pKa and acidity trends

Key structures for this topic — drawn live.

What pKa actually measures

The acid dissociation constant K_a measures how far an acid ionizes in water. Because K_a values span enormous ranges, we use pK_a = −log K_a to compress them onto a workable scale. The key habit to build: the lower the pK_a, the stronger the acid. A difference of one pK_a unit is a factor of ten in K_a, so small numeric gaps are large chemical differences.

Rough landmarks worth memorizing: strong mineral acids like HCl are around −7; carboxylic acids sit near 4–5; water is 15.7; alcohols around 16–18; terminal alkynes near 25; and typical C–H bonds are 45–50 (essentially not acidic). These anchors let you predict whether one species can deprotonate another.

The golden rule: it's about the conjugate base

Acidity is really a question about the conjugate base. When an acid H–A donates its proton, it becomes A⁻. Anything that stabilizes the resulting negative charge on A⁻ makes the acid stronger, because a stable conjugate base means the acid is happy to let the proton go. So to compare two acids, deprotonate each one on paper and ask which anion is more stable.

The four factors below all reduce to the same idea: how well is that negative charge stabilized? A common memory aid is ARIO — Atom, Resonance, Induction, Orbital. Apply them in order, and usually the first one that differs decides the outcome.

Factor 1: the atom bearing the charge

Which element holds the negative charge matters most. Two trends of the periodic table apply:

• Across a period (left to right), electronegativity rises, so a more electronegative atom holds the negative charge better. Acidity increases: CH₄ < NH₃ < H₂O < HF.
• Down a group, size dominates over electronegativity — a bigger atom spreads the charge over a larger volume, stabilizing it. Acidity increases down the halogens: HF < HCl < HBr < HI.

Size beats electronegativity down a group, which is why HI (large, diffuse I⁻) is a far stronger acid than HF even though fluorine is more electronegative.

Factor 2: resonance

If the conjugate base can delocalize its negative charge by resonance, it is much more stable, and the acid is much stronger. The classic case is a carboxylic acid (pK_a ≈ 4–5) versus an alcohol (pK_a ≈ 16). Both deprotonate at an O–H, but the carboxylate anion spreads its charge over two equivalent oxygens, while an alkoxide has the charge stuck on one oxygen. That resonance stabilization is worth roughly 11 pK_a units — a factor of 10¹¹.

Factor 3: induction

Electron-withdrawing groups pull electron density through the sigma bonds (the inductive effect), helping to spread out a nearby negative charge. Adding electronegative atoms near the acidic site lowers pK_a:

• Acetic acid (CH₃COOH) has pK_a ≈ 4.76; trichloroacetic acid (CCl₃COOH) drops to ≈ 0.7 because three electronegative chlorines stabilize the carboxylate.
• Induction weakens with distance — a chlorine on the carbon next to the acid group helps much more than one several carbons away.

Factor 4: orbital hybridization

The hybridization of the atom holding the lone pair matters. More s character holds electrons closer to the nucleus and stabilizes negative charge, so acidity follows sp > sp² > sp³. This is why a terminal alkyne C–H (sp carbon, pK_a ≈ 25) is dramatically more acidic than an alkene C–H (sp², ≈ 44) or an alkane C–H (sp³, ≈ 50). The sp carbanion is stabilized by its 50% s character.

Putting it together

To rank acids, deprotonate each, then walk down ARIO until you find the first factor that distinguishes the conjugate bases. Usually that factor decides the winner. For example, comparing ethanol (pK_a ≈ 16), acetic acid (≈ 4.8), and acetylene (≈ 25): acetic acid wins on resonance, ethanol's alkoxide has no resonance but sits on oxygen, and acetylene relies only on hybridization with charge on a less electronegative carbon — giving the order acetic acid > ethanol > acetylene in acid strength.

Once these four factors are second nature, you can predict which base will deprotonate which acid, and therefore which direction an acid–base equilibrium favors — the reaction always runs toward the weaker (more stable) acid and base.