Explain the relationship between the strength of an acid or base and the structure of the molecule or ion.
The protons on a molecule that will participate in acid–base reactions, and the relative strength of those protons, can be inferred from the molecular structure.
The organizing principle: an acid is strong when its conjugate base is stable. Anything that spreads out or stabilizes the negative charge left behind makes it easier for the proton to leave.
Binary acids H–X, down a group: HF < HCl < HBr < HI. Going down, the H–X bond gets longer and weaker, and the resulting X⁻ is larger, spreading its charge over more volume. Bond strength dominates electronegativity here — HF is a weak acid despite fluorine being the most electronegative element.
Binary acids across a period: CH₄ < NH₃ < H₂O < HF. Now electronegativity dominates, because bond length barely changes. A more electronegative atom stabilizes the negative charge on the conjugate base better.
Oxyacids with more oxygens: HClO < HClO₂ < HClO₃ < HClO₄. Each additional oxygen withdraws electron density, weakening the O–H bond and — more importantly — delocalizing the negative charge of the anion over more oxygen atoms through resonance. Perchlorate spreads its charge over four oxygens and is exceptionally stable, which is why HClO₄ is a strong acid.
Oxyacids with the same number of oxygens: HOI < HOBr < HOCl. A more electronegative central atom pulls electron density away from the O–H bond, weakening it.
Organic acids: a carboxylic acid (–COOH) is far more acidic than an alcohol (–OH) because the carboxylate anion is resonance stabilized across two equivalent oxygens, while an alkoxide has its charge localized on one oxygen. Adding electronegative substituents near the acidic group increases strength through the same charge-withdrawing effect: chloroacetic acid is much stronger than acetic acid.
Which protons are acidic? Only those bonded to highly electronegative atoms — O, N, or a halogen. In acetic acid, CH₃COOH, only the –OH proton ionizes; the three C–H protons do not.
Explain why (a) HBr is a stronger acid than HCl, (b) H₂SO₄ is a stronger acid than H₂SO₃, and (c) CH₃COOH is a much stronger acid than CH₃CH₂OH.
(a) HBr vs HCl. Bromine is larger than chlorine, so the H–Br bond is longer and weaker than the H–Cl bond and requires less energy to break. The resulting Br⁻ ion is also larger than Cl⁻, so its negative charge is spread over a greater volume, making it more stable. Both effects favor proton loss, so HBr is the stronger acid. (Electronegativity would predict the opposite, but bond strength dominates within a group.)
(b) H₂SO₄ vs H₂SO₃. Sulfuric acid has one more oxygen bonded to the central sulfur than sulfurous acid. That extra highly electronegative oxygen withdraws electron density from the S–O–H region, weakening the O–H bond. More importantly, the conjugate base HSO₄⁻ can delocalize its negative charge over more oxygen atoms through resonance than HSO₃⁻ can. The more stable conjugate base makes H₂SO₄ the stronger acid.
(c) CH₃COOH vs CH₃CH₂OH. When acetic acid loses its proton, the resulting acetate ion has its negative charge delocalized over two equivalent oxygen atoms by resonance, and both C–O bonds become identical with bond order 1.5. When ethanol loses its proton, the ethoxide ion has its full negative charge localized on a single oxygen with no resonance stabilization available. The far more stable acetate ion means the proton leaves acetic acid much more readily — acetic acid’s Ka is roughly 10¹⁰ times larger.