pKa and Predicting Reactions

pKa is the single most useful number in organic chemistry. Given the pKa of two acids, you can predict which way a reaction will run, which hydrogen will be pulled off preferentially, and whether a base is strong enough to do the job you need. Internalize the pKa values of a dozen common functional groups and you will save yourself from memorizing hundreds of individual reactions.

The MCAT will not ask you to calculate a pKa. What it will ask is whether a reaction occurs and which side of the equilibrium is favored. Both answers come from comparing pKa values.

Definition: Ka and pKa

An acid in water establishes an equilibrium:

HA ⇌ H⁺ + A⁻

The acid dissociation constant is:

Ka = [H⁺][A⁻] / [HA]

A bigger Ka means more dissociation, which means a stronger acid. But Ka values span twenty orders of magnitude, so chemists use the logarithm:

pKa = −log Ka

Smaller (or more negative) pKa = stronger acid. Bigger pKa = weaker acid.

pKa Values You Should Memorize

A handful of pKa values covers almost every MCAT question. Anchor these in your head:

Learn the family trends: hydrohalic acids first, then carboxylic acids (4-5), ammoniums (9-10), phenols (10), alcohols (16-18), alpha-carbons (20-25 for one carbonyl, 10-13 for two), amines (38), alkanes (50). You do not need exact values - you need relative order and typical magnitudes.

Why a Stronger Acid Has a Weaker Conjugate Base

Strong acid means the proton dissociates easily. That happens when the conjugate base is stable - it can hold the extra electron pair without trouble. Conversely, a weak acid has a conjugate base that is desperate to reclaim the proton (the conjugate base is unstable as an anion).

Rule to remember: strong acid ↔ weak conjugate base. Weak acid ↔ strong conjugate base.

That is why hydroxide (OH⁻, conjugate base of water at pKa 16) is a stronger base than chloride (Cl⁻, conjugate base of HCl at pKa −7). Water is a weaker acid than HCl, so its conjugate base hangs on to protons more tightly.

Predicting the Direction of an Acid-Base Reaction

Here is the central rule: the equilibrium lies on the side of the weaker acid (the side with the higher pKa).

In other words: a reaction proceeds in the direction that converts a stronger acid into a weaker one (and a stronger base into a weaker one).

Example: does acetic acid (pKa 4.76) react with sodium hydroxide (conjugate acid of hydroxide is water, pKa 15.7)?

• Forward direction: acetic acid (pKa 4.76) → acetate + water. Water is the product acid with pKa 15.7.
• The reaction converts acetic acid (stronger, pKa 4.76) into water (weaker, pKa 15.7).
• Equilibrium lies far to the right - favored by about 10^(15.7 − 4.76) ≈ $10^{11}$.

Yes, the reaction occurs and essentially goes to completion.

Second example: does methanol (pKa 16) react with sodium hydride (NaH - conjugate acid is H₂, pKa ≈ 36)?

• Forward: methanol → methoxide + H₂. H₂ is much weaker (pKa 36) than methanol (pKa 16).
• Equilibrium lies on the weaker-acid side (H₂), so strongly to the right.

Again yes - this is how methoxide is actually generated in a lab.

Third example: does ethanol (pKa 16) deprotonate acetone (alpha-H pKa ≈ 20)?

• Forward: acetone → enolate + ethanol. Ethanol is weaker acid (pKa 16) than acetone (pKa 20). Wait - that is the wrong way around. Let me re-read.
• Actually: if ethanol removes a proton from acetone, the forward direction is ethanol → ethoxide + acetone-enolate. But we started with acetone being the acid here. If ethanol’s O-H proton is what exists in solution, then ethanol is not the base; we need a base like NaH or LDA to remove the alpha-H.

The practical takeaway: if your proposed base is a weaker acid’s conjugate base than your target acid, the reaction is unfavorable. You need a stronger base (higher pKa conjugate acid) than your target acid. This is why deprotonating a ketone alpha-carbon (pKa 20) requires LDA (conjugate acid pKa 36), not hydroxide (pKa 15.7).

The 10-Unit Rule of Thumb

A quick mental estimate: if the two pKa values differ by at least 10, the reaction goes essentially to completion (equilibrium constant ~$10^{10}$ favoring the weaker acid). If they differ by 1-3 units, the reaction is close to 50-50 and may or may not be practically useful.

Example: acetic acid (pKa 4.76) + water (pKa 15.7) → acetate + H₃O⁺. The difference is about 11, so yes - essentially complete in the forward direction. But the reverse (acetate + H₃O⁺ → acetic acid + water) is the direction in a strongly acidic solution, so we flip the lens: at low pH, acetate exists mostly as acetic acid; at high pH, mostly as acetate. This is the Henderson-Hasselbalch logic (see General Chemistry Ch 10).

Ka, pKa, and pH Together

The Henderson-Hasselbalch equation ties these together:

pH = pKa + log([A⁻] / [HA])

When pH = pKa, the protonated and deprotonated forms are in equal amounts. At pH two units below pKa, the acid form dominates roughly 100:1. Two units above, the base form dominates 100:1.

For MCAT biology: the carboxylic acid of an amino acid has pKa around 2. At pH 7 (physiological), the COOH is 5 units above its pKa, so it is roughly 99.999% deprotonated (COO⁻). The amine has pKa around 9. At pH 7, the amine is 2 units below its pKa, so it is roughly 99% protonated (NH₃⁺). That is the zwitterion.

The next section explains WHY these pKa values differ so much - the ARIO framework for ranking acidity from molecular structure alone.