Appendix B: pKa Tables
Chapter 3 makes a deliberate choice not to lead with numbers: the goal there is to understand why one acid is stronger than another by evaluating the stability of the conjugate base. This appendix is the reference that goes with that framework — a set of representative pKa values, organized by compound class, to consult once the underlying reasoning is in place.
A caution before using this appendix: the Common Mistake named in Chapter 3 is memorizing pKa values in place of understanding stability. Use this table to check reasoning and to calibrate intuition (is this acid stronger or weaker than water? than acetic acid?) — not as a list to memorize outright. Exact values also vary somewhat by source and solvent; what matters is the relative order and the structural reasoning behind it.
How to Read pKa
pKa = −log(Ka). A lower pKa means a stronger acid — it dissociates more completely.
Each unit of pKa represents a factor of 10 in acid strength. An acid with pKa 5 is 10 times stronger than one with pKa 6, and 100,000 times stronger than one with pKa 10.
Basicity is the mirror image. The weaker the conjugate acid (higher its pKa, in the pKaH notation used for amines below), the stronger the base. There is no separate “pKb table” in this appendix — every base’s strength is read from the pKa of its conjugate acid.
A functional group is a moderate acid, base, both, or neither, depending on which factor from Chapter 3 applies to it — resonance, electronegativity, induction, or hybridization. The tables below are grouped by compound class so those factors can be compared side by side.
Master Table — Representative pKa Values
Ordered from strongest acid to weakest.
| Compound | Acidic Site | pKa (approx.) | Compound Class |
|---|---|---|---|
| HCl | H–Cl | −7 | Mineral acid (reference point) |
| H₃O⁺ | H–OH₂⁺ | −1.7 | Hydronium (reference point) |
| Trifluoroacetic acid | CF₃COOH | 0.2 | Carboxylic acid (heavily inductive) |
| Trichloroacetic acid | CCl₃COOH | 0.7 | Carboxylic acid (heavily inductive) |
| Dichloroacetic acid | Cl₂CHCOOH | 1.3 | Carboxylic acid |
| Chloroacetic acid | ClCH₂COOH | 2.9 | Carboxylic acid |
| Formic acid | HCOOH | 3.8 | Carboxylic acid |
| Benzoic acid | PhCOOH | 4.2 | Carboxylic acid |
| Anilinium ion | PhNH₃⁺ | 4.6 | Ammonium (conjugate acid of a weak amine) |
| Acetic acid | CH₃COOH | 4.8 | Carboxylic acid |
| Pyridinium ion | C₅H₅NH⁺ | 5.2 | Ammonium (aromatic amine) |
| p-Nitrophenol | 4-O₂N–C₆H₄OH | 7.2 | Phenol (electron-poor ring) |
| Acetylacetone (pentane-2,4-dione) | –CH₂– between two C=O | ~9 | 1,3-Dicarbonyl (doubly activated α-C–H) |
| Ammonium ion | NH₄⁺ | 9.25 | Ammonium |
| Phenol | C₆H₅OH | 10.0 | Phenol |
| Methylammonium ion | CH₃NH₃⁺ | 10.6 | Ammonium (alkyl amine) |
| Diethyl malonate | –CH₂– between two esters | ~13 | 1,3-Dicarbonyl (doubly activated α-C–H) |
| Water | H₂O | 15.7 | Reference solvent |
| Ethanol | CH₃CH₂OH | 16 | Alcohol |
| tert-Butanol | (CH₃)₃COH | 18 | Alcohol (more hindered/less stabilized) |
| Acetone | CH₃–CO–CH₃, α-C–H | 20 | Ketone α-carbon |
| Terminal alkyne | HC≡CH, C–H | 25 | Hydrocarbon (sp C–H) |
| Ethyl acetate | CH₃CO–OEt, α-C–H | 25 | Ester α-carbon |
| Ammonia | NH₃ | 38 | Reference (as an acid, not a base) |
| Ethylene | H₂C=CH₂, C–H | 44 | Hydrocarbon (sp² C–H) |
| Ethane | CH₃CH₃, C–H | 50 | Hydrocarbon (sp³ C–H) |
Carboxylic Acids
| Compound | pKa | Structural Note |
|---|---|---|
| Trifluoroacetic acid | 0.2 | Three highly electronegative F atoms pull electron density from the carboxylate |
| Trichloroacetic acid | 0.7 | Three inductively withdrawing Cl atoms |
| Chloroacetic acid | 2.9 | One inductively withdrawing Cl atom |
| Formic acid | 3.8 | No alkyl group donating electron density |
| Benzoic acid | 4.2 | Aromatic ring is mildly electron-withdrawing relative to alkyl |
| Acetic acid | 4.8 | Simple alkyl carboxylic acid — the typical reference value |
Why carboxylic acids are acidic at all: deprotonation gives a carboxylate ion whose negative charge is delocalized equally over both oxygens by resonance. This is the same reasoning developed in Chapter 3’s Gentle Exercise (“why are carboxylic acids stronger acids than alcohols?”) — an alkoxide has nowhere to delocalize its charge, while a carboxylate does.
Why the halogenated acids are so much stronger: induction, not resonance. Electronegative halogens near the carboxyl group pull electron density through the sigma bonds, stabilizing the negative charge on the carboxylate. This effect drops off quickly with distance — a halogen on the carbon adjacent to the carboxyl (α) has a much larger effect than one further away.
Phenols
| Compound | pKa | Structural Note |
|---|---|---|
| p-Nitrophenol | 7.2 | The nitro group is strongly electron-withdrawing and conjugated directly into the ring |
| Phenol | 10.0 | Reference value |
Why phenols are more acidic than alcohols (compare phenol, pKa 10, to ethanol, pKa 16): the phenoxide ion delocalizes its negative charge into the aromatic ring by resonance, spreading it onto the ortho and para carbons. An alkoxide has no comparable delocalization pathway. This is the same resonance argument used for carboxylate, applied to a different scaffold — see Appendix A’s entry on phenols and Chapter 17 for how ring substituents shift this value further.
Why electron-withdrawing ring substituents increase acidity further: a para-nitro group can accept negative charge directly by resonance (the charge can be drawn all the way onto the nitro oxygens), which is why p-nitrophenol is nearly 1,000 times more acidic than phenol itself. Electron-donating substituents (e.g., para-methoxy) have the opposite effect.
Alcohols and Water
| Compound | pKa | Structural Note |
|---|---|---|
| Water | 15.7 | Reference solvent |
| Ethanol | 16 | Typical primary/secondary alcohol |
| tert-Butanol | 18 | Bulkier alkoxide is less well solvated/stabilized |
Alcohols and water are only weakly acidic — their conjugate bases (alkoxide, hydroxide) carry a full negative charge with no resonance delocalization. Alkoxide is nonetheless a strong enough base to be useful synthetically (e.g., in E2 eliminations, Chapter 8), precisely because it is not resonance-stabilized.
Ammonium Ions and Amine Basicity
Amines themselves are not usually discussed by “pKa” directly — instead, chemists report the pKa of the conjugate acid (the ammonium ion), often written pKaH. A higher pKaH means a more basic amine, because its conjugate acid holds onto the proton more tightly.
| Amine | Conjugate Acid | pKaH | Basicity Note |
|---|---|---|---|
| Aniline | Anilinium ion | 4.6 | Weakly basic — nitrogen lone pair is delocalized into the aromatic ring, less available to accept a proton |
| Pyridine | Pyridinium ion | 5.2 | Weakly basic — lone pair sits in an sp² orbital in the ring plane, not part of the aromatic system, but still less basic than an alkyl amine |
| Ammonia | Ammonium ion | 9.25 | Reference value |
| Methylamine | Methylammonium ion | 10.6 | More basic than ammonia — the alkyl group donates electron density to nitrogen |
Why aromatic amines are weaker bases than alkyl amines: in aniline, the nitrogen lone pair is conjugated into the ring (the same resonance that makes aniline nitrogen a weaker nucleophile — see Appendix A’s note on amide resonance for the analogous effect on amides, which is far more extreme). That lone pair is less available to accept a proton, so aniline is a much weaker base than cyclohexylamine, its non-aromatic analog.
C–H Acids and Hybridization
| Compound | pKa | Hybridization at Carbon |
|---|---|---|
| Ethane | 50 | sp³ |
| Ethylene | 44 | sp² |
| Terminal alkyne (acetylene) | 25 | sp |
Why hybridization affects C–H acidity: an sp orbital has more s-character (50%) than sp² (33%) or sp³ (25%). Electrons in an orbital with more s-character sit closer to the nucleus, so the conjugate base (a carbanion) is more stable when the negative charge sits in an sp orbital. This is why a terminal alkyne C–H (Chapter 9) is acidic enough to be deprotonated by strong bases such as sodium amide, while an alkane or alkene C–H is not.
Carbonyl α-Carbons and Enolates
| Compound | pKa | Structural Note |
|---|---|---|
| Ethyl acetate (ester α-C–H) | 25 | One carbonyl stabilizes the enolate by resonance |
| Acetone (ketone α-C–H) | 20 | One carbonyl stabilizes the enolate by resonance; ketones are somewhat more acidic than esters at the α-position |
| Diethyl malonate (1,3-diester α-C–H) | ~13 | Two carbonyls both stabilize the same enolate — resonance delocalizes the negative charge onto two separate carbonyl oxygens |
| Acetylacetone (1,3-diketone α-C–H) | ~9 | Two carbonyls, same doubly-stabilized effect, and ketone carbonyls stabilize slightly better than ester carbonyls |
Why α-carbons are acidic at all: deprotonating a C–H adjacent to a carbonyl generates an enolate, whose negative charge is resonance-delocalized onto the carbonyl oxygen rather than sitting on carbon alone (Chapter 14). This is the same stability principle as carboxylate and phenoxide — resonance delocalization of the conjugate base — applied to a carbanion instead of an alkoxide.
Why 1,3-dicarbonyls are dramatically more acidic: with two carbonyls flanking the same C–H, the single resulting enolate can delocalize onto either carbonyl oxygen, effectively doubling the resonance stabilization. This is why diethyl malonate and acetylacetone are common starting materials for enolate alkylation chemistry (Chapter 14) — their α-protons are acidic enough to be removed by moderate bases like alkoxide, rather than requiring the very strong bases needed for a simple ketone or ester.
The Four Stability Factors, Revisited
Chapter 3 introduces four factors that stabilize a conjugate base, sometimes remembered by the mnemonic ARIO — Atom, Resonance, Induction, Orbital (hybridization). Every trend in this appendix is an example of one of these four:
| Factor | What It Means | Example From This Appendix |
|---|---|---|
| Atom | More electronegative atoms (further right on the periodic table) or larger atoms (further down the periodic table) stabilize negative charge better — electronegativity dominates across a period, size dominates down a group | Underlies why O–H is more acidic than N–H (electronegativity); not heavily featured in this course’s scope otherwise |
| Resonance | Delocalizing charge across multiple atoms lowers its energy | Carboxylic acids, phenols, and enolates are all far more acidic than their non-delocalized counterparts (alcohols, alkanes) |
| Induction | Nearby electronegative atoms withdraw electron density through sigma bonds, stabilizing adjacent charge | Trifluoroacetic acid vs. acetic acid; chloroacetic acid vs. acetic acid |
| Orbital (hybridization) | More s-character holds electrons closer to the nucleus, stabilizing negative charge | Terminal alkyne C–H (sp) vs. alkene C–H (sp²) vs. alkane C–H (sp³) |
When comparing two acids, working through these four factors in order — is one atom more electronegative, can the charge resonate, is there an inductive effect nearby, does hybridization differ — will explain nearly every pKa trend in this table without needing to memorize the numbers directly.
Cross-References
- Chapter 3 (Acids and Bases) — the conceptual framework (stability of the conjugate base) that this appendix supplies values for.
- Chapter 8 (Elimination) — alkoxide and other strong, non-resonance-stabilized bases as the reagents that drive E2 elimination.
- Chapter 9 (Addition) — terminal alkyne acidity and its use in alkylation chemistry.
- Chapter 14 (Enols and Enolates) — the mechanism behind α-carbon acidity and enolate resonance in full detail.
- Chapter 17 (Substituent Effects) — how ring substituents shift phenol and aniline acid/base strength.
- Appendix A (Functional Group Atlas) — polarity and reactivity notes for each compound class referenced here.