Acid strength is often confused with acidity. An acid releases hydrogen ions (\(\text{H}^+\)) when dissolved in water. Acidity is measured by the \(\text{pH}\) scale, which reflects the concentration of these released ions at a specific moment.
Acid strength is an inherent property of the molecule, representing its potential to dissociate and release \(\text{H}^+\) ions. A strong acid, like hydrochloric acid, dissociates almost completely. A weak acid, such as acetic acid, only partially dissociates, even at the same concentration. True strength is measured by quantitative metrics constant for that acid, whereas \(\text{pH}\) depends on both strength and concentration.
The Quantitative Measure: The Acid Dissociation Constant (\(\text{K}_a\))
The inherent strength of an acid is quantified by the Acid Dissociation Constant (\(\text{K}_a\)). This value is derived from chemical equilibrium, which balances the undissociated acid and its dissociated components in water. For a weak acid (\(\text{HA}\)), the reversible dissociation reaction is \(\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-\), where \(\text{A}^-\) is the conjugate base.
The \(\text{K}_a\) is the ratio of the concentration of the products (\(\text{H}^+\) and \(\text{A}^-\)) to the concentration of the undissociated acid (\(\text{HA}\)) at equilibrium. A high \(\text{K}_a\) value means the equilibrium favors the products, indicating the acid has largely dissociated and is stronger. Conversely, a very small \(\text{K}_a\) value means the acid remains mostly intact, indicating a weaker acid.
Hydrochloric acid, a strong acid, has an extremely large \(\text{K}_a\) value, indicating complete dissociation in water. Acetic acid, a weak acid, has a \(\text{K}_a\) around \(1.7 \times 10^{-5}\). The \(\text{K}_a\) is a constant for a given acid at a specific temperature, providing an absolute measure of its ability to donate a proton.
Interpreting Strength Using the \(\text{p}\text{K}_a\) Scale
Since \(\text{K}_a\) values span many orders of magnitude, chemists use the logarithmic \(\text{p}\text{K}_a\) scale for simpler comparisons. The \(\text{p}\text{K}_a\) is defined as the negative logarithm of the \(\text{K}_a\) (\(\text{p}\text{K}_a = -\log_{10}(\text{K}_a)\)). This transformation converts exponential \(\text{K}_a\) values into a more manageable range.
The interpretative rule is inverted: a lower \(\text{p}\text{K}_a\) value corresponds to a stronger acid. For example, strong hydrochloric acid has a \(\text{p}\text{K}_a\) of about \(-7\), while weak acetic acid has a \(\text{p}\text{K}_a\) of approximately \(4.76\). A difference of one unit on the \(\text{p}\text{K}_a\) scale represents a ten-fold difference in acid strength.
The \(\text{p}\text{K}_a\) scale should not be confused with \(\text{pH}\). \(\text{pH}\) measures the hydrogen ion concentration in a specific solution. \(\text{p}\text{K}_a\) is a fixed property of the acid molecule, indicating its potential to release ions. Any acid with a \(\text{p}\text{K}_a\) value less than about \(-2\) is considered a strong acid because it is almost completely dissociated in water.
Experimental Measurement Through Titration
The \(\text{p}\text{K}_a\) of an unknown weak acid is determined experimentally using titration. This procedure involves slowly adding a strong base of known concentration to a measured amount of the unknown acid. A \(\text{pH}\) meter monitors the acidity, and the resulting data is plotted to create a titration curve.
The initial stages show a slow rise in \(\text{pH}\) as the base neutralizes the acid. The most important point for determining strength is the half-equivalence point. The equivalence point is where the moles of added base exactly equal the initial moles of acid, marked by a steep vertical jump in \(\text{pH}\).
The half-equivalence point is reached when half the volume of base required for the equivalence point has been added. At this point, half of the original acid (\(\text{HA}\)) has been converted into its conjugate base (\(\text{A}^-\)), meaning their concentrations are equal. Under these conditions, the solution’s \(\text{pH}\) becomes mathematically equal to the \(\text{p}\text{K}_a\) of the weak acid. Reading the \(\text{pH}\) value directly from the titration curve at this volume determines the absolute strength of the unknown acid.
Molecular Factors That Determine Acid Strength
While \(\text{K}_a\) and \(\text{p}\text{K}_a\) quantify acid strength, the underlying cause lies in the acid’s molecular structure. The strength of any acid (\(\text{HA}\)) is determined by the stability of its conjugate base (\(\text{A}^-\)) after the hydrogen ion leaves. The more stable the resulting negative charge on the conjugate base, the stronger the original acid will be.
One significant factor is electronegativity, the atom’s ability to attract electrons. When comparing acids where the acidic hydrogen is attached to atoms in the same row, a more electronegative atom better stabilizes the negative charge of the conjugate base. This leads to a stronger acid; for example, water (\(\text{H}_2\text{O}\)) is stronger than ammonia (\(\text{NH}_3\)) because oxygen is more electronegative than nitrogen.
When comparing atoms within the same column, atomic size becomes the dominant factor. As the atom’s size increases down a column, the negative charge of the conjugate base is dispersed over a larger volume, increasing its stability. This effect makes hydroiodic acid (\(\text{HI}\)) stronger than hydrofluoric acid (\(\text{HF}\)), despite fluorine being more electronegative than iodine.
Resonance stabilization also increases acid strength. This occurs when the negative charge of the conjugate base is delocalized across multiple atoms, greatly increasing stability, as seen in carboxylic acids.