N-Heterocyclic Carbenes (NHCs) and Electronic Effects

Nucleophiles are electron-rich species that donate electrons.
They are typically negatively charged or neutral species with a pair of electrons they can donate.
The most common nucleophiles are neutral and contain a non-bonding lone pair of electrons on a heteroatom (O, N, S, P).
Examples include water ( $H2O$ ), ammonia ( $NH3$ ), phosphine ( $PH3$ ), alcohols ( $ROH$ ), and amines ( $RNH2$ ).
All these examples feature a lone pair of electrons in an $sp^3$ orbital (tetrahedral geometry).
Anions, such as those found on heteroatoms like O, S, or halogens, can also be nucleophiles and have multiple lone pairs of equivalent energy (e.g., hydroxide $OH^-$ , bromide $Br^-$ , cyanide $CN^-$ .
Neutral carbon nucleophiles usually have $\pi$ bonds as a source of electron density (e.g., alkenes, aromatics).

Electrophiles are electron-deficient species that accept electrons.
They can be neutral or positively charged with either an empty atomic orbital or a low-lying anti-bonding orbital ( $\pi^$ or $\sigma^$ ).
Examples include proton (H+), $BF3$ with an empty p orbital, and $AlCl3$ .
Another example with polarized bond (Oxygen more electronegative than Carbon) is carbonyl group with a low lying $\pi^*$ acceptor orbital.

A general example reaction of a nucleophile (Nu) with an electrophile (E+) involves the nucleophile donating electrons to the electrophile.
The curly arrow represents the movement of a pair of electrons.
The nucleophile, having given away electrons, becomes positively charged, while the electrophile, having accepted electrons, becomes neutral.

Curly arrows are the most important symbol in organic chemistry with a very specific meaning.
They indicate the displacement/movement of electron pairs (bonding and lone) and, most importantly, the direction of displacement.
A double-headed arrow represents the movement of two electrons.

The tail of an arrow begins at a lone pair or bonding pair.
The head of an arrow ends at an atom (generates lone pair) or between atoms (generates a single bond).
The movement of electrons cannot create or destroy charge.

Understanding the electronic properties of atoms/groups allows us to understand how molecules react.
We classify atoms/groups as electron-donating or electron-withdrawing.
The three factors that influence the electronic properties of an atom/group are:
- Inductive effects
- Hyperconjugation
- Mesomeric effects / delocalisation

Arise due to electronegativity differences between atoms/groups in a molecule.
Electrons are attracted to the most electronegative atom, resulting in bond polarisation.
This results in bond polarisation, leading to a partial negative charge on the electronegative atom and a partial positive charge on the carbon atom.
These are short-range effects that occur through a sigma bond ( $\sigma$ ).

This polarisation is a negative inductive effect or an electron-withdrawing inductive effect, as it draws electron density from the carbon chain.
The more electronegative the atom, or the more heteroatoms in a group, the more electron-withdrawing it is.
For example, the $-NO_2$ group is strongly electron-withdrawing.

Electropositive metals (Li, Mg) as found in Grignard reagents or organolithium reagents are electron-donating inductive groups.
Alkyl groups (Me, Et, etc.) are the most important class. The fact that alkyl groups are electron-donating explains the order of stability of carbocations.
Carbocations are electron-deficient and are stabilised by electron-donating groups.
Alkyl groups are better electron donor groups than H atoms because the $\sigma$ bonds of the alkyl groups stabilise the cations by hyperconjugation.

Describes the donation of electrons from adjacent C-C or C-H $\sigma$ bonds to the empty p orbital of a carbocation.
This can only happen when these orbitals align (are in the same plane).

Inductive effects push and pull electrons in the $\sigma$ bonds of organic molecules.
Mesomeric effects involve the delocalisation of electron density through $\pi$ bonds.
Both electron-donating and electron-withdrawing groups are known.

Stabilise adjacent carbocations by donation of electrons from a $\pi$ donor group as shown by curly arrows.
Consider the electron donor C=C group.
Only the positions of the electrons move to give resonance forms.
These are not distinct or interconverting structures - the actual structure is a mixture of the two.
The benzyl cation is stabilised by delocalisation of the positive charge around the ring.

Stabilise adjacent carbanions by accepting a pair of electrons (as shown by curly arrows).

Sometimes these two effects work in the same direction, but often in the opposite direction.
The mesomeric effect generally wins as this can spread charge (delocalise charge) over a number of atoms.

Consider acetamide; it contains an amide functional group.
What is the hybridisation of N? $sp^2$ or $sp^3$ ? Remember amines $NH3$ , $RNH2$ , $R1R2NH$ are $sp^3$ ….
Amides are $sp^2$ hybridised at N; the lone pair orbital on N ( $sp^2$ ) can overlap with the $\pi$ orbital of C=O to form a new, low energy molecular orbital = stability gained.
This can only happen if orbitals align - only if lone pair in ‘p’ orbital
The N lone pair has become a bonding pair; neither Lewis structures fully exist; the real structure is a composite and is called a resonance hybrid or delocalised structure.

Nuclear positions must not change, only electrons may move.
A molecule is described by its component resonance structures - weighted according to their relative stability. Equivalent resonance structures contribute equally to the real structure.
Generally, a resonance structure with a negative charge on the most electronegative atom will have the greatest stability (will contribute most to a resonance hybrid). Conversely, a resonance structure with a positive charge on the least electronegative atom will be more stable.

More than two adjacent p orbitals can combine to form a set of molecular orbitals where the electrons are shared, forming a delocalised molecular orbital (of lower energy).

Consider how the 3 x p orbitals can combine to form hybrid molecular orbitals.
3 x p orbitals combine to form 3 molecular orbitals; by overlapping these orbitals a low energy molecular orbital is formed
Lowest energy - all combine in phase to form a bonding orbital
Next highest energy - must have one node (change of sign in combination)
Highest energy orbital - all orbitals combine out-of-phase (2 nodal planes)
A molecular orbital energy diagram can be constructed using these orbitals.
Delocalisation in the allyl cation system can be represented in a simple way by drawing resonance forms.
The curly arrows show the positive charge is shared over both end/terminal atoms.
There are a range of systems where electrons are delocalised over a $\pi$ system (and resonance forms can be drawn).
Allyl anion systems