic acid ([latex]\text{HI}[/latex]): [latex]\text{pKa}[/latex] = -9.3
• Hydrobromic acid ([latex]\text{HBr}[/latex]): [latex]\text{pKa}[/latex] = -8.7
• Perchloric acid ([latex]\text{HClO}_4[/latex]₎: [latex]\text{pKa}[/latex] ≈ -8
• Hydrochloric acid ([latex]\text{HCl}[/latex]): [latex]\text{pKa}[/latex] = -6.3
• Sulfuric acid ([latex]\text{H}_2\text{SO}_4[/latex]): [latex]\text{pKa1}[/latex] ≈ -3 (first dissociation only)
• p-Toluenesulfonic acid: [latex]\text{pKa}[/latex] = -2.8
• Nitric acid ([latex]\text{HNO}_3[/latex]): [latex]\text{pKa}[/latex] ≈ -1.4
• Chloric acid ([latex]\text{HClO}_3[/latex]): [latex]\text{pKa}[/latex] ≈ 1.0

Weak Acids

A weak acid only partially dissociates in solution.

A weak acid is one that does not dissociate completely in solution; this means that a weak acid does not donate all of its hydrogen ions ([latex]\text{H}^+[/latex]) in a solution. Weak acids have very small values for [latex]\text{K}_a[/latex] (and therefore higher values for [latex]\text{pK}_a[/latex] ) compared to strong acids, which have very large [latex]\text{K}_a[/latex] values (and slightly negative [latex]\text{pK}_a[/latex] values).

The majority of acids are weak. On average, only about 1 percent of a weak acid solution dissociates in water in a 0.1 mol/L solution. Therefore, the concentration of [latex]\text{H}^+[/latex] ions in a weak acid solution is always less than the concentration of the undissociated species, [latex]\text{HA}[/latex]. Examples of weak acids include acetic acid ([latex]\text{CH}_a\text{COOH}[/latex]), which is found in vinegar, and oxalic acid ([latex]\text{H}_2\text{C}_2\text{O}_4[/latex]), which is found in some vegetables.