Chemical
Bonding

All bonds and intermolecular forces are electrostatic attractions between …

2.1Structure and bonding

Metals have a giant metallic lattice structure, with strong electrostatic forces of attraction between a sea of delocalised valence electrons and a lattice of positive metal cations.

AnalogyEach metal atom gives its valence electrons to a shared pool. The positive cations are held together by the "glue" of the electrons moving among them.

2.2Factors affecting strength

Strength of metallic bond ↑ when

1. Charge of cation ↑ and ionic radius ↓ → charge density of cation ↑
2. No. of valence e⁻ ↑ → more delocalised e⁻ in the sea
3. → stronger attraction → more energy required to overcome it → m.p. ↑

2.3Physical properties

• Conduct electricity in all states: delocalised valence e⁻ become mobile and migrate freely when a potential difference is applied. More delocalised e⁻ → higher conductivity.
• Good thermal conductors: delocalised e⁻ throughout the structure transfer kinetic energy quickly from one end to the other.
• Malleable and ductile: metallic bonds are non-directional, so layers of cations slide over one another without overcoming the attraction between cations and the sea of e⁻.

3.1Structure and bonding

Ionic compounds have a giant ionic lattice structure with strong electrostatic forces of attraction between cations and anions.

DefinitionCoordination number: the number of oppositely charged ions surrounding an ion in a crystal lattice.

Dot-and-cross diagram of MgO

• Mg loses its 2 valence e⁻ (crosses), O gains them.
• Show valence e⁻ only, as dots and crosses
• Square brackets with the charge outside
• No. of each ion as a coefficient, e.g. Ca²⁺ 2[F]⁻

3.2Factors affecting strength

lattice energy ∝ q₊ × q₋r₊ + r₋

Greater charge and smaller ionic radii → greater magnitude of lattice energy → more energy required to overcome the attraction.

AnalogyTwo magnets: stronger magnets, held closer together, are harder to pull apart.

3.3Physical properties

• Conducts when molten or aqueous: ions are mobile and migrate to the electrodes when a p.d. is applied. In the solid, ions are held in fixed positions.
• Hard: a large amount of energy is needed to overcome the strong attraction. Brittle: under force, layers slide, ions of like charges come together and repel → the lattice shatters.
• Soluble in polar solvents: energy released forming ion–dipole interactions compensates for overcoming the ionic attraction and the solvent–solvent forces. Insoluble in non-polar solvents.

3.4Ionic bonds with covalent character

• A small, highly charged cation polarises (distorts) the electron cloud of the anion → electron density drawn between the nuclei → partial sharing of e⁻ → covalent character
• Polarising power of cation ↑ with charge density ∝ q₊ / r₊
• Polarisability of anion ↑ with anion size

4.1Co-ordinate (dative) bonds

• Shown as an arrow from donor to acceptor, or as 2 e⁻ of the same type in a dot-and-cross diagram. Once formed, it is identical to a normal covalent bond.
• Lewis acid: electron pair acceptor. Lewis base: electron pair donor. E.g. BF₃ (acid) + NH₃ (base) → H₃N→BF₃
• Al₂Cl₆: each Al accepts a lone pair from a Cl of the other AlCl₃, so each Al gets an octet.

4.2Drawing steps: CO₂

1. Central atom: forms the most bonds, or is the least electronegative (not H)
2. Anion: add e⁻ to the more electronegative atom(s). Cation: remove e⁻ from the less electronegative atom(s). Square brackets, charge outside
3. Give each atom an octet with single, double or triple bonds
4. Period 2 atoms: maximum 8 valence e⁻

Your turnDraw the dot-and-cross diagrams of (a) H₂O and (b) N₂. Show outer electrons only.

4.3Electronegativity: covalent bonds with ionic character

• Across a period ↑: nuclear charge ↑, shielding about constant → bonding e⁻ more strongly attracted
• Down a group ↓: more shells → bonding e⁻ further away and more shielded
• Different electronegativities → bonding e⁻ drawn to the more electronegative atom → permanent dipole: a polar bond

1. Electron pairs around the central atom are arranged as far apart as possible to minimise repulsion
2. Repulsion: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair

• A lone pair repels more strongly: it is non-bonding, so its electron density is closer to the nucleus
• A double or triple bond counts as one region of electron density

AnalogyRival campers on a tiny round planet pitch their tents as far apart as they can: two end up at opposite poles, three spread round the equator, four sit at the corners of a tetrahedron.

1. Identify any polar bonds (bond dipoles)
2. Determine the shape of the molecule
3. Resolve the dipoles: net dipole → polar

AnalogyA tug of war: two equal teams pulling in opposite directions go nowhere (CO₂). Pull at an angle and the rope moves (H₂O).

4.7Bond energy and bond length

DefinitionBond dissociation energy: the enthalpy change when one mole of a particular covalent bond between atoms in a gaseous molecule is broken.  Cl₂(g) → 2Cl(g)  +244 kJ mol⁻¹

DefinitionBond length: the distance between the nuclei of the atoms in a covalent bond. Shorter bond = stronger bond.

1. Orbital overlap: larger bonding atoms → larger, more diffused valence orbitals → less effective overlap → weaker bond
2. Polarity: may ↑ or ↓ strength (deduce from data), e.g. N–H 390 < O–H 460 < F–H 562
3. Multiplicity: triple > double > single. BE(C=C) < 2 × BE(C–C) as the π bond is weaker than σ

5.1Instantaneous dipole–induced dipole (id–id)

1. Continuous movement of e⁻ → at an instant, one region has slightly higher electron density
2. This instantaneous dipole induces a dipole in a nearby molecule
3. The two molecules attract by weak electrostatic forces between opposite partial charges

What makes id–id stronger?

• No. of electrons ↑ (larger M_r) → electron cloud more polarisable → stronger id–id. F₂ −188 °C, Cl₂ −34 °C, Br₂ 59 °C
• Surface area ↑ → more extensive id–id. pentane 36 °C > 2,2-dimethylpropane 10 °C

AnalogyTwo strips of tape stick along their whole length, two balls touch at one point only.

5.3Hydrogen bonding (H bond)

1. At least one H atom directly bonded to F, O or N
2. The neighbouring molecule has at least one lone pair on F, O or N

Why only F, O and N? They are highly electronegative, so H carries a large δ+, and their lone pair sits in a small 2p orbital with high electron density.

5.3Strength of a hydrogen bond

5.3Extent of hydrogen bonding

AnalogyEach protonic H is a hand, each lone pair is a handle. One H bond = one hand holding one handle, so the number of H bonds per molecule is the smaller of the two numbers.

6.3Density of ice

1. Ice adopts an open, crystalline structure with hexagonal holes
2. Each H₂O forms H bonds with 4 other H₂O molecules, each O is surrounded tetrahedrally by 4 H atoms: 2 covalent bonds, 2 H bonds
3. The regular arrangement spaces molecules more widely than in the liquid → water expands on freezing → ice is less dense than water

6.3Intramolecular H bonds

2-nitrophenol (b.p. 216 °C) < 4-nitrophenol (279 °C)

1. The –OH and –NO₂ groups are in close proximity and correctly oriented for an intramolecular H bond
2. → the H atom of –OH is less available for intermolecular H bonds
3. Boiling overcomes intermolecular H bonds → less energy needed → lower b.p.