JEE Chemistry Chemical Bonding and Molecular Structure Notes Set 08

Back Bonding

Back bonding occurs when one atom of a bonded pair has a vacant orbital (usually a second- or third-period atom such as B or Si) and the other atom has a lone pair (usually a second-period atom such as F, O or N). The lone pair is partly donated back into the empty orbital, giving the bond some double-bond character, which makes it shorter and stronger.
The overlap is best when the two orbitals are of similar size. In boron trihalides the p-p back bonding falls in the order \( \ce{BF3} \) > \( \ce{BCl3} \) > \( \ce{BBr3} \) (2p-2p, 2p-3p, 2p-4p). So the Lewis acid strength runs the other way: \( \ce{BF3} \) < \( \ce{BCl3} \) < \( \ce{BBr3} \). The B-F bond in \( \ce{BF3} \) is shorter than in \( \ce{BF4^-} \), where boron has no vacant orbital left and no back bonding.
If the atom with vacant orbitals also has a lone pair, back bonding is weaker. For third-period atoms bonded to oxygen the extent of p-pi to d-pi bonding rises Si < P < S < Cl, as the d orbitals contract.

Back Bonding and Basicity or Shape

Trisilylamine \( \ce{N(SiH3)3} \) is planar and is not a Lewis base, because the nitrogen lone pair is delocalised into empty 3d orbitals of Si and N becomes sp2. Trimethylamine \( \ce{N(CH3)3} \) is pyramidal and basic, as carbon has no suitable vacant orbital. In \( \ce{N(PH3)3} \) the back bonding is weak, so N stays pyramidal.
In the same way silyl isocyanate \( \ce{SiH3NCO} \) is linear (N is sp, with back bonding to Si) while methyl isocyanate \( \ce{CH3NCO} \) is bent (N is sp2). \( \ce{(SiH3)2O} \) has a wider Si-O-Si angle and is a much weaker base than \( \ce{(CH3)2O} \).

Lewis Acids, Bases and Coordinate Bonds

A Lewis acid accepts an electron pair; a Lewis base donates one. In a coordinate (dative) bond both electrons of the shared pair come from the same atom, as in \( \ce{H3N->BF3} \) and \( \ce{NH4^+} \).

Electron-Deficient Bonding

Some molecules have fewer electrons than needed for ordinary 2-centre 2-electron bonds. They contain 3-centre 2-electron (3c-2e, banana) bonds: \( \ce{B2H6} \), \( \ce{Al2H6} \), \( \ce{(BeH2)n} \), \( \ce{Al2(CH3)6} \) and bridging metal carbonyls.
Diborane has four terminal 2c-2e B-H bonds and two bridging 3c-2e B-H-B bonds. Boron is sp3; the four terminal H and both B lie in one plane and the bridging H lie above and below it, in a perpendicular plane. The bridge B-H bonds are longer than the terminal ones, and the H-B-H angle between terminal bonds is larger than that between bridging bonds. In \( \ce{Al2(CH3)6} \) the bridges are made through the sp3 carbon atoms of the methyl groups.
In polymeric \( \ce{(BeH2)n} \) every Be is surrounded tetrahedrally by four bridging H atoms and all its bonds are 3c-2e. Halide bridges are different: in \( \ce{Al2Cl6} \), \( \ce{(BeCl2)n} \), \( \ce{I2Cl6} \) and solid \( \ce{SnCl2} \) the bridging halogen donates a lone pair to the neighbouring atom (3c-4e bonding), so these molecules are electron precise and there is no electron-deficient bond. The driving force is completion of the octet.

Hydrogen Bonding

A hydrogen bond forms when H, covalently bonded to a small, strongly electronegative atom (F, O or N), is attracted to a lone pair on another such atom. Its strength depends on the electronegativity and small size of the atom, so F-H...F > O-H...O > N-H...N. Typical strength is 5 to 50 kJ per mole, much less than a covalent bond; the strongest known is the symmetrical \( \ce{[F-H-F]^-} \) ion (about 160 kJ per mole). HCN also hydrogen bonds because the sp carbon makes the C-H polar enough.
Intermolecular hydrogen bonding links separate molecules: homo (water, HF chains, boric acid in layered sheets, carboxylic acid dimers) or hetero (alcohol and water). Intramolecular hydrogen bonding is within one molecule and gives a 5- or 6-membered chelate ring, as in o-nitrophenol, salicylaldehyde, o-fluorophenol, the enol form of ethyl acetoacetate and the \( \ce{HSO5^-} \) ion (which is why Caro's acid has K1 much greater than K2).

Effects of Hydrogen Bonding

Melting and boiling points: intermolecular hydrogen bonding raises them, so \( \ce{H2O} \) > \( \ce{CH3OH} \) > \( \ce{CH3OCH3} \), and glycerol > ethylene glycol > ethanol. Carboxylic acids boil higher than alcohols of similar mass because they form dimers. HF, \( \ce{H2O} \) and \( \ce{NH3} \) show anomalously high boiling points compared with the hydrides of heavier members of their groups. Intramolecular bonding lowers the boiling point relative to the isomer that associates: o-nitrophenol is steam volatile, p-nitrophenol is not.
Solubility: hydrogen bonding with water raises solubility (alcohols, acetic acid, glucose, p-hydroxybenzaldehyde) while intramolecular bonding lowers it (salicylaldehyde, o-nitrophenol). Viscosity increases with hydrogen bonding, as in glycerol. Ice is less dense than water because the open hydrogen-bonded cage structure holds the molecules further apart.

Worked Examples

1. Compare B-F bond length in \( \ce{BF3} \) and \( \ce{BF4^-} \): in \( \ce{BF3} \) back bonding gives partial double-bond character, so \( \ce{BF4^-} \) has the longer bond.
2. Why is \( \ce{BF3} \) the weakest Lewis acid of the boron trihalides? Its strong 2p-2p back bonding partly fills the vacant orbital on B, reducing its ability to accept a lone pair.
3. Why does \( \ce{C2H2} \) dissolve in acetone but not in water? Acetone can hydrogen bond with the acidic C-H of ethyne, but water molecules are so strongly associated with each other that ethyne cannot break into the network.