The strength of an acid is determined by its inherent ability to release a proton, or hydrogen ion (\(\text{H}^{+}\)), when dissolved in a medium, usually water. This tendency to dissociate is a fundamental chemical property that distinguishes a strong acid, which releases virtually all its protons, from a weak acid, which releases only a small fraction. Acid strength is not the same as concentration; a strong acid is inherently more potent because its molecular structure dictates how easily the bond holding the proton breaks. These structural features are the true determinants of acid strength.
The Central Role of Conjugate Base Stability
The question of why one acid is stronger than another is answered by examining what remains after the proton is released. When an acid (\(\text{HA}\)) loses a proton (\(\text{H}^{+}\)), it forms its conjugate base (\(\text{A}^{-}\)). The position of this chemical equilibrium is dictated by the stability of the resulting conjugate base. A stronger acid is one whose conjugate base is more stable. A stable conjugate base comfortably accommodates the negative charge left behind. If the resulting anion is unstable, it quickly re-associates with the proton, shifting the equilibrium back toward the original acid, making the acid weak. Conversely, a highly stable conjugate base remains dissociated, pushing the equilibrium toward the products and making the original acid strong.
How Atomic Size and Electronegativity Influence Strength
The stability of the conjugate base is profoundly affected by the inherent properties of the atom to which the acidic hydrogen is bonded: its electronegativity and its atomic size.
Electronegativity (Across a Period)
When comparing atoms across the same row (period) of the periodic table, electronegativity is the dominant factor. As electronegativity increases, the atom becomes better at accommodating the negative charge on the conjugate base, making the acid stronger. For example, acidity increases from methane (\(\text{CH}_{4}\)) to ammonia (\(\text{NH}_{3}\)) to water (\(\text{H}_{2}\text{O}\)) to hydrofluoric acid (\(\text{HF}\)) because the negative charge is placed on increasingly electronegative atoms (\(\text{C}^{-}\) to \(\text{F}^{-}\)).
Atomic Size (Down a Group)
When comparing atoms down the same column (group), atomic size becomes the overriding factor. As the atom gets larger, its ability to disperse the negative charge over a greater volume increases significantly. This charge dispersal, often called polarizability, results in a lower charge density and a much more stable conjugate base. This is why hydroiodic acid (\(\text{HI}\)) is the strongest of the hydrohalic acids; the immense size of the iodide ion (\(\text{I}^{-}\)) makes its conjugate base far more stable than the small fluoride ion (\(\text{F}^{-}\)). The increasing atomic size from fluorine to iodine means the acid strength follows the trend \(\text{HF} < \text{HCl} < \text{HBr} < \text{HI}[/latex].
The Impact of Electron Delocalization (Resonance and Induction)
Beyond the inherent properties of a single atom, the overall molecular structure can also stabilize the conjugate base through the distribution of electron density.
Resonance
One powerful mechanism is resonance, where the negative charge on the conjugate base is spread out, or delocalized, over multiple atoms. The ability to share this charge dramatically lowers the energy of the anion, leading to high stability and a much stronger acid. A clear comparison is seen between an alcohol and a carboxylic acid. The conjugate base of an alcohol localizes the negative charge entirely on a single oxygen atom, making it unstable. In contrast, the conjugate base of a carboxylic acid, called a carboxylate ion, can delocalize the negative charge across both oxygen atoms through resonance. This sharing makes the carboxylate ion significantly more stable, which is why carboxylic acids are much stronger acids than simple alcohols.
Inductive Effect
Another structural effect is the inductive effect, which involves the pull of electron density through the molecular bonds by nearby electronegative atoms or groups. These electronegative groups act as electron-withdrawing groups that help to pull electron density away from the negative charge on the conjugate base. This withdrawal effect disperses the charge, stabilizing the anion. The closer and more numerous these electron-withdrawing groups are, the stronger the effect. For instance, chloroacetic acid, which has one chlorine atom, is a stronger acid than acetic acid, and the addition of a second or third chlorine further increases the acidity.
Quantifying Acid Strength with Ka and pKa
While the structural factors provide a qualitative understanding of acid strength, chemists use quantitative measures to precisely compare the strengths of different acids. The primary measure is the acid dissociation constant, [latex]\text{K}_{a}\). The \(\text{K}_{a}\) is the mathematical expression of the acid’s equilibrium position in water. A larger \(\text{K}_{a}\) value indicates that the equilibrium lies further toward the dissociated ions, meaning the acid releases its proton more readily and is therefore a stronger acid. For example, the \(\text{K}_{a}\) for hydrochloric acid (\(\text{HCl}\)) is very large (\(10^{4}\)), while the \(\text{K}_{a}\) for acetic acid is much smaller (\(1.7 \times 10^{-5}\)). To manage the vast range of \(\text{K}_{a}\) values, the \(\text{pK}_{a}\) scale is used (\(\text{pK}_{a} = -\log_{10} \text{K}_{a}\)). The \(\text{pK}_{a}\) scale has an inverse relationship with strength: a smaller \(\text{pK}_{a}\) value indicates a stronger acid. The \(\text{pK}_{a}\) of \(\text{HCl}\) is about \(-4\) to \(-6\), compared to \(4.8\) for acetic acid. This numerical difference demonstrates how structural factors translate into measurable differences in acid strength.