The Strengths of Acids and Bases
A Complete Guide for Competitive Exams
Concepts, Trends, Structures and Exam Strategies — Explained from First Principles
Acid–Base Definitions: From Arrhenius to Lewis
Three major frameworks describe acid–base behaviour, and competitive exams test all three. Understanding why each exists — and its limitations — is far more powerful than memorising definitions.
1.1 Arrhenius Definition (Limited but Foundational)
- Acid: A substance that releases H⁺ (hydrogen ions) in aqueous solution. Example: HCl → H⁺ + Cl⁻
- Base: A substance that releases OH⁻ ions in aqueous solution. Example: NaOH → Na⁺ + OH⁻
- Limitation: Only works in water; cannot explain NH₃ as a base (no OH⁻ released).
1.2 Brønsted–Lowry Definition (Most Used in Organic Chemistry)
Brønsted defined an acid as a proton donor and a base as a proton acceptor. This elegantly explains the concept of conjugate pairs.
Fig 1.1 — First ionisation of H₂SO₄ showing conjugate acid–base pairs (Brønsted–Lowry)
1.3 Lewis Definition (Most General)
- Lewis Acid: Electron-pair acceptor. Examples: BF₃, AlCl₃, SnCl₄, ZnCl₂, FeCl₃
- Lewis Base: Electron-pair donor. Examples: NH₃, amines, ethers, water
Fig 1.2 — Lewis acid–base adduct formation: BF₃ (acid) + N(CH₃)₃ (base)
Lewis acids must have empty orbitals to accept electrons. BF₃ (empty 2p on B), AlCl₃ (empty 3p on Al), carbocations (empty orbital), transition metal ions — all are Lewis acids. NH₃, H₂O, ROH, R₂S, CN⁻ are classic Lewis bases. This comes up in Friedel–Crafts reactions and coordination chemistry questions.
Understanding pKₐ — The Master Acidity Scale
For an acid HA dissociating in water, the equilibrium is:
Kₐ = [H₃O⁺][A⁻] / [HA]
pKₐ = –log₁₀(Kₐ)
Relation to free energy: –ΔG° = 2.303 RT log Kₐ
Also: ΔG° = ΔH° – TΔS°
The Golden Rule of pKₐ
- Lower pKₐ = Stronger Acid (more dissociated, larger Kₐ)
- Higher pKₐ = Weaker Acid
- Acids with pKₐ > ~16 cannot be detected as acids in water (they produce less H₃O⁺ than water's own autolysis)
- Very strong acids (HCl, HNO₃, HClO₄) all appear equally strong in water — this is the levelling effect of water
Fig 2.1 — Relative pKₐ scale for common organic and inorganic acids
Many students assume Ka depends only on bond energy (ΔH°). But the data reveals something remarkable: for ethanoic acid, ΔH° ≈ –0.5 kJ/mol (almost zero!) while TΔS° = –27.6 kJ/mol dominates ΔG°. The real driver of acid strength differences in carboxylic acids and haloacids is entropy of solvation, not bond enthalpy. This is a favourite IIT-JAM concept.
Why Organic Compounds Are Acidic — Four Governing Factors
The acidity of any organic compound HA is controlled by four factors working together. Understanding each separately — then in combination — is the key to predicting relative acidities without memorisation.
- Bond strength of H–A: Rarely the deciding factor among organic compounds of the same family
- Electronegativity of A: Higher electronegativity → stronger acid (better at retaining the negative charge after proton loss)
- Stabilisation of conjugate base A⁻ vs HA: This is the MOST IMPORTANT factor in organic chemistry
- Nature of solvent: Solvation can override structural factors entirely
Resonance Stabilisation of Carboxylate Anion — The Key Principle
Formic acid (methanoic acid, pKₐ = 3.77) is a relatively strong organic acid because its conjugate base, the formate ion, is stabilised by complete charge delocalisation over two equivalent oxygen atoms.
Fig 3.1 — Resonance stabilisation of formate (HCOO⁻) anion. Both canonical structures are equivalent, giving maximum delocalisation and hence maximum stability.
Why Alcohols Are Far Less Acidic Than Carboxylic Acids
In alkoxide ions (RO⁻), the negative charge sits entirely on one oxygen atom — no delocalisation is possible. So the alkoxide is not significantly more stable than the parent alcohol, and the equilibrium strongly disfavours ionisation.
Students often think "electronegativity of O explains carboxylic acid strength." Wrong order of importance! The resonance stabilisation of the carboxylate anion is the dominant effect. Both carboxylic acids and phenols have an O–H bond, but carboxylic acids are ~10,000× stronger because their anion has two equivalent O atoms sharing the charge, while phenoxide distributes charge onto less electronegative carbon atoms of the ring.
The Role of the Solvent — Water as the Ionising Medium
Despite all the structural reasoning above, the solvent is often the dominant factor determining actual acid strength. Water is uniquely effective because of two properties:
- High dielectric constant (ε = 80): Reduces the electrostatic energy holding ion pairs together — ions stabilise more easily in high-ε solvents
- Ion-solvating ability: Water molecules wrap around ions, delocalising charge via hydrogen bonding (especially powerful for anions)
Hydrogen-Bonded Solvation of H⁺
Fig 4.1 — Solvation shell around H₃O⁺ via hydrogen bonding with surrounding water molecules
When a very strong acid dissolves in water, it ionises completely regardless of its intrinsic strength. HCl, HBr, HI, HNO₃, HClO₄ all appear equally strong in water. To distinguish their true strengths, one must use a weaker base solvent (like glacial acetic acid) where they don't fully ionise.
Simple Aliphatic Acids — Alkyl Groups and Acidity
Replacing the H in formic acid by an alkyl group introduces an electron-donating inductive effect (+I effect), which pushes electrons toward the carboxylate, destabilising the anion and reducing acidity.
| Acid | Structure | pKₐ | Key Feature |
|---|---|---|---|
| Methanoic (Formic) | HCOOH | 3.77 | Reference — strongest simple acid |
| Ethanoic (Acetic) | CH₃COOH | 4.76 | –CH₃ donates electrons (+I) |
| Propanoic | CH₃CH₂COOH | 4.88 | Slightly weaker — longer chain |
| 2-Methylpropanoic | (CH₃)₂CHCOOH | 4.86 | Branching, minor effect |
| 2,2-Dimethylpropanoic | (CH₃)₃CCOOH | 5.05 | Bulky, but mainly steric/solvation |
| Propenoic (Acrylic) | CH₂=CHCOOH | 4.25 | sp² carbon — weaker +I than sp³ |
| Propynoic (Propiolic) | HC≡CCOOH | 1.84 | sp carbon — strong –I, very acidic |
Why Does sp Hybridisation Increase Acidity So Dramatically?
The hybridisation state of the carbon adjacent to –COOH determines how strongly it pulls electrons away from (or donates electrons to) the carboxyl group.
Fig 5.1 — Increasing s-character in hybrid orbitals causes stronger electron withdrawal (–I effect), increasing acidity. sp has 50% s-character vs 25% for sp³.
Substituted Aliphatic Acids — The Power of Electron-Withdrawing Groups
Introducing electron-withdrawing groups (EWG) — especially halogens — near the carboxyl dramatically increases acidity by stabilising the carboxylate anion through spreading negative charge.
Halogen Substituted Acetic Acids
| Compound | pKₐ | Relative Strength vs CH₃COOH |
|---|---|---|
| CH₃COOH (Acetic acid) | 4.76 | Reference (1×) |
| FCH₂COOH (Fluoroacetic) | 2.57 | ~150× stronger |
| ClCH₂COOH (Chloroacetic) | 2.86 | ~80× stronger |
| BrCH₂COOH (Bromoacetic) | 2.90 | ~70× stronger |
| ICH₂COOH (Iodoacetic) | 3.16 | ~40× stronger |
| Cl₂CHCOOH (Dichloroacetic) | 1.25 | ~3200× stronger |
| Cl₃CCOOH (Trichloroacetic) | 0.65 | ~13,000× stronger |
| CF₃COOH (Trifluoroacetic) | 0.23 | ~340,000× stronger |
Why Does the Effect Diminish with Distance?
The inductive effect operates through bonds and decreases rapidly (attenuates) as the number of intervening carbon atoms increases.
Fig 6.1 — Inductive effect of Cl diminishes rapidly with distance from the –COOH group
For α-halo acids: F > Cl > Br > I in terms of acid-strengthening (matches electronegativity order). For same halogen, more halogens = much stronger acid. For position effect: α > β > γ ≫ unsubstituted. These two rules together can answer most ordering questions in 10 seconds.
Phenols — Where Resonance and Induction Battle
Phenol (pKₐ = 9.95) sits between alcohols (pKₐ ~16) and carboxylic acids (pKₐ ~4–5), because the phenoxide anion is partially stabilised by delocalisation into the ring — but less effectively than carboxylate.
Effect of Substituents on Phenol Acidity
Electron-withdrawing groups (–NO₂) increase acidity; electron-donating groups (–CH₃) decrease acidity slightly. Position matters enormously for mesomeric effects.
| Phenol | pKₐ | Reason |
|---|---|---|
| Phenol (C₆H₅OH) | 9.95 | Reference |
| o-Nitrophenol | 7.23 | –I and –M (o-position) both withdraw electrons |
| m-Nitrophenol | 8.35 | Only –I effect; no mesomeric effect at meta |
| p-Nitrophenol | 7.14 | Both –I and –M; strongest withdrawing at para |
| 2,4-(NO₂)₂C₆H₃OH | 4.01 | Two nitro groups — very large effect |
| 2,4,6-(NO₂)₃C₆H₂OH (Picric acid) | 1.02 | Mineral acid strength! |
| o-Methylphenol | 10.28 | +I effect — slight decrease in acidity |
| p-Methylphenol | 10.19 | +I effect; +M actually destabilises anion |
For nitrophenols: p-NO₂ < o-NO₂ < m-NO₂ in terms of pKₐ (both o and p are more acidic than m because they get both inductive AND mesomeric electron withdrawal). Note: o-nitrophenol can form intramolecular H-bonds, making it less soluble in water but stronger as an acid. For o compounds always look for intramolecular H-bonding!
Aromatic Carboxylic Acids — Benzoic Acid and Derivatives
Benzoic acid (pKₐ = 4.20) is stronger than cyclohexane carboxylic acid (pKₐ = 4.87), confirming that the sp² carbon of the ring is less electron-donating than a saturated sp³ carbon. The same logic as acrylic > propionic acid (Section 5).
Effect of Substituents on Benzoic Acid
| Compound | pKₐ | Effect |
|---|---|---|
| Benzoic acid (C₆H₅COOH) | 4.20 | Reference |
| o-Nitrobenzoic acid | 2.17 | Very strong: –I (short distance) + direct interaction |
| m-Nitrobenzoic acid | 3.45 | –I only at meta |
| p-Nitrobenzoic acid | 3.43 | –I + –M at para |
| p-Methoxybenzoic acid | 4.47 | +M (OMe donates at para) dominates –I |
| p-Hydroxybenzoic acid | 4.58 | +M from OH weakens acid — weaker than benzoic! |
| o-Hydroxybenzoic acid (Salicylic) | 2.98 | Intramolecular H-bond stabilises anion |
| 2,6-Dihydroxybenzoic acid | 1.30 | H-bonding from both ortho-OH groups |
o-Hydroxybenzoic acid (salicylic acid) is much more acidic than m- or p-isomers because, after ionisation, the carboxylate anion is stabilised by intramolecular hydrogen bonding with the adjacent –OH group. The m- and p-isomers cannot do this. This is a classic JEE question on intramolecular vs intermolecular H-bonding effects.
Dicarboxylic Acids — A Tale of Two pKₐ Values
Every dicarboxylic acid has two ionisation steps: pKₐ₁ (first COOH loss) and pKₐ₂ (second COOH loss). The second step is always harder because you are removing H⁺ from an already-negative species.
| Acid | pKₐ₁ | pKₐ₂ | Reason for large difference |
|---|---|---|---|
| Ethanedioic (Oxalic) HOOC-COOH | 1.23 | 4.19 | Very close COOH groups: strong EWG mutual effect |
| Propanedioic (Malonic) HOOCCH₂COOH | 2.83 | 5.69 | 1 carbon between — still strong |
| Butanedioic (Succinic) HOOC(CH₂)₂COOH | 4.19 | 5.48 | 2 carbons — inductive effect reduced |
| Maleic acid (cis) | 1.92 | 6.23 | Intramolecular H-bond stabilises cis mono-anion |
| Fumaric acid (trans) | 3.02 | 4.38 | No intramolecular H-bond possible |
Maleic vs Fumaric — Geometric Isomers, Radically Different Acidity
Fig 9.1 — Maleic acid (cis) mono-anion is stabilised by intramolecular H-bonding, making it far more acidic (pKₐ₁ = 1.92) than fumaric acid (trans, pKₐ₁ = 3.02)
Bases — The pKₐ of the Conjugate Acid
Modern practice measures base strength using the pKₐ of the conjugate acid (BH⁺), rather than the older pKb. This creates one unified scale for both acids and bases.
Kₐ(BH⁺) = [B:][H₃O⁺] / [BH⁺]
Rule: Larger pKₐ(BH⁺) = Stronger base (BH⁺ less willing to lose H⁺ → B: more willing to hold it)
pKₐ(HA) + pKₐ(BH⁺) at 25°C: a stronger conjugate acid (lower pKₐ of BH⁺) means a weaker base. Remember: pKa of NH₄⁺ = 9.25, so NH₃ is the reference aliphatic base.
Aliphatic Amines — The Solvation Paradox
You might predict basicity order: NH₃ < RNH₂ < R₂NH < R₃N (more alkyl groups → more electron density on N → stronger base). Reality in water is more complex due to solvation.
Observed pKₐ Values (Conjugate Acid)
The tertiary amine drops back! The two competing effects are:
- Inductive effect (+I): More alkyl groups → more electrons on N → better proton acceptor → increases basicity
- Solvation effect: More alkyl groups → fewer N–H bonds in cation BH⁺ → weaker hydrogen bonding with water → less stabilisation of cation → decreases basicity
Fig 11.1 — Decreasing H-bond solvation of ammonium ions as N–H bonds are replaced by N–C bonds
Guanidine HN=C(NH₂)₂ has pKₐ ≈ 13.6 — one of the strongest organic bases. Why? On protonation, the resulting cation H₂N⁺=C(NH₂)₂ has its positive charge symmetrically delocalised over three equivalent nitrogen atoms via three equivalent resonance structures. This exceptional stabilisation of the cation makes protonation highly favourable. Compare: the neutral molecule's resonance involves charge separation, so it's less stabilised.
Amides — Extremely Weak Bases
Amides (RCONH₂) have pKₐ ≈ 0.5 — vastly weaker than amines (pKₐ ~10). The C=O group withdraws the N lone pair via resonance (–M effect), making it unavailable for protonation.
The N lone pair is tied up in π system → NOT available for H⁺
Aromatic Bases — Aniline and the Resonance Trap
Aniline (pKₐ = 4.62) is a dramatically weaker base than cyclohexylamine (pKₐ = 10.68) or ammonia (pKₐ = 9.25). The reason is elegant and fundamental.
In aniline, the nitrogen lone pair is delocalised into the benzene ring via resonance. This stabilises the free amine. When aniline is protonated, this delocalisation is destroyed — the lone pair is consumed by H⁺. So protonation is energetically unfavourable.
Substituent Effects on Aniline Basicity
| Aniline derivative | pKₐ | Effect |
|---|---|---|
| Aniline (PhNH₂) | 4.62 | Reference |
| p-Nitroaniline | 0.98 | –M + –I of NO₂ at para; further stabilises free amine |
| m-Nitroaniline | 2.45 | –I only (no mesomeric effect at meta) |
| o-Nitroaniline | –0.28 | –M + –I + steric; conjugate acid unstable |
| p-Methoxyaniline | 5.29 | +M from OMe increases electron density on N |
| p-Methylaniline | 5.10 | +I from CH₃ — slightly more basic |
| Diphenylamine (Ph₂NH) | 0.8 | Two rings pulling lone pair away |
| Triphenylamine (Ph₃N) | ~–5 | Not basic by ordinary standards |
Heterocyclic Bases — Pyridine and Pyrrole Contrasted
Pyridine — sp² N, Lone Pair in Plane
In pyridine, nitrogen is sp² hybridised. One electron goes into the π system (making it aromatic). The lone pair occupies an sp² orbital in the plane of the ring — it does NOT contribute to aromaticity, so it IS available for protonation. However, sp² orbitals have more s-character than sp³, pulling the lone pair closer to the nucleus → weaker base than aliphatic amines.
Pyrrole — sp² N, Lone Pair in π System
Pyrrole's nitrogen contributes its lone pair to complete the 6π aromatic system (4n+2 with n=1). This lone pair is fully delocalised and NOT available for protonation. If forced to protonate, pyrrole does so at the α-carbon (not nitrogen), and loses its aromaticity — energetically very costly. Hence pyrrole is actually an extremely weak base (pKₐ = –0.27), but can act as a weak acid (pKₐ ~17) since the pyrrolide anion retains aromaticity!
Fig 13.1 — Pyridine lone pair is in the plane (sp² orbital), available for protonation. Pyrrole lone pair is part of the aromatic π cloud — unavailable for protonation.
Acid/Base Catalysis — Specific vs General
Catalysis in homogeneous solution works by providing an alternative reaction path of lower activation energy. The acid–base distinction determines which step is rate-limiting.
Specific Acid Catalysis
- Rate depends only on [H₃O⁺], not on other acids present
- Characteristic when: fast, reversible protonation first → then slow rate-limiting step
- Example: Acetal hydrolysis — Rate = k[H₃O⁺][acetal]
General Acid Catalysis
- Rate depends on [H₃O⁺] AND concentration of other proton donors (HA)
- Characteristic when: protonation itself is the slow, rate-limiting step
- Example: Orthoester hydrolysis — Rate = k₁[H₃O⁺][substrate] + k₂[HA][substrate]
Specific Base Catalysis
- Rate depends only on [⁻OH]
- Characteristic when: fast, reversible proton removal → slow step follows
- Example: Aldol reversal — Rate = k[⁻OH][substrate]
General Base Catalysis
- Rate depends on [⁻OH] AND other bases (B:) present
- Characteristic when: proton removal is slow (rate-limiting)
- Example: Base-catalysed bromination of acetone in acetate buffer — Rate = k₁[⁻OH][acetone] + k₂[MeCO₂⁻][acetone]
The key distinction: if adding extra acid (HA) at constant pH speeds up the reaction → general acid catalysis. If only [H₃O⁺] matters → specific acid catalysis. Same logic applies for bases. The mechanism indicator: specific catalysis → proton transfer is fast; general catalysis → proton transfer is rate-limiting.
🎯 High-Value Exam Tips, Tricks & Key Facts
🏆 Acidity Order (Must Know)
Carboxylic acids ≫ phenols ≫ alcohols ≫ water ≫ C–H acids. Memorise: HCOOH (3.77) < CH₃COOH (4.76) < PhOH (9.95) < MeOH (16) < CH₄ (43).
🏆 Hybridisation & Acidity
sp > sp² > sp³ in acidity (and basicity reversed for N atoms). C–H in acetylene (sp, pKₐ≈25) > ethylene (sp², pKₐ≈44) > ethane (sp³, pKₐ≈50).
🏆 Halogen Effects
α-Halo acid order: F > Cl > Br > I. More halogens = much stronger acid. Inductive effect falls off rapidly with distance: α ≫ β > γ.
🏆 Ortho Effect in Benzoic Acids
Ortho substituents often show anomalous behaviour due to (1) intramolecular H-bonding stabilising anion or (2) steric inhibition of resonance. Always check for H-bonding!
🏆 Amine Basicity in Water
R₂NH > RNH₂ > R₃N > NH₃ is the typical order in water. Tertiary drops because its conjugate acid R₃NH⁺ has only ONE N–H bond for H-bonding with water.
🏆 Pyridine vs Pyrrole
Pyridine = good base (sp² lone pair in plane, pKₐ=5.21). Pyrrole = extremely weak base (lone pair in aromatic π cloud, pKₐ=–0.27) but can act as weak acid (pKₐ≈17, anion is aromatic).
🏆 Guanidine Strength
Guanidine (pKₐ≈13.6) is among the strongest organic bases because its protonated form has charge spread symmetrically over THREE equivalent N atoms — maximum delocalisation.
🏆 Maleic < Fumaric pKₐ₁
Maleic (cis) is more acidic (pKₐ₁=1.92) because intramolecular H-bonding stabilises its mono-anion. For pKₐ₂, fumaric is more acidic (4.38 vs 6.23) because maleate must lose H⁺ from a cyclic intramolecularly H-bonded anion.
🏆 Levelling Effect
HCl, HBr, HI, HNO₃, HClO₄ all appear equally strong in water. Use non-aqueous solvents (glacial acetic acid, DMSO) to distinguish their true intrinsic strengths.
🏆 ΔH° vs ΔS° in Acidity
For carboxylic acids in water, ΔH° ≈ 0 (bond energy ≈ solvation energy). The actual difference in Ka between acids comes from ΔS° — entropy of solvation. This is a standard IIT-JAM theoretical question.
Quick-Reference pKₐ Summary Table
| Compound | Type | pKₐ |
|---|---|---|
| Trifluoroacetic acid (CF₃COOH) | Acid | 0.23 |
| Trichloroacetic acid (CCl₃COOH) | Acid | 0.65 |
| Picric acid (2,4,6-(NO₂)₃PhOH) | Acid | 1.02 |
| Oxalic acid (pKₐ₁) | Acid | 1.23 |
| Propynoic acid (HC≡CCOOH) | Acid | 1.84 |
| Formic acid (HCOOH) | Acid | 3.77 |
| Benzoic acid (PhCOOH) | Acid | 4.20 |
| Acrylic acid (CH₂=CHCOOH) | Acid | 4.25 |
| Acetic acid (CH₃COOH) | Acid | 4.76 |
| Pyridine (conjugate acid) | Base | 5.21 |
| Phenol (PhOH) | Acid | 9.95 |
| Ammonium ion (NH₄⁺) | Base ref. | 9.25 |
| MeNH₂ (methylamine conj. acid) | Base | 10.64 |
| Methanol (MeOH) | Acid | ~16 |
| Guanidine (conjugate acid) | Base | ~13.6 |
| Methane (CH₄) | Acid | ~43 |
50 Previous Year Questions
Acids & Bases — Organic Chemistry
Scientifically validated answers with detailed explanations
50 PYQs — Acids & Bases
The correct order of acid strength among the following is:
(I) CHCl2COOH (II) CCl3COOH (III) CH3COOH (IV) CH2ClCOOH
More Cl atoms at the α-carbon = stronger –I (inductive) effect = greater stabilisation of the carboxylate anion = stronger acid.
pKₐ values: CH₃COOH (4.76) > ClCH₂COOH (2.86) > Cl₂CHCOOH (1.25) > Cl₃CCOOH (0.65)
Rule: Each additional Cl spreads the negative charge further over the anion, reducing the tendency of H₃O⁺ to recombine — equilibrium shifts right, stronger acid.
Which of the following is the strongest acid?
Picric acid (2,4,6-trinitrophenol) has pKₐ = 1.02 — far lower than any of the others listed. Three –NO₂ groups exert both –I (inductive) and –M (mesomeric) electron-withdrawing effects, massively stabilising the picrate anion by delocalising negative charge over multiple oxygen atoms of three nitro groups plus the ring.
Remember: lower pKₐ = stronger acid.
Among the following the weakest acid is:
Hybridisation governs the electron-donating ability: sp³ > sp² > sp in +I effect. The tert-butyl group (all sp³ C) has the strongest +I effect, destabilising the carboxylate anion most — weakest acid. sp carbon in propynoic acid has 50% s-character, withdrawing electrons strongly — strongest acid in this set.
Order: HC≡CCOOH < CH₂=CHCOOH < CH₃CH₂COOH < (CH₃)₃CCOOH (acidity decreasing)
The correct order of basic strength of the following amines in aqueous solution is:
NH₃, CH₃NH₂, (CH₃)₂NH, (CH₃)₃N
In water, two effects compete: (1) +I effect of alkyl groups (increases basicity) and (2) solvation of RₙNH(3-n)⁺ cation via H-bonding (decreases as N–H bonds are replaced).
pKₐ values: NH₃ (9.25) < (CH₃)₃N (9.80) < CH₃NH₂ (10.64) < (CH₃)₂NH (10.77)
Tertiary amine drops back because its conjugate acid (CH₃)₃NH⁺ has only one N–H bond — minimal hydrogen bonding with water, poorly solvated, less stable cation.
Which of the following pairs correctly represents an acid and its conjugate base?
By Brønsted–Lowry definition: an acid donates H⁺ to give its conjugate base (differs by one H⁺). H₂SO₄ → HSO₄⁻ + H⁺ ✓ and H₂O → OH⁻ + H⁺ ✓. Both pairs are valid conjugate acid–base pairs.
Option D is wrong because NH₃/NH₄⁺ shows base→conjugate acid (reverse direction), and H₂O/H₃O⁺ is also base→conjugate acid. These are conjugate base/acid pairs, not acid/conjugate base.
Aniline (C₆H₅NH₂) is a much weaker base than cyclohexylamine (C₆H₁₁NH₂). This is because in aniline:
In aniline, the N lone pair overlaps with the benzene ring π orbitals (mesomeric donation, +M). Four resonance structures can be drawn, placing positive charge on N and negative charge at ortho/para positions. This stabilises free aniline. On protonation (→ anilinium cation), this delocalisation is completely destroyed — protonation is energetically unprofitable.
pKₐ(aniline) = 4.62 vs pKₐ(cyclohexylamine) = 10.68 — a difference of ~10⁶ in basicity!
For ionisation of acetic acid in water, ΔH° ≈ –0.5 kJ/mol and TΔS° ≈ –27.6 kJ/mol. Which statement is correct?
ΔH° is tiny (≈0) because energy to break O–H bond is nearly cancelled by solvation energy of the ions produced. The positive ΔG° (acid only partially ionised) comes almost entirely from the large negative TΔS° term — solvation of RCO₂⁻ and H₃O⁺ severely orders surrounding water molecules, reducing entropy.
This is a defining IIT-JAM concept: acid strength differences in the carboxylic acid series are entropy-controlled, not enthalpy-controlled.
The pKₐ of maleic acid (cis-butenedioic acid, pKₐ₁ = 1.92) is much lower than that of fumaric acid (trans-butenedioic acid, pKₐ₁ = 3.02). The reason is:
In the cis (maleate) mono-anion, the –COO⁻ and –COOH groups are on the same side of the double bond. The O⁻ of the carboxylate can form an intramolecular hydrogen bond with the O–H of the adjacent carboxyl group. This extra stabilisation of the anion lowers ΔG° for ionisation → lower pKₐ₁.
Trans (fumaric) geometry places the two –COOH groups on opposite sides — intramolecular H-bonding is geometrically impossible, so no such stabilisation exists.
The order of acidity of the following nitrophenols is:
o-NO₂C₆H₄OH, m-NO₂C₆H₄OH, p-NO₂C₆H₄OH
pKₐ values: o-NO₂PhOH = 7.23, m-NO₂PhOH = 8.35, p-NO₂PhOH = 7.14
At ortho and para positions, –NO₂ exerts both –I (inductive) and –M (mesomeric) electron withdrawal — stabilising the phenoxide anion through extended delocalisation. At meta, only –I effect operates (no mesomeric pathway). Hence meta isomer is the weakest acid.
Note: o-isomer has additional intramolecular H-bonding contribution, making it slightly more acidic than p-.
Pyrrole (pKₐ ≈ –0.27 as base) is a much weaker base than pyridine (pKₐ = 5.21) despite both having a ring nitrogen. The correct explanation is:
In pyridine: N is sp² hybridised; one electron enters the π system; the lone pair is in an sp² orbital in the plane of the ring — NOT part of aromatic cloud — available for H⁺.
In pyrrole: N contributes both electrons of its lone pair to complete the 6π aromatic system (4n+2, n=1). The lone pair is fully delocalised around the ring — unavailable for protonation. If forced to protonate, it does so at C-2 (α-carbon), destroying aromaticity — extremely unfavourable energetically.
Pyrrole can act as a weak acid (pKₐ ≈ 17): deprotonation at N–H gives pyrrolide anion which retains aromaticity.
Which of the following is a Lewis acid but NOT a Brønsted acid?
BF₃ has an empty 2p orbital on boron — it accepts an electron pair (Lewis acid). However, it has no transferable H⁺ — it cannot donate a proton, so it is NOT a Brønsted acid.
HCl, H₂SO₄, CH₃COOH are all Brønsted acids (proton donors) AND Lewis acids (the H⁺ they release is itself a Lewis acid). BF₃ is the only example that is exclusively a Lewis acid.
Salicylic acid (o-hydroxybenzoic acid, pKₐ = 2.98) is much stronger than m-hydroxybenzoic acid (pKₐ = 4.08). The primary reason is:
After ionisation of the –COOH group, the resulting –COO⁻ in salicylate forms an intramolecular hydrogen bond with the adjacent –OH group at the ortho position. This extra stabilisation of the anion (not available in the m-isomer) significantly lowers ΔG° of ionisation → lower pKₐ → stronger acid.
The effect is even more dramatic for 2,6-dihydroxybenzoic acid (pKₐ = 1.30) where H-bonding occurs from both ortho positions.
Guanidine [HN=C(NH₂)₂] has pKₐ ≈ 13.6, making it one of the strongest organic bases. The exceptional basicity is best explained by:
On protonation, guanidinium cation [H₂N–C(=NH₂)–NH₂]⁺ has its positive charge symmetrically delocalised over all three nitrogen atoms via three equivalent resonance structures of equal energy. This is far more effective than in the neutral molecule (where charge separation makes two of the three structures higher energy).
The cation is greatly stabilised relative to the neutral base → protonation is strongly energetically favourable → extremely strong base. Compare: amidines RC(=NH)NH₂ are also strong bases (pKₐ ≈ 12) for the same reason but with only 2 N atoms.
Which statement about the levelling effect of water is correct?
The levelling effect occurs because water is a sufficiently strong base to completely remove H⁺ from any acid stronger than H₃O⁺ (pKₐ ≈ –1.7). All such acids appear fully ionised and thus identically strong. To distinguish true strengths of HCl, HBr, HI, HNO₃, HClO₄, one must use a weaker base solvent such as glacial acetic acid where they are not fully ionised.
In the series: NH₃ → RNH₂ → R₂NH → R₃N, measurements of basicity in chlorobenzene (a non-H-bonding solvent) give the order BuNH₂ < Bu₂NH < Bu₃N. But in water, Bu₃N (pKₐ = 9.87) is weaker than Bu₂NH (pKₐ = 11.28). This is because:
In chlorobenzene (no H-bonding), only the +I inductive effect of alkyl groups operates → basicity increases steadily: primary < secondary < tertiary. In water, an additional factor: solvation of the ammonium cation via N–H···OH₂ hydrogen bonds. Tertiary ammonium R₃NH⁺ has only ONE N–H bond → weakest H-bonding → poorest solvation → least stabilised cation → lowest basicity in water. This perfectly explains the reversal and confirms solvation (entropy effect) as the cause.
The pKₐ of p-hydroxybenzoic acid (4.58) is greater than benzoic acid (4.20). This means p-hydroxybenzoic acid is a weaker acid than benzoic acid. The reason is:
–OH is a +M group (electron donor by resonance) when at ortho or para positions. At para, it donates electron density into the ring, which is then relayed to the –COO⁻ group — increasing negative charge on an already negative carboxylate, destabilising it and making proton removal harder.
The –I effect of –OH would increase acidity, but at para the +M effect dominates the –I effect → net acid-weakening → pKₐ rises above benzoic acid. At meta, only –I operates → m-hydroxybenzoic acid (pKₐ = 4.08) is stronger than benzoic acid (4.20).
Acetic acid (CH₃COOH) is a stronger acid than methanol (CH₃OH) primarily because:
After CH₃COOH loses H⁺, the acetate anion CH₃COO⁻ delocalises its negative charge over both oxygen atoms equally via resonance (two equivalent canonical forms). This massive stabilisation of the anion shifts the ionisation equilibrium to the right.
After CH₃OH loses H⁺, the methoxide CH₃O⁻ has the full negative charge on a single oxygen — no delocalisation possible. Much less stable anion → equilibrium lies far to the left → pKₐ(MeOH) ≈ 16 vs pKₐ(AcOH) = 4.76.
In specific acid catalysis, the rate of reaction depends on:
Specific acid catalysis: mechanism involves rapid, reversible protonation of substrate by H₃O⁺ to form a more reactive intermediate, followed by the slow rate-limiting step. Since protonation is fast and equilibrium-controlled, only [H₃O⁺] (which sets the equilibrium position) matters. Adding NH₄⁺ or any other weak acid at the same pH does NOT change the rate.
Rate = k [H₃O⁺][substrate] — e.g., acetal hydrolysisContrast with general acid catalysis: protonation IS the slow step → any acid HA can catalyse → Rate = k₁[H₃O⁺][S] + k₂[HA][S]
Benzoic acid (pKₐ = 4.20) is a stronger acid than cyclohexane carboxylic acid (pKₐ = 4.87). This is due to:
The C-1 of benzene ring (sp² hybridised, 33% s-character) is less electron-donating (+I) than the sp³ C of cyclohexyl (25% s-character). More s-character = electrons held closer to nucleus = less electron donation to –COOH. Result: benzoic acid's carboxylate is slightly better stabilised → lower pKₐ → stronger acid. Same principle explains why acrylic acid (sp² α-C) is stronger than propionic acid (sp³ α-C).
The pKₐ of formic acid (HCOOH) is 3.77 and that of acetic acid (CH₃COOH) is 4.76. Formic acid is stronger because:
The –CH₃ group in acetic acid exerts a +I (electron-donating inductive) effect. This pushes electrons toward the already-negative carboxylate oxygen atoms, destabilising the acetate anion and promoting recombination with H⁺. Formate (HCOO⁻) lacks this destabilising effect — its anion is better stabilised → lower pKₐ → stronger acid.
The actual difference is dominated by the solvation entropy of the two anions, but the +I destabilisation of the anion is the structural explanation.
The second dissociation of maleic acid (pKₐ₂ = 6.23) is harder than the second dissociation of fumaric acid (pKₐ₂ = 4.38) because:
The maleate mono-anion forms a cyclic intramolecular H-bond between –COO⁻ and –COOH — this system is more stable (lower energy) and therefore the second H⁺ (from –COOH within this ring) is harder to remove. The fumarate mono-anion has no such cyclic stabilisation → second H⁺ is more accessible → lower pKₐ₂.
Which of the following is the correct relationship? (Kₐ = acidity constant, Kb = basicity constant of conjugate base)
This is a fundamental thermodynamic identity. A stronger acid (larger Kₐ) has a weaker conjugate base (smaller Kb), and vice versa.
Among FCH₂COOH, ClCH₂COOH, BrCH₂COOH, ICH₂COOH, the order of acid strength is:
The inductive (–I) effect order of halogens follows electronegativity: F > Cl > Br > I. More electron withdrawal → better stabilisation of –CH₂COO⁻ anion → lower pKₐ → stronger acid. Fluoroacetic acid (pKₐ = 2.57) is ~100× stronger than acetic acid (4.76).
Triphenylamine (Ph₃N) is essentially non-basic in water. This is because:
Each phenyl group withdraws the N lone pair via resonance (+M of N into ring = –M on N from ring's perspective). With three phenyl groups, the combined resonance withdrawal of the lone pair is so extensive that essentially nothing remains for proton acceptance. pKₐ ≈ –5. Compare: Ph₂NH (pKₐ = 0.8), PhNH₂ (pKₐ = 4.62).
Water is an effective ionising solvent for organic acids because of its: (Select the MOST complete answer)
Two synergistic properties: (1) High ε = 80 lowers electrostatic attraction between ions, stabilising separated ion pairs. (2) Small, polar H₂O molecules form hydrogen-bonded solvation shells around both cations (via O lone pairs) and anions (via O–H···A⁻ H-bonds). Anionic solvation is especially powerful. Together these effects stabilise the dissociated ions enormously, driving ionisation equilibrium to the right.
The pKₐ of oxalic acid (HOOCCOOH) pKₐ₁ = 1.23 is much lower than malonic acid pKₐ₁ = 2.83. This is primarily because:
The carboxyl group (–COOH) is itself a powerful electron-withdrawing group (–I and –M). When two –COOH groups are directly bonded (oxalic acid), each strongly withdraws electrons from the other — maximum inductive effect on the ionising O–H. As the chain lengthens (malonic, succinic), the –I effect attenuates rapidly across saturated C–C bonds and the pKₐ₁ rises toward normal carboxylic acid values.
Amides (RCONH₂) are very weak bases (pKₐ ≈ 0.5) compared to aliphatic amines (pKₐ ≈ 10). This is because in amides:
The resonance delocalisation of the N lone pair into the C=O group is so effective that N has essentially partial double bond character with C. The lone pair is tied up in the π system — not available for protonation. If two C=O groups flank N (imides), the compound may even act as an acid (e.g., phthalimide forms alkali metal salts).
2,4,6-trinitro-N,N-dimethylaniline is about 40,000 times (ΔpKₐ = 4.6) stronger a base than 2,4,6-trinitroaniline. The explanation involves:
The large NMe₂ group cannot fit coplanar with the ring because its methyl groups clash with the adjacent o-NO₂ groups. The molecule twists — N's p orbital is no longer parallel to ring p orbitals. Mesomeric withdrawal of the N lone pair by NO₂ groups is blocked. The NMe₂ lone pair stays on N, available for protonation. The NH₂ group in trinitroaniline IS coplanar (small enough) → full –M withdrawal → lone pair fully delocalised → non-basic. This is a classic example of SIR (steric inhibition of resonance).
The pKₐ of acetic acid at 25°C is 4.76. If the temperature is raised to 50°C, the pKₐ will:
Kₐ is an equilibrium constant and must vary with temperature (van't Hoff equation). Since ΔG° = ΔH° – TΔS°, changing T changes ΔG° and hence Kₐ. Even relative acidities reverse: ethanoic acid is weaker than Et₂CHCOOH below 30°C but stronger above 30°C. This is why pKₐ values are always quoted at a specified temperature (usually 25°C).
General base catalysis (as distinguished from specific base catalysis) is characterised by:
In general base catalysis, proton abstraction from the substrate is the slow, rate-limiting step. Any base (B:) present can participate — not just OH⁻. Classic example: base-catalysed bromination of acetone in acetate buffer.
Rate = k₁[⁻OH][CH₃COCH₃] + k₂[CH₃CO₂⁻][CH₃COCH₃]Specific base catalysis: proton removal is fast and reversible (pre-equilibrium), slow step follows → only [OH⁻] matters.
AlCl₃ acts as a Lewis acid because:
Al in AlCl₃ has only 6 electrons around it (three bond pairs) — the 3p orbital is empty. This empty orbital readily accepts an electron pair from a Lewis base (e.g., Cl⁻ from another AlCl₃, or an organic substrate in Friedel–Crafts reactions). AlCl₃ is a classic Lewis acid catalyst; others include FeCl₃, BF₃, SnCl₄, ZnCl₂.
Among the following which has the highest pKₐ (weakest acid)?
Propanoic acid has no electron-withdrawing halogen substituents — only the electron-donating +I ethyl group, which destabilises the anion. CF₃COOH has three F atoms with maximum –I effect → most stable anion → strongest acid. Remember: higher pKₐ = weaker acid.
The pKₐ of phenol (9.95) is much lower than methanol (~16). Despite both having O–H bonds, phenol is ~10⁶ times more acidic. The BEST explanation is:
Phenoxide (C₆H₅O⁻) has four resonance contributors: O⁻ on oxygen, and negative charge at ortho (×2) and para positions of the ring. This delocalisation stabilises the anion significantly (though less than carboxylate — ring carbons are less electronegative than O atoms). Methoxide (CH₃O⁻) has no resonance — full charge on one O atom — far less stable. Greater anion stability → lower ΔG° ionisation → lower pKₐ.
The conjugate base of H₂PO₄⁻ is:
Conjugate base is always formed by removing ONE proton (H⁺) from the acid. H₂PO₄⁻ loses H⁺ → HPO₄²⁻. The conjugate acid would be H₃PO₄ (one more H⁺ added). PO₄³⁻ is three deprotonation steps away.
Pyrrolidine (fully saturated pyrrole ring, pKₐ = 11.27) has nearly the same basicity as diethylamine (pKₐ = 11.04). This is because:
Pyrrolidine is the fully reduced (saturated) version of pyrrole. All ring carbons are sp³; N is sp³ hybridised with its lone pair in an sp³ orbital — NOT contributing to any aromatic system. It behaves exactly like a normal secondary aliphatic amine (R₂NH). Hence pKₐ ≈ 11.27, very close to diethylamine (11.04). This contrasts sharply with pyrrole (pKₐ = –0.27) where N is sp² and lone pair is in the aromatic π cloud.
The inductive effect of a halogen substituent on the acidity of an aliphatic acid diminishes rapidly with chain length. Which set of pKₐ values best illustrates this?
The inductive effect operates through bonds and attenuates by roughly a factor of 2.8 per additional C–C bond. α-Cl (adjacent) has dramatic effect (pKₐ drops from 4.82 to 2.86). β-Cl shows much less effect (pKₐ = 4.52 — closer to unsubstituted 4.82). γ-Cl has barely measurable effect. The negative charge in the carboxylate becomes progressively more concentrated (less spread) as Cl moves farther away.
o-Nitrobenzoic acid (pKₐ = 2.17) is much stronger than m-nitrobenzoic acid (pKₐ = 3.45) or p-nitrobenzoic acid (pKₐ = 3.43). The extra acidity at ortho is mainly due to:
At the ortho position, the –NO₂ group is very close to –COOH (through-space distance ≈ 2.5 Å). The powerful –I effect of –NO₂ is most effective at short distances. Additionally, direct through-space interaction between the adjacent –NO₂ and –COOH groups cannot be excluded. At para, –M and –I both operate but at longer effective distance. The anomalously high acidity of many o-substituted benzoic acids is a well-established pattern called the ortho effect.
Which is NOT true about the Lewis definition of acids and bases?
Not every Lewis acid is a Brønsted acid. BF₃, AlCl₃, FeCl₃, carbocations — all are Lewis acids (electron pair acceptors) but have no transferable H⁺, so they are NOT Brønsted acids. The Lewis definition is broader: every Brønsted acid is a Lewis acid (the released H⁺ is a Lewis acid), but not vice versa.
Arrange in increasing order of basicity (pKₐ of conjugate acid): pyridine, aniline, ammonia, trimethylamine
pKₐ values: PhNH₂ (4.62) < Pyridine (5.21) < NH₃ (9.25) < (CH₃)₃N (9.80)
Aniline: lone pair delocalised into ring → very weak base. Pyridine: sp² N → lone pair less available than sp³. NH₃: sp³ N, no +I groups. (CH₃)₃N: three +I methyl groups, but solvation limits it. In gas phase, (CH₃)₃N would be the strongest by far; in water, solvation effect reduces the advantage.
Which of the following correctly explains why CF₃COOH (pKₐ = 0.23) has ΔG° ≈ 1.3 kJ/mol while CH₃COOH has ΔG° ≈ 27.2 kJ/mol, yet their ΔH° values are nearly identical?
Both acids have ΔH° ≈ 0 (O–H bond energy ≈ solvation energy). The massive difference in ΔG° is entirely due to ΔS°. The CF₃COO⁻ anion has its charge delocalised over the whole molecule (F atoms spread it) → the anion imposes less rigidity on surrounding water molecules → smaller decrease in entropy on ionisation → TΔS° is less negative → ΔG° is much smaller → Kₐ is much larger → much stronger acid. This is the thermodynamic entropy argument at its most elegant.
Which among these is the strongest base in aqueous solution?
In water, secondary amines are the strongest aliphatic bases because they benefit from two +I methyl groups (more electron density on N) while their conjugate acid (CH₃)₂NH₂⁺ still has TWO N–H bonds for effective H-bonding solvation. The tertiary amine gains from three methyl groups (+I) but its cation (CH₃)₃NH⁺ has only ONE N–H — poor solvation drops its pKₐ back.
Propiolic acid (HC≡C–COOH, pKₐ = 1.84) is a much stronger acid than propionic acid (CH₃CH₂COOH, pKₐ = 4.88). The reason is:
sp hybridised carbon (50% s-character) is far more electronegative than sp³ (25% s-character). Greater s-character = electrons drawn closer to nucleus = stronger –I effect = better anion stabilisation. This also explains why acetylene (HC≡CH, pKₐ ≈ 25) is much more acidic than ethylene (H₂C=CH₂, pKₐ ≈ 44) or ethane (pKₐ ≈ 50).
Pyrrole can act as a weak acid (pKₐ ≈ 17) when treated with strong bases like NaNH₂. The acidic H is on the N–H, and the pyrrolide anion is stable because:
When pyrrole (18a) loses the N–H proton, the resulting pyrrolide anion (20) has 6 π electrons total (4 from ring carbons + 2 from N lone pair) in a cyclic conjugated system — fully aromatic by Hückel's rule (4n+2, n=1). The anion is thus aromatically stabilised. In contrast, when pyrrole gains H⁺ (acts as a base), it must disrupt this aromatic system → much more unfavourable. This is why pyrrole is a better acid than it is a base.
Acetal hydrolysis (MeCH(OEt)₂) shows specific acid catalysis. If one adds NH₄⁺ (a weak acid) to the reaction at constant pH, the rate:
This is the defining test for specific acid catalysis. The mechanism involves fast, reversible protonation by H₃O⁺ only — the pre-equilibrium concentration of the protonated intermediate depends solely on [H₃O⁺] (pH). NH₄⁺ does not change [H₃O⁺] at constant pH, therefore does not affect the rate. If NH₄⁺ DID accelerate the reaction at constant pH, it would be general acid catalysis instead.
The Brønsted base in the reaction: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻ is:
CH₃COO⁻ accepts a proton from H₂O → becomes CH₃COOH. Since it accepts H⁺, it is the Brønsted base. H₂O is the acid here (it donates H⁺). OH⁻ is the conjugate base of H₂O, and CH₃COOH is the conjugate acid of CH₃COO⁻. This reaction shows that acetate ion is a base — the basis of buffer chemistry.
Phthalimide (benzene-1,2-dicarboximide) forms alkali metal salts, indicating it acts as an acid. This is because:
Two C=O groups flanking N each withdraw the N lone pair by –M (resonance) and –I (inductive) effects. The combined withdrawal makes N electron-poor and activates the N–H bond toward ionisation. The resulting phthalimide anion is stabilised by extensive delocalisation over both carbonyl groups and the ring. This makes phthalimide acidic (pKₐ ≈ 8.3) — it reacts with KOH to form potassium phthalimide (key in Gabriel synthesis of primary amines).
The correct order of acidity of hydrogen atoms in ethane, ethylene, and acetylene is:
pKₐ: Ethane (~50) > Ethylene (~44) > Acetylene (~25). As s-character increases (sp³→sp²→sp), C becomes more electronegative → C–H bond becomes more polar → H⁺ more easily released → more acidic. Acetylide anion (sp, 50% s-character) is the most stable carbanion of the three. Terminal alkynes are acidic enough to react with strong bases like NaNH₂ or n-BuLi.
The dielectric constant of a solvent affects acid ionisation primarily by:
Coulomb's law: electrostatic energy between ions ∝ 1/ε. Higher dielectric constant (ε) → lower electrostatic attraction between oppositely charged ions → ion pairs less stable as a pair → they separate more readily → equilibrium favours ionised form. Water (ε = 80) is exceptional; methylbenzene (toluene, ε ≈ 2.4) barely supports ionisation — HCl is almost completely un-ionised in toluene.
Tetraalkylammonium hydroxides (R₄N⁺OH⁻) are as strong as mineral alkalis. This is because:
Tertiary amines can revert from their protonated forms: R₃NH⁺ + OH⁻ → R₃N + H₂O (deprotonation equilibrium). But R₄N⁺OH⁻ has NO hydrogen on N — there is literally no pathway for the positive nitrogen to lose its charge and revert to a neutral amine. The hydroxide is completely dissociated — 100% ionised — giving a solution of OH⁻ concentration equivalent to its molar concentration. Hence basic strength = NaOH = KOH.
FINAL QUESTION — Comprehensive: Arrange the following in decreasing order of acidity:
(I) CF₃COOH (II) CCl₃COOH (III) 2,4,6-(NO₂)₃C₆H₂OH (IV) CH₃COOH (V) C₆H₅OH
pKₐ values (lower = stronger acid):
• CF₃COOH (I): pKₐ = 0.23 — three F atoms (highest electronegativity) maximally stabilise anion via –I, plus entropy effect
• CCl₃COOH (II): pKₐ = 0.65 — three Cl atoms, strong but weaker than F (lower electronegativity)
• 2,4,6-(NO₂)₃C₆H₂OH (III) [Picric acid]: pKₐ = 1.02 — three NO₂ groups (–I + –M), phenol acid strengthened to mineral acid level
• CH₃COOH (IV): pKₐ = 4.76 — resonance-stabilised carboxylate, moderate acid
• C₆H₅OH (V): pKₐ = 9.95 — phenoxide partially stabilised by ring, weaker than carboxylic acids
Final order: I > II > III > IV > V ✓
Based on Chapter 3 of "A Guidebook to Mechanism in Organic Chemistry" by Peter Sykes
