What Electronic Effects Are
A chemical reaction in organic chemistry is accompanied by the breaking of some bonds and the making of others. For that bond fission to occur under the influence of an attacking reagent, the substrate — the molecule under attack — must first acquire centres of high and low electron density. According to NIOS, this polarity can only be developed by the displacement, partial or complete, of bonding electrons due to certain effects.
These displacements of electrons in the substrate are collectively called electron displacement effects or simply electronic effects. The single most useful way to organise them is by their permanence. Some effects are permanent — they are built into the molecule and present whether or not a reagent is nearby. Others are temporary — they appear only at the instant of attack and vanish when the reagent is withdrawn.
| Effect | Nature | Acts in | Symbol |
|---|---|---|---|
| Inductive | Permanent | $\sigma$ bond | I (+I, −I) |
| Electromeric | Temporary | $\pi$ bond (multiple bond) | E (+E, −E) |
| Resonance / Mesomeric | Permanent | delocalised $\pi$ / lone pairs | M or R (+M, −M) |
| Hyperconjugation | Permanent | $\sigma$(C–H)–$\pi$ overlap | — |
The Inductive Effect
In a covalent bond between two dissimilar atoms, the shared electron pair is attracted more strongly towards the atom of higher electronegativity. Consider a haloalkane higher than the halomethane. The halogen atom $\ce{X}$, being more electronegative than carbon, pulls the bonding electrons of the $\ce{C-X}$ bond towards itself. The carbon then carries a partial positive charge $(\delta+)$ and the halogen a partial negative charge $(\delta-)$.
This positively charged first carbon now attracts the bonding electrons of the next $\ce{C-C}$ bond, making the second carbon a little less positive than the first. The polarisation is relayed onward to the third carbon, but with a much smaller magnitude — NIOS notes that nearly zero positive charge is present after the third atom. This transmission of induced charge along a chain of $\sigma$-bonded carbon atoms is the inductive effect: a permanent polarisation that decreases steadily as one moves away from the electronegative atom.
Effects are classified by their direction relative to hydrogen, which is taken as the reference. Any atom or group that withdraws electrons more strongly than the H-atom shows a −I effect (electron-withdrawing). Any group that repels electrons more strongly than hydrogen shows a +I effect (electron-releasing). NIOS gives the standard decreasing orders below.
| Direction | Meaning | Decreasing order (NIOS) |
|---|---|---|
| −I (electron-withdrawing) | pulls electrons harder than H | $\ce{(CH3)3N+} > \ce{-NO2} > \ce{-CN} > \ce{-F} > \ce{-Cl} > \ce{-Br} > \ce{-I} > \ce{-OH} > \ce{-OCH3} > \ce{-C6H5} > \ce{-H}$ |
| +I (electron-releasing) | repels electrons harder than H | $\ce{(CH3)3C-} > \ce{(CH3)2CH-} > \ce{CH3CH2-} > \ce{-CH3} > \ce{-H}$ |
The Electromeric Effect
The electromeric effect is a temporary electron displacement that takes place only in compounds containing multiple covalent bonds — such as $\ce{C=C}$, $\ce{C=O}$ and $\ce{C=N}$ — and only in the presence of an attacking reagent. Unlike the inductive effect, it involves the complete transfer of the $\pi$ electron pair, producing definite positive and negative charges within the molecule.
The shift takes place in the direction of the more electronegative atom and is shown by a curved arrow that begins at the original position of the electron pair and ends at its new position. In a carbonyl group it operates as $\ce{C=O -> {}^+C-O^-}$. The effect is given the symbol E: it is the +E effect when the electron pair is displaced away from the atom or group, and the −E effect when displaced towards it. In the carbonyl example, the carbon experiences a +E effect and the oxygen a −E effect. Once the reagent is removed, the electrons return to their original position.
Inductive vs Electromeric vs Resonance — the three confusion points
Candidates routinely mix these up. Inductive is a partial shift in a σ bond, permanent and distance-decaying. Electromeric is a complete shift of a π pair, temporary and seen only when a reagent attacks. Resonance is a permanent delocalisation of π electrons over the whole conjugated system, present at all times.
Memory hook: I = permanent + σ + partial; E = temporary + π + complete + reagent-driven; M/R = permanent + π + delocalised.
Resonance and the Mesomeric Effect
Resonance is shown by molecules that can be represented by two or more structures, called resonating or canonical structures, none of which alone explains all the properties of the compound. These structures are obtained by redistributing the valence electrons; the real molecule is an intermediate of them all — a single, more stable resonance hybrid. The permanent polarisation produced by this delocalisation is the mesomeric effect, denoted +M (electron-donating into the system) or −M (electron-withdrawing from it).
The classic example is benzene, $\ce{C6H6}$, written as two Kekulé structures whose hybrid is the true molecule. The bond-length evidence is decisive: every $\ce{C-C}$ bond in benzene is 139 pm, an intermediate value between a single bond (154 pm) and a double bond (130 pm), confirming that each bond has partial double-bond character. NIOS gives two further standard examples — the ethanoate (acetate) ion and nitromethane — where a charge or a lone pair is delocalised over two equivalent oxygen atoms.
Resonance and hyperconjugation are the two pillars of carbocation, carbanion and radical stability — read that next to see the effects applied to reactive intermediates.
Hyperconjugation
Hyperconjugation, also called no-bond resonance, involves the conjugation of a $\sigma$ (C–H) bond with an adjacent $\pi$ bond or an empty p-orbital. NIOS illustrates it with propene, where the contributing structures II to IV have no bond between one of the H-atoms and its carbon atom — the hydrogen is held by the conjugated system through delocalisation, which is exactly why the phenomenon is named no-bond resonance. NIOS also asks students to describe hyperconjugation in terms of resonance, underlining the parallel.
The effect is most important for cations. In a carbocation, the $\sigma$ electrons of an $\alpha$ C–H bond can overlap with the empty p-orbital on the positively charged carbon, partially neutralising the charge and stabilising the ion. The greater the number of $\alpha$ C–H bonds available, the more hyperconjugative structures can be drawn and the more stable the cation.
Master Table of the Four Effects
The four effects are best fixed in memory side by side. The summary below maps each effect to its permanence, the bond it acts in, its directional forms, and a representative example with its consequence, all grounded in the NIOS account.
| Effect | Permanent / Temporary | Bond involved | + / − forms | Example & consequence |
|---|---|---|---|---|
| Inductive (I) | Permanent | $\sigma$ bond; partial shift; decays with distance | +I (release, e.g. alkyl); −I (withdraw, e.g. $\ce{-NO2}$, halogen) | $\ce{C-X}$ in haloalkane; explains acid strength of carboxylic acids |
| Electromeric (E) | Temporary (reagent-driven) | $\pi$ bond; complete transfer of the pair | +E (away from atom/group); −E (towards it) | $\ce{C=O -> {}^+C-O^-}$ during attack; sets the site of addition |
| Resonance / Mesomeric (M, R) | Permanent | delocalised $\pi$ / lone pairs over the system | +M (donate into system); −M (withdraw from system) | benzene (equal 139 pm bonds); ethanoate ion (charge over two O) |
| Hyperconjugation | Permanent | $\sigma$(C–H) overlapping adjacent $\pi$ / empty p | electron-releasing (no-bond resonance) | propene; tertiary cation stabilised by nine $\alpha$ C–H bonds |
Applications: Stability and Acidity
The whole purpose of cataloguing these effects is prediction. Two NEET-favourite applications follow directly from them: the stability of carbocations and the relative acidity of organic compounds.
Why is the tertiary butyl carbocation more stable than the secondary butyl carbocation?
The tertiary butyl cation $\ce{(CH3)3C+}$ has three methyl groups, i.e. nine $\alpha$ C–H bonds, all able to hyperconjugate with the empty p-orbital. The secondary butyl cation has fewer such bonds. More hyperconjugative structures spread the positive charge more widely, so the tertiary cation is more stable. The +I effect of the alkyl groups assists, but the dominant factor — and the answer NEET 2020 accepted — is hyperconjugation.
Acidity is governed by the stability of the conjugate base. A −I or −M group near an acidic site disperses the negative charge of the anion, stabilising it and raising acidity; a +I group concentrates the charge and lowers acidity. This is precisely why carboxylic acids bearing electron-withdrawing substituents are stronger acids — a property NIOS explicitly attributes to the inductive effect.
Don't credit carbocation stability to the wrong cause
When asked why a tertiary cation outranks a secondary one, the expected primary answer is hyperconjugation, not the +I effect alone, and certainly not resonance (an isolated alkyl cation has no $\pi$ system to delocalise into). Reserve resonance stabilisation for cations such as allyl or benzyl, where a genuine conjugated $\pi$ system exists.
Alkyl cation order $3^\circ > 2^\circ > 1^\circ$ → say hyperconjugation. Allyl / benzyl extra stability → say resonance.
Five things to carry into the exam
- Electronic effects create the polarity a substrate needs before a nucleophile or electrophile can attack (NIOS §23.3.2).
- Inductive = permanent, $\sigma$ bond, partial shift, decays to near-zero after the third carbon; +I releases, −I withdraws.
- Electromeric = temporary, $\pi$ bond, complete transfer, only when a reagent attacks; +E away from the atom, −E towards it.
- Resonance/mesomeric = permanent delocalisation over a conjugated system; the hybrid (benzene, ethanoate ion) is the real, stabler structure.
- Hyperconjugation = no-bond resonance of $\sigma$(C–H) with $\pi$ or an empty p-orbital; the prime reason a tertiary carbocation beats a secondary one.