The strengths of acids and bases determine their reactivity and dissociation in solution. Strong acids and bases, like hydrochloric acid and sodium hydroxide, completely dissociate, while weak ones, like acetic acid and ammonia, only partially dissociate. Recognising these strengths is key to understanding their behaviour in chemical reactions and their diverse applications.
Acids can be categorised as strong acids or weak acids depending on their ability to completely transfer protons in an aqueous solution. The extent of hydrogen transfer depends on the molecular structure of the acid. Similarly, bases can be categorised as strong bases or weak bases depending on their ability to completely accept donated protons in aqueous solution.
Examples of strong and weak acids and bases are shown in the table.
Strong acid
Weak acid
Strong base
Weak base
Hydrochloric acid, \(\ce{HCl}\)
Acetic acid, \(\ce{CH3COOH}\)
Ammonia, \(\ce{NH3}\)
Sodium hydroxide, \(\ce{NaOH}\)
Hydrobromic acid, \(\ce{HBr}\)
Formic acid, \(\ce{HCOOH}\)
Methylamine, \(\ce{CH3NH2}\)
Potassium hydroxide, (\(\ce{KOH}\)
Nitric acid, \(\ce{HNO3}\)
Carbonic acid, \(\ce{H2CO3}\)
Aniline, \(\ce{C6H5NH2}\)
Calcium hydroxide, \(\ce{Ca(OH)2}\)
Sulfuric acid, \(\ce{H2SO4}\)
Phosphoric acid, \(\ce{H3PO4}\)
Pyridine, \(\ce{C5H5N}\)
Barium hydroxide, \(\ce{Ba(OH)2}\)
Perchloric acid, \(\ce{HClO4}\)
Hydrofluoric acid, \(\ce{HF}\)
Ethylamine, \(\ce{C2H5NH2}\)
Lithium hydroxide, \(\ce{LiOH}\)
Strong acids
Strong acids can transfer nearly all (\(100\%\)) of their protons (hydrogen ions) to water. If the hydrogen is bound to a highly electronegative atom like oxygen or chlorine, the hydrogen is acidic and is able to undergo complete ionisation. The dominant species in solution would be \(\ce{H3O}^{+}\) and the conjugate base \(\ce{A}^{-}\) (the right side of the reaction arrow). An example is hydrochloric acid (\(\ce{HCl}\)). Acid–base reactions involving strong acids are written using reaction arrows (\(\rightarrow\)).
Weak acids can only transfer some of their protons to the water. If the hydrogen is bound to a less electronegative atom like carbon, the bond is not polar enough to release the hydrogen. The dominant species in solution would be the acid \(\ce{HA}\) (the left side of the reaction arrow). An example is the hydrogen bound to carbon in acetic acid (\(\ce{CH3COOH}\)), which makes it a weak acid. Acid–base reactions involving weak acids are written using equilibrium arrows (\(\rightleftharpoons\)).
The harpoon pointing towards the products (forward reaction) is larger than the one pointing to the reactants (backwards reaction), indicating the preference for formation of products.
Weak acid
Chemical equation: HA + H2O ⇌ H3O+ + A–
The harpoon pointing towards the products (forward reaction) is smaller than the one pointing to the reactants (backwards reaction), indicate the preference for formation of reactants.
Strong bases
Strong bases show a high affinity for protons in an aqueous solution. \(\ce{NaOH}\), \(\ce{KOH}\) and hydroxides of group \(1\) and \(2\) elements in the periodic table are strong bases. In a solution of a strong base, the dominant species in the solution would be \(\ce{OH}^{-}\) and the conjugate acid \(\ce{BH}^{+}\). Acid–base reactions involving strong bases are written using reaction arrows (\(\rightarrow\)).
Weak bases show less affinity for protons in aqueous solutions. \(\ce{NH3}\) gas is a weak base. In a solution of a weak base, the dominant species in the solution is the base \(\ce{B}\). Acid–base reactions involving weak bases are written using equilibrium arrows (\(\rightleftharpoons\)).
The harpoon pointing towards the products (forward reaction) is larger than the one pointing to the reactants (backwards reaction), indicating the preference for formation of products.
Weak base
Chemical equation: B + H2O ⇌ BH+ + OH–
The harpoon pointing towards the products (forward reaction) is smaller than the one pointing to the reactants (backwards reaction), indicate the preference for formation of reactants.
Relationship between acid–base strength
There is an inverse relationship between the strength of an acid and the strength of a base.
The conjugate base of a strong acid is always a weak base.
The conjugate base of a weak acid is always a strong base.
This relationship can be explained using the definition of a strong/weak acid and strong/weak base. The strong acid has the capacity to completely (nearly 100%) dissociate in water to donate protons, which means its conjugate base has less affinity to protons as the reaction favours the formation of products. A base showing less affinity for protons is a weak base.
Similarly, in water, the weak acid partially dissociates to produce protons, which means its conjugate base displays a high affinity for the protons. The reaction favours the formation of reactants.
Therefore, the stronger the acid, the weaker its conjugate base, and the weaker the acid, the stronger its conjugate base.
Ionisation constants
The strengths of acids and bases can be quantified using ionisation constants. These constants are essentially the ratio of products (dissociated acid or base) compared to reactants (undissociated acid or base).
Acid ionisation constants – \(K_{a}\)
Recall that the dissociation of a weak acid in water can be written as:
Once the concentrations of \(\ce{H3O}^{+}\), \(\ce{A}^{-}\) and \(\ce{HA}\) are experimentally determined, the acid ionisation constant (\(K_{a}\)) can be calculated using:
The strength of an acid is directly proportional to the acid ionisation constant—that is, the higher the constant, the stronger the acid. Strong acids have larger \(K_{a}\) values (greater than \(1\)), as the formation of products is favoured; there is a larger concentration of \(\ce{H3O}^{+}\) and \(\ce{A}^{-}\) than \(\ce{HA}\), Weak acids have \(K_{a}\) values less than \(1\), as ionisation is not favoured.
Water is not included in the ionisation constant expression because it is the solvent and its concentration is effectively constant.
Base ionisation constants – \(K_{b}\)
Recall that the dissociation of a weak base in water can be written as:
The strength of a base can also be predicted by an ionisation constant, the base ionisation constant (\(K_{b}\)). The general expression for \(K_{b}\) is:
Step 3: Calculate the concentrations of \(\ce{H3O}^{+}\), \(\ce{A}^{-}\) and \(\ce{HA}\) using the percentage ionisation.
The concentration of \(\ce{H3O}^{+}\) will be \(10\%\) of the original concentration of \(\ce{HA}\).
\[\begin{align*}
\textrm{[}\ce{H3O}^{+}\textrm{]} & = 10\% \times 0.0200\textrm{M}\\
& = 0.0200\textrm{M}
\end{align*} \]
Using the molar ratio:
\[\begin{align*}
\textrm{[}\ce{A}^{-}\textrm{]} & = \frac{1}{1} \times \textrm{[}\ce{H3O}^{+}\textrm{]}\\
& = 0.00200\textrm{M}
\end{align*} \]
Since the concentrations are molar, \(\textrm{[}\ce{H3O}^{+}\textrm{]} + \textrm{[}\ce{HA}^{+}\textrm{]} = \textrm{[}\ce{A}^{-}\textrm{]}\). We can calculate:
\[\begin{align*}
\textrm{[}\ce{HA}\textrm{]} & = 0.0200\textrm{M}-0.00200\textrm{M}\\
& =0.0180\textrm{M}
\end{align*}\]
Step 4 Substitute the concentration values into the acid ionisation constant expression:
The units for ionisation constants can vary. In this case, the numerator has \(\textrm{M}\times\textrm{M}\) and the denominator has just \(\textrm{M}\). \(\frac{\textrm{M}^{2}}{\textrm{M}}=\textrm{M}\), so the units for this example are just \(\textrm{M}\). Thus, the correct answer is \(2.2\times10^{-4}\textrm{M}\).
