Acidity & Basicity of Organic Compounds - Chemistry for JEE Main & Advanced

Arrhenius theory of Acids & Bases

According to the Arrhenius concept:

• Acid: A substance that releases H⁺ ions on dissolving in water. Greater the number of H⁺ ions produced in the solution, the stronger is the acid.
• Base: A substance that releases OH^- ions on dissolving in water. Greater the number of OH^- ions produced in the aqueous solution, the stronger is the base.

Dissociation of Acids & Bases

Weak electrolytes (weak acids and weak bases) ionise only to a small extent in water. Most solute molecules remain undissociated and ionic concentrations are relatively low. The extent of dissociation is described by the degree of ionisation (α) or by equilibrium (ionisation) constants.

For a weak monoprotic acid:

HA + H₂O ⇌ H₃O⁺ + A^-

For a weak base:

B + H₂O ⇌ BH⁺ + OH^-

Equilibrium constants of ionisation

The acid ionisation constant, Ka, for HA ⇌ H⁺ + A^- is defined as:

Ka = [H⁺][A^-] / [HA]

The base ionisation constant, Kb, for B + H₂O ⇌ BH⁺ + OH^- is defined as:

Kb = [BH⁺][OH^-] / [B]

For conjugate acid-base pairs in water, the relation holds:

Ka × Kb = Kw where Kw is the ionic product of water (≈ 1.0 × 10^-14 at 298 K).

Use of logarithmic scale: pKa = -log Ka. A smaller pKa (or larger Ka) indicates a stronger acid.

Acid and base ionization constants

Bronsted-Lowry Theory of Acids and Bases

• A Brønsted-Lowry acid is a proton (H⁺) donor.
• A Brønsted-Lowry base is a proton (H⁺) acceptor.
• When an acid donates a proton it forms its conjugate base. When a base accepts a proton it forms its conjugate acid.
• The strength of an acid is measured by its ability to donate protons; the strength of a base is measured by its ability to accept protons.
• Generally, the stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid.

Lewis Theory of Acids & Bases

Lewis acid

• A Lewis acid is an electron-pair acceptor (often has an empty orbital).
• Typical examples include species with an empty p-orbital or electron deficiency, e.g., BR₃ (R = halide or organic group).
• Some species (e.g., H₂O) can act both as Lewis acids and Lewis bases depending on the reaction.

Examples of Lewis acids

Some common Lewis acids that can accept electron pairs include:

• H⁺ (proton) and onium ions like H₃O⁺.
• Cations of d-block elements in high oxidation states (e.g., Fe³⁺).
• Metal cations such as Mg²⁺, Li⁺ in aqua complexes where water acts as ligand.
• Carbocations such as CH₃⁺.
• Pentahalides of group 15 elements (PCl₅, AsCl₅, SbCl₅).
• Electron-deficient π-systems such as enones.

Lewis base

• A Lewis base is an electron-pair donor (species with a lone pair or high-lying HOMO).
• Common Lewis bases: ammonia, alkyl amines and other conventional amines.
• Often Lewis bases are anionic; their base strength is related to the pKa of the corresponding parent acid.
• Lewis bases are nucleophiles; Lewis acids are electrophiles.

Examples of Lewis bases

Examples include:

• Pyridine and its derivatives.
• Group-16 anions or molecules with O, S, Se, Te in oxidation state -2 (e.g., water, ketones).
• Simple anions with lone pairs (H^-, F^-) and complex anions like SO₄^2-.
• Electron-rich π-systems such as benzene, ethyne, ethene.
• Lone-pair bearing species such as CH₃^- and OH^-.

Factors Affecting Acidity (Organic Compounds)

Key factors that influence acidity of an organic acid (ability to donate H⁺) are:

• Electronegativity: More electronegative atom attached to the acidic hydrogen withdraws electron density and stabilises the conjugate base, increasing acidity.
• Inductive effect (-I or +I): Electron withdrawing substituents (-I) stabilise the conjugate base and increase acidity; electron donating groups (+I) decrease acidity.
• Resonance (delocalisation): If the negative charge of the conjugate base can be delocalised by resonance, the conjugate base is stabilised and acidity increases.
• Hybridisation: The greater the s-character of the orbital bearing the negative charge, the more stabilised the conjugate base; thus acidity increases in the order sp > sp² > sp³ (so alkynes > alkenes > alkanes).
• Size / polarizability: For elements down a group, larger atoms better stabilise negative charge by dispersion of charge; acidity increases down the group for hydrides (e.g., H₂O < H₂S < H₂Se < H₂Te).
• Solvation: Stabilisation of ions by solvent molecules (especially water) affects acidity; better-solvated conjugate base increases acidity.
• Aromaticity and special effects: Loss or gain of aromaticity on deprotonation or presence of ortho-substitution can alter acidity (see Ortho effect).

Illustrative examples and guided comparisons

Ex.1 Compare the acidic strength of the following acids.

(a) C - C - C - COOH (b) C = C - C - COOH (c) C≡C-C-COOH

Sol. The acid whose conjugate base is most stable will be more acidic.

After forming conjugate base from the above acids.

(a)

(b)

(c)

Explanation: An sp-hybridised carbon adjacent to the carboxylate stabilises the negative charge most effectively due to higher s-character and greater electronegativity. Hence acidity order:

c > b > a

Ex.2 Which is more acidic between the two:

(a) CHF₃ (b) CHCl₃

Sol. CHF₃ > CHCl₃

After removal of proton:

Note: In haloforms, inductive and other stabilising effects can be complex. The input comment about vacant d-orbitals and pπ-dπ bonding is not generally applicable to carbon; for halogens such as Cl and Br, pπ-dπ interactions are not a primary stabilising factor for conjugate bases of haloforms. The important point is the relative -I effect and polarisation stabilising the negative charge. Thus CHF₃ is generally more acidic than CHCl₃.

Ex.3 Compare the acidic strength of the following:

(a) CHF₃ (b) CHCl₃ (c) CHBr₃

Sol. CHCl₃ > CHBr₃ > CHF₃

Explanation: Although F is more electronegative than Cl or Br, other factors such as polarisation and inductive stabilisation, and the overall ability to stabilise the conjugate base result in this order. For haloforms, CHCl₃ is often more acidic than CHBr₃, and CHF₃ is least acidic among these in many comparisons.

Ex.4 Compare the acidic strength of the following

(a) CH(CN)₃ (b) CH(NO₂)₃ (c) CHCl₃

Sol. After removing H⁺

(Resonance) In its resonating structure, negative charge will be on N)

(Resonance) Negative charge will reside on O (more effective resonance)

→ more effective resonance stabilisation in nitro derivative than cyano derivative.

(pπ-dπ interaction in CHCl₃ is weaker)

Therefore: b > a > c

Reason: Negative charge on O is stabilised better than on N because O is more electronegative; nitro groups withdraw by powerful -I and -M effects giving strong stabilisation.

Ex.5 Compare the acidic strength of the following:

(a) CH≡CH (b) CH₂=CH₂ (c) CH₃-CH₃

Sol.

(Stability of the conjugate base)

Order: a > b > c (acidic strength)

Explanation: The conjugate base of alkyne has negative charge on an sp carbon (50% s-character) and is most stabilised, then sp², then sp³.

Ex.6 Compare the acidic strength of the following :

Sol. d > c > b > a

Ex.7 Compare the acidic strength of the following :

(a) H₂O (b) H₂S (c) H₂Se (d) H₂Te

Sol. Conjugate base stability order:

Therefore: H₂O < H₂S < H₂Se < H₂Te (acidic strength)

Explanation: Down the group, larger size and greater polarizability stabilise the negative charge better, increasing acidity.

Ex.8 Compare the acidic strength of the following compound:

(a)

(b)

(c)

(d)

Sol. After forming conjugate bases:

Order: c > d > b > a

Ex.9 Compare the reactivity of the following compounds with 1 mole of AgNO₃

(a)

(b)

(c)

(d)

Sol. After removing Cl^-:

(Negative charge not stabilised by resonance - least stable)

(Most stable as lone pair of Cl coordinates to positive charge completing octet and stabilising carbocation)

Extent of positive charge decreases as stability increases.

Ex.10 Compare the acidic strength:

(a)

(b)

(c)

(d)

Sol. After making conjugate base:

Order: c > b > a > d

Basic Strength (General Principles)

Basic strength depends directly on the availability of the lone pair for bonding with H⁺ (protonation). Factors affecting basicity include:

• Electronegativity: For atoms bearing the lone pair, lower electronegativity increases availability of the lone pair and thus basicity.
• Hybridisation: Lone pairs on orbitals with greater s-character are held closer to the nucleus and are less available; basicity decreases in the order sp³ > sp² > sp.
• Inductive effect: Electron donating groups increase electron density on the basic atom and increase basicity (in gas phase), while electron withdrawing groups decrease basicity.
• Solvation: In aqueous solution, solvation stabilises conjugate acids; solvation effects can alter basicity order observed in gas phase.
• Steric effects: Bulky substituents can hinder solvation and protonation, reducing basicity (Steric inhibition of protonation).
• Resonance: If the lone pair is delocalised by resonance, it becomes less available for protonation and basicity decreases.

Ex.11 Compare the basic strength of following:

Sol.

Ex.12 Compare the basic strength of the following

(a)

(b)

(c)

(d)

Sol.

Note: The acidity order given in input: CH₄ < NH₃ < H₂O < HF (acidic strength). From conjugate acid viewpoint, HF is strongest acid among these; hence F^- is weakest base among them.

* Strong acids have weak conjugate bases.

* For atoms in the same period, lower electronegativity implies greater nucleophilicity (greater tendency to donate electron pair).

Ex.13 Which is more basic

or

Sol.

>

Which is more basic: NH₃ or ? (forming conjugate acid)

Comparison of Basicity of Ammonia and Alkyl Amines

Ex.14 Compare the basic strength of NH₃, CH₃NH₂, (CH₃)₂NH, (CH₃)₃N

Factors: (1) Steric effect (2) Inductive effect (3) Solvation effect.

The base whose conjugate acid is more stable will be more basic (in that phase/medium).

Stability order of conjugate acids (gas phase):

Therefore basic strength in gas phase or non-polar solvent:

(CH₃)₃N > (CH₃)₂NH > CH₃NH₂ > NH₃

In aqueous solution (polar solvent), due to solvation and hydrogen bonding of the conjugate acids, the order changes:

(CH₃)₂NH > CH₃NH₂ > (CH₃)₃N > NH₃

Explanation: Conjugate acids of 1° amines form more H-bonds with water (more N-H available for hydrogen bonding), so 1° amine conjugate acid is more stabilised in water. Steric hindrance in 3° amine reduces solvation of conjugate acid and lone pair accessibility, lowering basicity in aqueous medium.

Aromatic amines are least basic as the lone pair on N is delocalised into the aromatic ring and less available for protonation.

Combined effects typically yield: 2° > 1° > 3° > NH₃ in aqueous solution.

Ex.15 Compare the basic strength of the following:

(a)

(b)

(c)

(If lone pair participates in resonance, the molecule becomes aromatic.)

Hence lone pair participates in resonance and is less available for H⁺; this compound will be least basic.

Ex.16 Compare the basic strength of the following:

Sol. sp hybridised carbon being most electronegative attracts electron density from nitrogen, making lone pair less available for protonation. Hence basicity decreases.

Order: c > b > a

Ex.17 Compare the basic strength:

(a)

(b)

Answer: a < b

Ex.18 Compare the basicity of the following compounds:

(a) CH₃-CH₂-CH = CH -

(b)

(c)

(d)

Sol. In part (a) the lone pair of nitrogen is in resonance, therefore less available for protonation, making it least basic among all; follow trend based on hybridisation of carbon atoms attached to N.

Order: b > c > d > a

Ex.19 Compare the basicity of the numbered nitrogen atoms.

Sol. The planarity of the ring will be destroyed if lone pair takes part in resonance. Basicity order of nitrogen follows:

N(sp³) > N(sp²) > N(sp)

(In this example, the sp² nitrogen lone pair is in resonance with the ring and thus is least available for protonation - least basic.)

Ex.20 Compare the basic strength of the following:

(a)

(b)

(c)

Sol. In (a) nitro group at para position withdraws electron density by both -M and -I effects, reducing availability of lone pair on N most. In (b) nitro at meta position withdraws by -I only. In (c) there is no such effect.

Order: c > b > a

Ortho effect

Ortho-substituted anilines are less basic than aniline, and ortho-substituted benzoic acids are more acidic than benzoic acid. The ortho effect arises from both steric hindrance and intramolecular interactions; it is particularly significant for benzoic acids and anilines.

Examples:

Ex.21 Compare the basic strength of the following :

(a)

(b)

(c)

(d)

Sol. a > b > d > c

* Due to ortho effect d > c

If c is less basic than d then it will be certainly less basic than b as b is more basic than d.

Ex.22 Compare the basic strength of the following :

(a)

(b)

(c)

(d)

Sol. Do yourselves.

Steric inhibition concepts

S.I.P - Steric inhibition of Protonation (ortho effect)

Bulky ortho substituents hinder protonation; after protonation, steric repulsion increases and the ortho-substituted aniline is less basic than aniline.

S.I.R - Steric inhibition of Resonance

Bulky ortho substituents can also prevent effective resonance by twisting rings or substituents out of conjugation. Examples:

(a)

(b)

Summary (optional)

This chapter covered three classical definitions of acids and bases (Arrhenius, Brønsted-Lowry and Lewis), dissociation and ionisation constants (Ka, Kb, pKa), and the principal structural and electronic factors that determine acidity and basicity in organic compounds: electronegativity, inductive and resonance effects, hybridisation, size and polarizability, solvation and steric effects. Worked examples illustrate how to compare acids and bases using these principles. Use conjugate base stability and lone-pair availability as guiding rules when comparing acidity and basicity respectively.