Electron Displacement Effects in Co-valent Bonds

So far, the unit has looked at electron movement during a reaction. But even before any reagent arrives, an organic molecule's own covalent bonds are not electronically neutral or static -- properties such as stability, reactivity and basicity are all shaped by how electron density is already displaced (or becomes displaced) within a molecule's bonds.

This displacement of electrons can be driven in two different ways:

1. By atoms or groups already present in close proximity to the bond, purely as a structural feature of the molecule itself.
2. By an approaching reagent, i.e. only in the presence of an attacking species.

Correspondingly, electron-displacement effects come in two flavours:

• Permanent effects -- a lasting polarisation built into the molecule's structure by a substituent, present whether or not any reagent is nearby, and capable under the right conditions of leading all the way to bond fission.
• Temporary effects -- a polarisation that appears only while a reagent is actually attacking, and vanishes again once the reagent is removed.

The book categorises electron-displacement effects into four named types, covered one by one in the following sections:

1. Inductive effect (I) -- permanent, operating through sigma bonds.
2. Electromeric effect (E) -- temporary, operating in unsaturated systems in the presence of an attacking reagent. …

The inductive effect (I) is defined as the change in the polarisation of a covalent bond caused by the presence of an adjacent bond, atom or group elsewhere in the molecule. Crucially, this is a permanent phenomenon -- it exists in the molecule's ground state, independent of any reagent.

Worked example -- ethane vs. ethyl chloride. The C-C bond in ethane ($CH_3$-$CH_3$) is non-polar, because both carbons are identical. But in ethyl chloride ($CH_3$-$CH_2$-Cl), the C-Cl bond is polar: chlorine is considerably more electronegative than carbon, so it pulls the shared C-Cl electron pair toward itself, developing a slight negative charge ($\delta^-$) on chlorine and a slight positive charge ($\delta^+$) on the carbon directly attached to it ($C_2$). To partly compensate for this induced positive charge, $C_1$ (the methyl carbon, one bond further away) in turn draws the $C_1$-$C_2$ bonding pair toward itself, picking up a smaller induced charge, $\delta\delta^+$. This chain of induced polarisation, rippling outward from the original electronegative substituent, is exactly what is meant by the inductive effect.

Key properties of the inductive effect:

• It does not transfer electrons wholesale from one atom to another the way ionic bond formation does -- it is a polarisation, a shift in electron density along existing sigma bonds, and this displacement is permanent.
• Its magnitude is greatest at the adjacent bond and falls off rapidly with distance: it is still clearly observable two carbons away from the substituent, but becomes almost insignificant beyond about four bonds.
• It reflects a substituent's inherent ability to either withdraw or donate electron density to the carbon it is attached to.

+I and -I groups. Based on this ability, substituents are classified into two families:

• -I groups (electron-withdrawing): highly electronegative atoms, and atoms/groups carrying a formal positive charge. Examples: $-F$, $-Cl$, $-COOH$, $-NO_2$, $-NH_3^+$. The higher a substituent's electronegativity, the stronger its $-I$ effect. A commonly quoted order of $-I$ strength is:

• +I groups (electron-donating): highly electropositive atoms, and atoms/groups carrying a formal negative charge. Examples: alkali metals; alkyl groups such as methyl and ethyl; negatively-charged groups such as $-OCH_3^-$ (methoxide-type), $-OC_2H_5^-$ and $-COO^-$. The lower a substituent's electronegativity, the stronger its $+I$ effect. Among simple alkyl groups, more branching (more attached alkyl carbons) gives a stronger $+I$ effect:

Effect on reactivity. When a strongly electronegative atom such as a halogen is attached to a carbon, its $-I$ effect polarises the C-X bond, which facilitates attack by an incoming nucleophile at that carbon -- increasing reactivity toward nucleophilic substitution. Likewise, a $-I$ group positioned near a carbonyl carbon reduces the electron density available at that carbon, which increases the rate of nucleophilic addition to the carbonyl. …

The electromeric effect (E) is a temporary effect -- unlike the inductive, resonance and hyperconjugation effects covered elsewhere in this unit, it exists only while a reagent is actually attacking, and disappears again the instant the reagent is removed. It operates specifically in unsaturated compounds -- those containing a multiple bond such as $>C=C<$ or $>C=O$ -- and only in the presence of an attacking reagent.

Mechanism. When a reagent approaches a multiple bond, the mobile $\pi$-electron pair shifts instantaneously -- completely, in a single step, not gradually -- either toward or away from the attacking reagent, forming a new bond in the process.

Worked example -- a nucleophile attacking a carbonyl compound. When a nucleophile (such as $CN^-$) approaches a carbonyl group ($>C=O$), the $\pi$-electrons between carbon and oxygen shift instantaneously onto the more electronegative oxygen. This simultaneously makes the carbonyl carbon electron-deficient, which is exactly what allows -- and accelerates -- the formation of a new bond between the incoming nucleophile and that carbon.

Worked example -- an electrophile attacking an alkene. When an electrophile such as $H^+$ approaches an alkene ($>C=C<$), the $\pi$-electrons shift instantaneously onto the attacking electrophile, forming a new C-H bond; the other alkene carbon, having lost its share of the $\pi$ electrons, becomes electron-deficient and acquires a positive charge -- generating a carbocation.

+E and -E effect. Just like the inductive effect, the electromeric effect is classified by direction:

• +E (positive electromeric effect): the $\pi$-electron pair shifts toward the attacking reagent. The addition of $H^+$ to an alkene, described above, is the textbook example. …

Resonance (also called mesomerism or the mesomeric effect) is a phenomenon seen in certain organic compounds that carry a double bond at a suitable position, and in which the true structure cannot be captured by any single Lewis structure. Such a compound can be drawn as more than one structure, each differing only in where the bonding electrons and lone pairs are placed -- these individual structures are called canonical (resonance) structures, and the compound is said to exhibit resonance.

Resonance hybrid, not equilibrium. It is important to be clear about what resonance actually means: the real molecule is not flipping back and forth between the canonical structures, nor is it a mixture of molecules each existing as one of those structures. The true molecule is a single, fixed structure -- the resonance hybrid -- that lies somewhere between the individual canonical forms, which are themselves only convenient (and individually inaccurate) approximations drawn because a single ordinary Lewis structure cannot represent it.

Worked example -- 1,3-butadiene. In 1,3-butadiene ($H_2C=CH$-$CH=CH_2$, carbons numbered 1 to 4), a single localised structure would predict that the $C_1$-$C_2$ and $C_3$-$C_4$ bonds (drawn as double bonds) should be noticeably shorter than the central $C_2$-$C_3$ bond (drawn as a single bond). But experimentally, all the bond lengths come out nearly the same. This can only be explained if the $\pi$ electrons are not confined to two separate, localised double bonds, but are instead delocalised across all four carbons -- exactly what the resonance description captures, by averaging over the several canonical structures in which the double-bond character shifts along the chain.

Aromatic systems. Benzene is the other classic case the book highlights: like 1,3-butadiene, its true structure (an aromatic ring with uniform bond lengths and exceptional stability) cannot be represented by any single Kekulé structure, and is properly understood only as the resonance hybrid of its canonical forms.

+M and -M (or +R and -R) effect. Just as with the inductive and electromeric effects, the mesomeric effect is classified by the direction in which a substituent's electrons move relative to the conjugated system it is attached to:

• Positive mesomeric effect (+M / +R): occurs when the substituent's electrons move away from the substituent and into the conjugated system -- i.e. the substituent releases electron density by resonance. This happens with electron-releasing substituents, typically ones bearing a lone pair adjacent to the conjugated system: $-OH$, $-SH$, $-OR$, $-SR$, $-NH_2$, $-O^-$, and so on.
• Negative mesomeric effect (-M / -R): occurs when electrons move toward the substituent, out of the conjugated system -- i.e. the substituent withdraws electron density by resonance. This happens with electron-withdrawing substituents, typically ones with a multiple bond to a more electronegative atom, able to accept the incoming electron density: $-NO_2$, $>C=O$, $-COOH$, $-C \equiv N$, and so on. …

Hyperconjugation is the delocalisation of electrons from a sigma bond (typically a C-H or C-C bond) into an adjacent, empty non-bonding p-orbital, or into an adjacent anti-bonding $\sigma^*$ or $\pi^*$-orbital -- in effect extending the molecular orbital system beyond the atoms directly involved in the multiple bond. It is a special stabilising effect, and -- unlike the electromeric effect -- it is a permanent feature of the molecule's structure, present whether or not a reagent is attacking.

Structural requirement. Hyperconjugation needs a specific structural arrangement: an $\alpha$-C-H group (a C-H bond on the carbon directly attached to the unsaturated centre), or alternatively a lone pair on an adjacent atom such as N or O, positioned next to an $sp^2$-hybridised (pi-bonded) carbon. It occurs by the sigma-bonding orbital (or the lone-pair-holding orbital) overlapping with the adjacent $\pi$-orbital or empty p-orbital.

Example 1 -- propene ('no-bond resonance'). In propene ($CH_3$-$CH=CH_2$), the $\sigma$-electrons of each C-H bond of the methyl group can, in turn, delocalise into the $\pi$-orbital of the adjacent C=C double bond. Because the methyl group has three equivalent C-H bonds, this gives three additional canonical structures (structures II, III and IV, alongside the ordinary structure I), one for each hydrogen taking its turn to hyperconjugate. In each of these contributing structures, one particular $\alpha$-C-H bond is shown as effectively broken -- its electrons have moved into the $\pi$ system, leaving that hydrogen as a bare, unbonded (but still nearby) $H^+$, with the vacated bonding position instead shown as part of a shifted double bond, and a negative charge now sitting on the far (terminal) carbon. Because these canonical structures literally show 'no bond' where a C-H bond used to be, hyperconjugation is also known as 'no-bond resonance', or, after the chemists who proposed the explanation, the Baker-Nathan effect. All three of these no-bond canonical structures are inherently polar in nature.

Example 2 -- a lone pair conjugated with a $\pi$ bond (vinyl chloride). Hyperconjugation-type effects are also seen when an atom or group bearing a lone pair is attached by a single bond and sits in conjugation with a $\pi$ bond. In vinyl chloride ($H_2C=CH$-Cl), one of chlorine's lone pairs can enter into resonance with the adjacent C=C $\pi$ bond: the lone pair moves in to form part of a new bond to carbon, displacing the original $\pi$ electrons onto the far (terminal) $CH_2$ carbon. The result is a resonance structure in which the terminal carbon carries a negative charge and chlorine (now doubly bonded to the adjacent carbon) carries a positive charge -- showing that chlorine's lone pair, though formally 'just' a substituent lone pair, genuinely donates electron density into the conjugated system.

Example 3 -- an electronegative substituent withdrawing through conjugation (acrylonitrile). The opposite situation arises when an electronegative atom or group is in conjugation with a $\pi$ bond: instead of donating, it pulls $\pi$-electrons toward itself, away from the multiple bond. In acrylonitrile ($H_2C=CH$-$C \equiv N$), the electronegative nitrile group withdraws $\pi$-electron density from the C=C double bond through the conjugated system, giving a resonance structure with a positive charge on the terminal $CH_2$ carbon and a negative charge on nitrogen, with an extended, cumulated multiple-bond system in between. …