pKa and Acid–Base Chemistry

The question this page answers: how do we evaluate the acidity and basicity of molecules — and predict which side of an acid–base reaction is favored?

Deeper reading: Clayden, Organic Chemistry 2e, ch. 8 (pp. 163–181) — see our chapter-by-chapter practice map for Clayden.

Acids, bases, and conjugate pairs

What happens when an acid meets a base?

In a Brønsted–Lowry acid–base reaction, an acid and a base exchange a proton (H⁺), forming a conjugate base and a conjugate acid.

Bronsted-Lowry acid-base reaction between water (acid, HA) and ammonia (base, B) giving hydroxide (conjugate base, A-) and ammonium (conjugate acid, HB+), with the transferred proton highlighted and electron pairs shown converting between sigma bond and lone pair — Water (the acid) transfers the highlighted proton to ammonia (the base). Watch the electron pairs: a σ bond becomes a lone pair on the conjugate base, and a lone pair becomes a σ bond in the conjugate acid.

Organic chemists care about acid–base reactions for two reasons: trends in acid and base strength correlate with trends in reactivity more broadly, and many organic reactions contain acid–base steps inside their mechanisms.

Ka and pKa: putting a number on acidity

What does a pKa value actually measure — and why not just use pH?

pKa compresses the enormous range of measured acidities onto a convenient log scale: pKa = −log₁₀ Ka. Stronger acids have lower pKa values.

Ka is the equilibrium constant for an acid's dissociation, Ka = [A⁻][H⁺]/[HA]. Because measured Ka values span dozens of orders of magnitude, it is far more convenient to compare acids by exponent than by absolute value — that is all pKa is.

Two consequences worth internalizing:

A difference of 1 pKa unit is a factor of 10 in acidity. Given pKa(HCl) = −7 and pKa(acetic acid) = 5, HCl is 10¹² times the stronger acid.
You may be used to pH from general chemistry — but pH describes a solution. Organic chemists work in many solvents besides water, so we characterize the molecule using pKa instead.

Predicting equilibria with pKa

Which side of an acid–base equilibrium is favored?

Acid–base reactions favor the side with the weaker acid (the higher pKa).

Acetic acid (acid) reacting with ethoxide (base) to give acetate (conjugate base) and ethanol (conjugate acid) — Acetic acid + ethoxide ⇌ acetate + ethanol. Which side wins?

Acetic acid has pKa about 5; ethanol has pKa about 16 — The acid on the left (pKa ≈ 5) is much stronger than the conjugate acid on the right (pKa ≈ 16), so the equilibrium lies far to the right — by about 11 orders of magnitude.

The six factors that control acidity

What structural factors make an acid stronger?

Stronger acids have more stable conjugate bases. Every pKa trend below is really a statement about how well the lone pair in the conjugate base is stabilized.

Here they are in rough order of importance. In each figure, the pKa of the bolded blue H is shown in blue.

1. Formal charge

Acids with formal (or partial) positive charge are more acidic:

Water has pKa 15.7; protonated water (hydronium) has pKa -1.7

2. Element effects

H's on more electronegative atoms are more acidic (across a row of the periodic table):

pKa values across a period: HF 3.2, H2O 15.7, NH3 38, CH4 50

H's on more polarizable atoms are more acidic (down a column):

pKa values down a group: HF 3.2, HCl -8, HBr -9, HI -10

3. Hybridization

Hybrid orbitals with more s character hold their electrons closer to the nucleus — they behave as if more electronegative:

pKa values by hybridization: ethane (sp3) 50, ethylene (sp2) 45, acetylene (sp) 24

4. Resonance

Delocalization can dramatically stabilize the conjugate base:

pKa values showing resonance stabilization: methane 50, toluene benzylic H 41, indene 20

5. Inductive effects

Electronegative atoms two to three bonds from the acidic site stabilize the conjugate base through σ bonds:

pKa values of acetic acid 4.8, fluoroacetic acid 2.6, difluoroacetic acid 1.3, trifluoroacetic acid -0.3

6. Solvation

Ions like to be solvated, so steric bulk that blocks solvation makes the conjugate base less stable — and the acid weaker:

pKa values of alcohols increasing with steric bulk: methanol 15.5, ethanol 15.9, isopropanol 17.1, tert-butanol 18

Lewis acids and bases

What if no proton is transferred — what are Lewis acids and bases?

Lewis acid–base reactions are the broader category: any lone pair (Lewis base) interacting with an empty orbital (Lewis acid).

Lewis acid-base reaction between acetone (lone pair donor) and lithium cation, forming a new sigma bond from the oxygen lone pair; Li+ behaves like a big H+ — A carbonyl lone pair (Lewis base) binds Li⁺ (Lewis acid). You can think of Li⁺ as behaving like a big H⁺.

Organic chemists use Lewis acids to accentuate positive charge in a molecule and to stabilize species through bond formation. Because both electrons of the new bond come from one atom, this bond is sometimes drawn with special notation: