Which is a stronger acid, sulfurous acid or sulfuric acid? Why?

Understanding the acid dissociation constant (Ka) is fundamental when comparing the strength of acids. The Ka value measures an acid's tendency to donate protons to water, thus forming hydronium (H3O+) ions in solution. In other words, the greater the Ka value, the stronger the acid.

For example, a high Ka value implies a greater concentration of hydronium ions, indicative of a strong acid that dissociates readily in water. On the other hand, a low Ka value signifies a weak acid, which doesn't dissociate as much, retaining more of its original molecules. This measurement allows chemists to quantitatively compare the strength of acids in an aqueous environment.

Proton donation is the hallmark of acidic behavior. When an acid donates a proton (H+ ion), it demonstrates its acidity. This process is part of a chemical equilibrium in water, where acids and their conjugate bases exist in a balance. Strong acids are characterized by their willingness to easily and fully donate their protons, shifting the equilibrium far towards the side of the conjugate base and the hydronium ions (H3O+).

By assessing how readily different acids donate their protons, we can rank their strength. This is why sulfuric acid, being able to donate protons completely and efficiently in an aqueous solution, is stronger than sulfurous acid, which doesn't donate protons as readily.

Sulfurous acid (H2SO3) and sulfuric acid (H2SO4) may seem similar at a glance, but they exhibit notably different acid strengths. These differences arise from their molecular structures and the number of oxygen atoms bonded to the sulfur atom.

When comparing the two, sulfuric acid is a much stronger acid. This is evidenced by its behavior in water, where it completely dissociates into H+ ions and sulfate ions (SO4^2-). This complete dissociation shows a high degree of proton donation and accounts for sulfuric acid's high Ka value. Sulfurous acid's lower Ka value indicates partial dissociation and thus a weaker acid strength. The higher number of oxygen atoms in sulfuric acid also contributes to the greater stability of its conjugate base.

The strength of an acid is intrinsically linked to the stability of its conjugate base. A stable conjugate base forms when an acid effectively donates its proton, and the remaining anion is able to spread the negative charge evenly across its structure.

The sulfate ion (SO4^2-), which is the conjugate base of sulfuric acid, has multiple resonance structures, allowing the negative charge to be delocalized over four oxygen atoms. This delocalization confers great stability to the sulfate ion. In contrast, the sulfite ion (SO3^2-), being the conjugate base of sulfurous acid, can only spread its negative charge across three oxygens. This lesser degree of delocalization results in a less stable conjugate base. The stability of the conjugate base is a direct measure of an acid's strength, with more stable conjugate bases aligning with stronger acids.