Following the footsteps of Gilbert N. Lewis in understanding the nature of the covalent bond, Linus N. Pauling stumbled upon an interesting observation. Pauling noticed that the nature of the chemical bond could be explained using a scale or continuum.
A permanent dipole is an inherent feature of the molecule due to the nature of the participating atoms forming the two-electron covalent bond.
Some atoms withdraw more of the bond electrons towards themselves, increasing their electron density and leaving the other end electron deficient. This unequal electron charge distribution creates positive and negative poles (denoted with δ+, δ- signs), causing the molecule to have permanent poles or a dipole.
The redistribution of electrons in an atom, bond, or molecule creates two ends (or poles), one electron-rich negative and the other electron-deficient positive; such an atom, bond, or molecule is said to have a dipole (two poles).
Two electron displacement effects in a covalent bond framework: one polarises sigma bonds along a chain, the other relocates pi electrons through conjugation.
The bond electrons stay in place; only their density shifts. Diminishes rapidly with each bond away from the electronegative source. Permanent.
Electrons move position via overlapping p-orbitals. Transmits across the whole conjugated framework without strong distance decay. Permanent.
The sigma (σ) or covalent bond electrons are pulled or pushed so that the electron density shifts toward the most electronegative atom.
The pi (π) electrons are pulled and pushed through the p-orbitals along the sigma (covalent) bond framework.
The sigma bond electrons do not change their positions. They are only polarised. The bond stays where it is; the electron density along it gets pulled toward one end.
δ+ δ+ δ+ δ+C4H3—δ+ δ+ δ+C3H2—δ+ δ+C2H2—δ+C1H2—δ−NO2
The π electrons change their positions along a conjugated system. The nonbonding electrons, when conjugated, are also involved in resonance.
It is a permanent effect. The polarisation persists as long as the electronegative substituent is part of the molecule.
It is a permanent effect. The delocalisation persists wherever conjugation is maintained through overlapping p-orbitals.
Categorised as negative (−I) and positive (+I) inductive effects. The inductive effect is positive when the substituent is an electron-donating group and negative when the substituent is an electron-withdrawing group.
Categorised as negative (−R/−M) and positive (+R/+M) resonance effects. The resonance effect is positive when the substituent is an electron-donating group and negative when the substituent is an electron-withdrawing group.
−NO2 attached through a sigma bond is electron-withdrawing (−I). The nitrogen pulls sigma electron density along the chain back toward itself.
−NO2 attached to a benzene ring is also electron-withdrawing (−R). The pi electrons of the ring delocalise into the nitro group's pi system, draining electron density from the ring.
−CH3 attached through a sigma bond is electron-donating (+I). The methyl group's electron density is pushed along the sigma framework toward the rest of the molecule.
−OH attached to a benzene ring is electron-donating (+R). The oxygen's lone pair delocalises into the pi system of the ring, raising electron density on the ring.
Decays rapidly with distance. Each successive sigma bond carries less of the polarisation, shown by the diminishing δ+ symbols above. Beyond three or four bonds, the effect becomes negligible.
Transmits through the whole conjugated system. As long as the p-orbital chain is unbroken, the resonance effect reaches every position the conjugation touches, regardless of how many atoms apart they are.
A hydroxyl group (−OH) is attached directly to a benzene ring. Which effect dominates in determining how it influences the ring's electron density?
The inductive and resonance effects are the two permanent ways electron density can be displaced inside a covalent molecule. Both are responses to electronegativity differences and to the structural arrangement of bonds, but they operate through different bond types and at different ranges. The inductive effect runs through sigma bonds; the resonance effect runs through the pi system and conjugated lone pairs. The two effects often coexist in the same molecule and can either reinforce or oppose each other.
The atoms or group of atoms are classified based on the Inductive effect as electron-withdrawing (-I) or electron-donating (+I) relative to Hydrogen.
The common functional groups showing +I and -I effect are:
+I group | -I group |
|---|---|
O- |
Inductive, along with resonance effects, are permanent effects.
The atoms or groups of atoms causing the inductive effect are part of the molecule. Depending on the atom's nature (electron-withdrawing or donating), the groups causing the molecule's inductive effect can impact the molecule's stability and physical and chemical properties.
The Inductive effect is the outcome of Electronegativity; therefore, they are different and not the same.
The +I effect is an electron-donating (or an electron-pushing) inductive effect by an atom or group of atoms relative to Hydrogen. It means that a +I group like methyl (-CH3) will push electrons away from itself more than the hydrogen atom would if it occupied the exact position in the molecule.
The push of electrons is represented with an arrow over the bond, (>) indicating the direction of the electron flow.
Pre-Requisite Reading: Lewis Structures, Types of reactions, Using curly arrows for electron movement, Identifying Functional groups.
Pre-Requisite Reading: Resonance
Electronegativity measures on a scale of 0.8 – 4 an atom’s or group of atoms’ tendency to attract the bond electron pair towards itself, thereby creating partial negative (δ-) and positive (δ+) terminals.
An atom or group of atoms that can pull the bond electrons towards itself or push the bond electrons from itself and decreasingly transmit the effect along the sigma (σ) bonds of the carbon chain inducing permanent polarization in the molecule. Such an effect is called the Inductive effect.
Resonance theory explains various observed properties in a molecule using the electron delocalization concept and multiple Lewis structures, which a single Lewis structure cannot.
A single Lewis structure can only describe some but not all of a molecule's observed properties. Resonance theory is helpful in molecules that can be expressed using several Lewis structures, like Benzene or CO2.
Sharing resources is essential to build a harmonious world. When the resources are shared ineffectively conflicts emerges. A similar principle extends to Chemistry.
Atoms are in a state of harmony when the neighbouring atoms shares electrons. When the sharing stops, Chemical reactions triggers.
The imbalance in sharing electrons influences the molecule’s polarity, reactivity, and physical properties. The disagreement on sharing is brought about by electronegativity.
Pre-Requisite Reading: Electronegativity
A covalent bond is formed between two similar or dissimilar atoms (Ex: Carbon-Carbon and Carbon-Halogen).
The replacement of the hydrogen in acetic acid (H-CH2-COOH) with the chloro gives chloroacetic acid (Cl-CH2-COOH). Therefore, Acetic acid and Chloroacetic acid have carboxylic acid (R-COOH) as the principal functional group.
The dipole moment, a product of charge difference (q) and the distance (d) between the centres of positive and negative charges (µ = q x d), is also directional. The direction is indicated by an arrowhead, which points towards the most electronegative atom, representing the direction of the dipole moment.