Acids and Bases

Acid-base chemistry underlies most of organic chemistry's reaction mechanisms — from proton transfers to the electron-pushing arrows used to track nucleophilic attack. This article covers the two definitions of acids and bases used in organic chemistry (Brønsted-Lowry and Lewis), how Ka and pKa quantify acid strength, and how acidity compares across common organic functional groups, from weak alkanes to strongly acidic carboxylic acids.

Key Takeaways

• Brønsted-Lowry: acid = proton donor, base = proton acceptor. Forms a conjugate base (from the acid) and a conjugate acid (from the base).

• Lewis: acid = electron pair acceptor, base = electron pair donor — broader than Brønsted-Lowry, includes reactions without proton transfer. Lewis acid-base pairs form a coordinate covalent bond, with both electrons from the Lewis base.

• Ka (acid dissociation constant) measures how much an acid dissociates; larger Ka = stronger acid. pKa = −log(Ka); lower pKa = stronger acid.

• Stronger acids have more stable conjugate bases; the conjugate base of a strong acid is weak, and vice versa.

• Functional-group acidity, weakest to strongest: alkanes → alkenes ≈ H₂ → amines → alkynes ≈ esters → aldehydes/ketones → alcohols ≈ water → carboxylic acids → hydronium ion.

• Resonance stabilization of the conjugate base explains the biggest jumps in acidity: carboxylate ions (carboxylic acids) and enolates (aldehydes/ketones) are both far more stable than a non-resonance-stabilized conjugate base like an alkoxide.

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry definition describes an acid as a species that donates a proton (H⁺) and a base as a species that accepts a proton. This definition focuses on proton transfer and covers many of the acid-base reactions encountered in organic chemistry.

For example, when a Brønsted-Lowry acid, such as HA, donates a proton to water, it forms hydronium (H₃O⁺) and its conjugate base, A⁻. Similarly, when a Brønsted-Lowry base, B, accepts a proton, it forms its conjugate acid, HB⁺.

Lewis Acids and Bases

The Lewis definition describes acids and bases in terms of electron pairs rather than protons. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. This definition broadens the scope of acid-base chemistry beyond proton transfer alone — every Brønsted-Lowry acid-base reaction is also a Lewis acid-base reaction, but the Lewis definition also covers reactions that don't involve a proton at all.

When a Lewis acid reacts with a Lewis base, they form a coordinate covalent bond, where both electrons in the new bond originate from the Lewis base. This type of interaction is especially relevant in organic reactions involving metal catalysts or electrophilic species, and it's the same electron-pair-donation logic used to describe nucleophiles attacking electrophiles.

Ka and pKa: Quantifying Acid Strength

To quantify acid strength, organic chemists use the acid dissociation constant, Ka, which describes the extent to which an acid donates a proton in solution. For a general acid HA dissociating in water, the equilibrium expression is:

Ka = [H₃O⁺][A⁻] / [HA]

A larger Ka value indicates a stronger acid, because more of the acid dissociates into ions, increasing the concentration of H₃O⁺ in solution.

To simplify comparisons, chemists define pKa as the negative logarithm of Ka:

pKa = −log(Ka)

A lower pKa value corresponds to a stronger acid, while a higher pKa value indicates a weaker acid. This relationship is useful for predicting acid-base behavior in organic molecules, particularly in reactions where proton transfer plays a key role.

Acid Strength and Conjugate Base Stability

The pKa of a molecule provides a direct measure of its acidity: lower pKa values indicate stronger acids, and higher pKa values indicate weaker acids. This also means that the conjugate base of a strong acid is weak, whereas the conjugate base of a weak acid is strong — the more readily an acid gives up its proton, the less reactive (more stable) its resulting conjugate base tends to be. Understanding this trend is essential for predicting reaction mechanisms and proton transfer events in organic chemistry.

Comparing Acidity Across Functional Groups

Acidity varies substantially across organic functional groups. Ranked from weakest acid (highest pKa) to strongest acid (lowest pKa):

• Alkanes — pKa ≈ 50. Extremely weak proton donors; their conjugate bases, alkyl carbanions, are highly unstable and strong bases.

• Alkenes — pKa ≈ 43. Still weak acids, but slightly more acidic than alkanes; their conjugate bases (alkene anions) are also highly reactive, strong bases.

• Molecular hydrogen (H₂) — pKa ≈ 42. Also a very weak acid.

• Amines, such as ammonia (NH₃) — pKa ≈ 35. More acidic than alkanes and alkenes; the conjugate base, NH₂⁻ (amide ion), is a strong base frequently used in organic synthesis.

• Alkynes (terminal, R-C≡C-H) — pKa ≈ 25. A significant drop from amines — terminal alkynes are much stronger acids. Their conjugate bases, alkynyl anions (R-C≡C⁻), are still strong bases but more stable than amide ions.

• Esters — pKa ≈ 25, similar acidity to alkynes. Their conjugate bases are resonance-stabilized but still reactive.

• Aldehydes — pKa ≈ 17-20, and ketones — pKa ≈ 20-24. More acidic than esters and alkynes, due to resonance stabilization of the enolate conjugate base.

• Alcohols — pKa ≈ 17. Significantly more acidic than amines or alkanes.

• Water — pKa ≈ 16. Similar acidity to alcohols.

• Carboxylic acids — pKa ≈ 4. Much stronger acids than alcohols and ketones (see below for why).

• Hydronium ion (H₃O⁺) — pKa ≈ −1.7. A very strong acid, meaning its conjugate base, water, is extremely weak.

Why Carboxylic Acids Are So Acidic

Carboxylic acids owe their strong acidity to resonance stabilization of the carboxylate ion (R-COO⁻) formed when they donate a proton. The negative charge is delocalized equally across both oxygen atoms, making the carboxylate ion far more stable than an alkoxide ion (from an alcohol), which has no comparable resonance stabilization. This stabilization is why carboxylic acids readily donate protons in aqueous solution, giving them a pKa around 4 compared to an alcohol's pKa around 17.

The same principle — resonance stabilization of the conjugate base — explains why aldehydes and ketones are more acidic than simple alkanes, alkenes, or alkynes: their conjugate base (an enolate) delocalizes negative charge between carbon and oxygen.

Common MCAT Mistakes

• Treating Lewis acids/bases as a separate category from Brønsted-Lowry ones. Every Brønsted-Lowry acid-base reaction is also a Lewis acid-base reaction — the Lewis definition just extends further, to reactions with no proton transfer at all.

• Reversing the Ka/pKa relationship. A larger Ka means a stronger acid, but because pKa = −log(Ka), a lower pKa means the stronger acid. Mixing these up flips every acidity comparison.

• Assuming a strong acid has a strong conjugate base. The opposite is true: the more readily an acid donates its proton, the more stable (weaker) its conjugate base is.

• Misordering functional-group acidity by ignoring resonance. Carboxylic acids (pKa ≈ 4) are far more acidic than alcohols (pKa ≈ 17) not because of some minor electronic effect, but because the carboxylate conjugate base delocalizes charge across two oxygens by resonance — the same logic makes aldehydes/ketones more acidic than alkanes, alkenes, or alkynes.

MCAT-Style Concept Check

Question: A carboxylic acid (pKa ≈ 4) is a much stronger acid than a structurally similar alcohol (pKa ≈ 17). Which of the following best explains this difference?

• A) The carboxylic acid's O-H bond is longer and physically weaker than the alcohol's O-H bond.

• B) The carboxylate conjugate base delocalizes its negative charge by resonance across two oxygen atoms, while the alkoxide conjugate base has no comparable resonance stabilization.

• C) The carboxylic acid has a greater molecular weight, which directly increases acidity.

• D) The alcohol's conjugate base is aromatic, making it unusually stable and unreactive.

Answer: B

Explanation: Acid strength tracks conjugate base stability. Deprotonating a carboxylic acid gives a carboxylate ion (R-COO⁻) whose negative charge is delocalized equally across both oxygen atoms by resonance, making it far more stable than an alkoxide ion (R-O⁻), which has no resonance stabilization to spread out its charge. This stability difference — not bond length, molecular weight, or aromaticity — is why carboxylic acids (pKa ≈ 4) are so much more acidic than alcohols (pKa ≈ 17).

FAQ

What's the difference between Brønsted-Lowry and Lewis acids and bases?

Brønsted-Lowry acids and bases are defined by proton transfer — an acid donates H⁺, a base accepts it. Lewis acids and bases are defined by electron pairs instead — an acid accepts an electron pair, a base donates one. Every Brønsted-Lowry reaction is also a Lewis acid-base reaction, but the Lewis definition also covers reactions involving no proton at all, such as those with metal catalysts or electrophilic species.

What does pKa actually tell you about an acid?

pKa is the negative logarithm of the acid dissociation constant Ka. A lower pKa means a larger Ka, meaning the acid dissociates more in solution — so lower pKa always corresponds to a stronger acid.

Why are carboxylic acids more acidic than alcohols?

When a carboxylic acid loses a proton, the resulting carboxylate ion delocalizes its negative charge by resonance across two oxygen atoms, which stabilizes it substantially. An alcohol's conjugate base, an alkoxide ion, has no comparable resonance stabilization, so it's much less stable — making the alcohol a much weaker acid (pKa ≈ 17 vs. ≈ 4 for carboxylic acids).

Why are aldehydes and ketones more acidic than alkanes or alkenes?

Deprotonating an aldehyde or ketone at the carbon adjacent to the carbonyl produces an enolate, which delocalizes negative charge between carbon and oxygen by resonance. Alkanes and alkenes have no such resonance-stabilized conjugate base, so their pKa values (≈ 43-50) are far higher than those of aldehydes and ketones (≈ 17-24).