The pKa table is a list of results. This section is the explanation behind it. Given any two hydrogens, you should be able to say which is more acidic and why, without looking anything up — and the reasoning is always the same: the more stable the conjugate base, the stronger the acid. Everything below is a way of asking how stable that anion is.
The ARIO framework
Four structural factors determine anion stability. Check them in this order, because they are roughly in order of size, and the first one that differs usually settles the question.
- A — Atom. Which atom carries the negative charge in the conjugate base?
- R — Resonance. Can the charge be delocalized over more than one atom?
- I — Induction. Are there electronegative atoms nearby pulling density away through sigma bonds?
- O — Orbital. What is the hybridization of the atom holding the charge?
Atom effects are usually the largest and resonance is usually next; induction and orbital effects most often break ties. But the order is a heuristic, not a law — a strongly resonance-stabilized carbanion can beat an unstabilized alkoxide, and the point of naming all four is that you can tell which are in play.
A — Atom: two trends, two different reasons
Across a period, acidity increases with electronegativity: C–H (pKa ≈ 50) < N–H (≈ 38) < O–H (≈ 16) < H–F (3.2). Twenty-two orders of magnitude from nitrogen to oxygen alone. A more electronegative atom is more willing to hold a negative charge, so its anion is more stable and its acid is stronger.
Down a group, acidity increases even though electronegativity decreases: H–F (3.2) < H–Cl (−7) < H–Br (−9) < H–I (−10). Here size wins. A larger atom spreads the charge over a much bigger, more polarizable electron cloud, and that dilution stabilizes the anion far more than the loss of electronegativity costs. The weakening C–H or H–X bond down the group pushes the same direction.
R — Resonance
Delocalizing a negative charge over several atoms is powerfully stabilizing, and resonance is the biggest single lever in the table after the choice of atom.
Compare ethanol (pKa 16) with acetic acid (pKa 4.76). Both give up an O–H proton, and in both cases the charge lands on oxygen — so factor A is identical. The difference is entirely in what happens next: acetate spreads its charge equally over two equivalent oxygens, while ethoxide must keep all of it on one. Eleven pKa units, a hundred billion-fold, from delocalization alone.
Phenol (pKa 10) sits between them, and the reason is instructive. Phenoxide delocalizes its charge into the aromatic ring — but onto carbon atoms, which hold negative charge much less comfortably than oxygen does. Real stabilization, but less of it than acetate gets. Six units more acidic than an alcohol, five units less acidic than a carboxylic acid.
A ketone's alpha C–H has a pKa around 20 — already remarkable for a C–H, because the resulting enolate delocalizes onto the carbonyl oxygen.
Put a second carbonyl on the other side, as in pentane-2,4-dione, and the alpha hydrogen between them drops to pKa 9. The anion now delocalizes onto two oxygens rather than one.
That is roughly as acidic as phenol, and more acidic than ammonium — from a hydrogen on carbon. This is why 1,3-dicarbonyls are the workhorse nucleophiles of Module 11: an ordinary base like ethoxide deprotonates them completely.
I — Induction
An electronegative atom pulls electron density through the sigma framework, helping to spread out a nearby negative charge. The effect is real, additive, and short-ranged.
| Acid | pKa | Effect |
|---|---|---|
| CH₃COOH | 4.76 | baseline |
| ClCH₂COOH | 2.86 | one Cl, adjacent |
| Cl₂CHCOOH | 1.29 | two Cl |
| Cl₃CCOOH | 0.65 | three Cl |
| FCH₂COOH | 2.59 | F beats Cl |
| ClCH₂CH₂CH₂COOH | 4.52 | Cl three carbons away |
Two lessons from that table. Induction is cumulative — each chlorine adds roughly another 1.5 pKa units — and it dies off sharply with distance. Move the chlorine three carbons from the acid group and almost nothing is left of the effect. That distance dependence is the fingerprint that distinguishes induction from resonance, which propagates along a conjugated system with little loss.
O — Orbital
Compare the acidity of three carbon–hydrogen bonds where nothing differs but hybridization: terminal alkyne C–H (sp, pKa ≈ 25), alkene C–H (sp², ≈ 44), alkane C–H (sp³, ≈ 50).
More s-character means the orbital holding the lone pair in the conjugate base is closer to the nucleus and lower in energy, which stabilizes the negative charge. An sp orbital is 50% s; an sp³ orbital is 25%. Twenty-five orders of magnitude of acidity difference from nothing but the mixing ratio of the orbitals.
The same effect explains why protonated pyridine (pKaH 5.2) is far more acidic than a protonated alkylamine (pKaH ≈ 10): pyridine's lone pair sits in an sp² orbital and is held more tightly, making it a weaker base.
Which is more acidic: ethanol or ethanethiol?
A: the charge lands on O in one case and S in the other — same group, so size dominates. Sulfur is bigger and more polarizable, so ethanethiol should be the stronger acid. R: no resonance in either. I: no electronegative substituents in either. O: both sp³.
Only factor A differs, and it points to the thiol. Measured: ethanol 16, ethanethiol 10.6 — the thiol is about 250,000 times more acidic. The framework got it right, and it took one step.
What carries forward
These four factors are the general theory of anion stability, and anion stability is what sets leaving-group ability, nucleophile strength, enolate formation and carbanion chemistry. You will apply them directly to carboxylic acids (Module 10), to alpha hydrogens (Module 11), to amine basicity (Module 12), and to the substituent effects that control aromatic substitution (Module 13). Very little in the second half of the course is new physics; most of it is this section applied to a new structure.