Sulfurous acid, chemically represented as H₂SO₃, is classified as a weak acid. It forms when sulfur dioxide gas (SO₂) dissolves in water, creating an aqueous solution that is mildly acidic. Unlike its stronger counterpart, sulfuric acid, sulfurous acid does not fully release its protons into the solution. This partial interaction results in a lower concentration of free hydrogen ions, the chemical species responsible for acidity.
Defining Acid Strength
The classification of an acid as strong or weak depends on its behavior when dissolved in water. Acids donate a proton (H⁺) to a water molecule, forming the hydronium ion (H₃O⁺). A strong acid, such as hydrochloric acid, undergoes virtually complete ionization, meaning nearly every acid molecule breaks apart to release its proton. This process is irreversible and creates a high concentration of H₃O⁺ ions in the solution.
Weak acids, by contrast, only partially ionize in an aqueous environment. The dissociation reaction is a reversible process, establishing a chemical equilibrium between the intact acid molecules and the dissociated ions. The majority of the acid molecules remain undissociated, holding onto their protons. This partial ionization results in a much lower concentration of hydronium ions compared to a strong acid of the same concentration.
Sulfurous Acid: The Chemistry of Weakness
When H₂SO₃ dissolves, most of the molecules remain in their original form rather than donating their hydrogen ions. The equilibrium between the intact acid and its ions is quantified by the acid dissociation constant, known as Kₐ.
The first dissociation step for sulfurous acid is represented by a Kₐ₁ value of approximately \(1.5 \times 10^{-2}\). This small numerical value indicates that the equilibrium strongly favors the undissociated H₂SO₃ molecules over the dissociated ions. Since acids with Kₐ values less than one are considered weak, this value places sulfurous acid firmly in that category.
The structure of sulfurous acid allows it to release two protons, making it a diprotic acid. However, the second proton is significantly harder to remove than the first. The Kₐ₂ for the second dissociation step is much smaller, around \(6.4 \times 10^{-8}\), indicating that the second proton is only released to a negligible extent.
Sulfurous Acid’s Existence in Water
Sulfurous acid is highly unstable and cannot be isolated in a pure, water-free form. When sulfur dioxide (SO₂) is dissolved in water (H₂O), the resulting solution is often referred to as sulfurous acid. The reaction produces the acidic solution, following the equation SO₂ + H₂O ⇌ H₂SO₃.
A significant portion of the dissolved substance does not actually form the H₂SO₃ molecule. Instead, it exists primarily as dissolved SO₂ molecules in equilibrium with the bisulfite ion (HSO₃⁻) and the hydronium ion. This means that the majority of the substance in the solution exists in a form that cannot readily donate protons. The difficulty in even forming the H₂SO₃ molecule further limits its ability to release H⁺ ions, reinforcing its classification as a weak acid.
Contrast with Sulfuric Acid
A common point of confusion arises when comparing sulfurous acid (H₂SO₃) with sulfuric acid (H₂SO₄). Sulfuric acid is a strong acid, meaning its first proton dissociates completely in water. The difference in strength stems directly from the molecular structure of the two compounds. Both are oxoacids, but sulfuric acid has four oxygen atoms attached to the central sulfur atom, while sulfurous acid has only three.
The greater number of highly electronegative oxygen atoms in sulfuric acid pulls electron density away from the sulfur atom. This electron withdrawal effect stabilizes the negative charge on the resulting conjugate base (HSO₄⁻) after the first proton is released. The resulting stability makes it much easier for sulfuric acid to donate its proton, which explains its strong acid character. Sulfurous acid, with one less oxygen atom, has a less stable conjugate base (HSO₃⁻), causing it to hold onto its proton more tightly and resulting in its weak acid status.