Chemical Bonding

Physical Chemistry Weightage: AS · Papers 1 and 2 (assumed in Papers 4 and 5) Topic 3

“Welcome! Bonding is the topic Cambridge uses to test whether you can explain, not just recall. Every ‘explain why’ here has a fixed three-step answer: name the particles, name the force between them, and say why that force is strong or weak. Learn that pattern, draw the diagrams with every lone pair and partial charge, and this topic becomes reliable marks.”

— SCORECHEM ACADEMIC TEAM

0. What Examiners Want

How marks are earned in bonding questions

Name the particles and the force. “Strong electrostatic attraction between Mg²⁺ ions and delocalised electrons” scores; “strong bonds” does not.
Compare, then explain. For a difference in melting point, state which force is stronger and why (higher charge, smaller ion, more electrons, more hydrogen bonds).
Diagrams have marking points. Lone pairs, partial charges, dotted hydrogen bonds and correct shapes are each credited separately. A missing lone pair loses a mark.
Keywords must be used correctly. “Ionic” for NaCl, “metallic” for Mg, “simple molecular” for I₂. A correct keyword in the wrong place earns nothing.
Do not mix up bond and force. Covalent, ionic and metallic bonds are strong. Hydrogen bonds, pd–pd and id–id are intermolecular forces.

1. Electronegativity and Bond Polarity

Definition to learn word for word: Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond.

Across a period electronegativity increases: the nuclear charge increases, the shielding stays about the same and the atomic radius decreases, so the nucleus attracts the bonding pair more strongly.
Down a group it decreases: more shells, more shielding and a larger radius.
Most electronegative: F, then O, N and Cl. This is why hydrogen bonding needs N, O or F.
A bond between atoms of different electronegativity is polar: the more electronegative atom has a partial negative charge (δ−) and the other has δ+. The bigger the difference, the more polar the bond, until at a large difference the electron is effectively transferred and the bond is ionic.

Bonding is a continuum. The bigger the difference in electronegativity between the two atoms, the more the bonding pair is pulled towards one atom. A polar bond has partial charges (δ+ and δ−); a symmetrical molecule can have polar bonds and still be non-polar overall (dipoles cancel).

A molecule is polar only if the dipoles do not cancel. CO₂, BF₃, CH₄ and CCl₄ have polar bonds but are symmetrical, so the dipoles cancel and the molecule is non-polar. H₂O, NH₃, HCl and CHCl₃ are polar.

⚠️ Examiner Trap: Polar bonds, non-polar molecule. “CO₂ has no dipole because C and O have the same electronegativity” is wrong. Each C=O bond is polar; the two dipoles point in opposite directions in a linear molecule, so they cancel.

2. Ionic Bonding

Definition: Ionic bonding is the electrostatic attraction between oppositely charged ions. It happens between a metal (loses electrons) and a non-metal (gains electrons), and the ions are arranged in a giant lattice.

Strength of the ionic bond depends on the charges of the ions and their radii. Higher charges and smaller ions give stronger attraction. So MgO (2+ and 2−) melts at a much higher temperature than NaCl (1+ and 1−).
Physical properties: high melting and boiling points (a lot of energy is needed to overcome the strong attraction throughout the lattice); hard but brittle (a shifted layer brings like charges together, which repel and split the crystal).
Electrical conductivity: poor as a solid (ions fixed in the lattice) but good when molten or in aqueous solution (ions are free to move and carry charge).
Solubility: many dissolve in water because the polar water molecules attract and surround the ions. They are insoluble in non-polar solvents.

⚠️ Examiner Trap: Which particles carry the current? In molten NaCl the charge carriers are the ions. In a metal they are delocalised electrons. Do not write “free electrons” for an ionic liquid.

3. Metallic Bonding

Definition: Metallic bonding is the electrostatic attraction between the positive metal ions in a lattice and the sea of delocalised electrons.

Metallic bonding: the electrostatic attraction between the lattice of positive ions and the delocalised electrons. More charge on the ion and more delocalised electrons per atom, with a smaller ion, give a stronger bond and a higher melting point (Al > Mg > Na). The mobile electrons carry charge, so metals conduct when solid.

Strength increases with the number of delocalised electrons per atom, the charge on the ion and a smaller ionic radius. So melting points rise Na < Mg < Al.
Conduct electricity as solids and liquids because the delocalised electrons are mobile. Conduct heat because the electrons and vibrating ions pass energy on quickly.
Malleable and ductile: layers of ions slide over one another and the delocalised electrons keep the ions bonded, so the metal changes shape without breaking.

4. Covalent and Dative Bonding

Covalent bond: the electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms. Dative (coordinate) covalent bond: a shared pair of electrons where both electrons are supplied by the same atom. Once formed it is identical to any other covalent bond. It is shown with an arrow from the donor: NH₃ → H⁺, and H₃N → BF₃.

Dot-and-cross diagrams show outer electrons only. Use dots and crosses to show which atom each electron came from, draw every lone pair, and show the charge with square brackets for an ion. A dative covalent bond is a shared pair in which both electrons come from the same atom.

Drawing dot-and-cross diagrams

Draw only the outer shell of each atom, using dots for one atom's electrons and crosses for the other's.
Put the shared pair(s) in the overlap. A double bond has two shared pairs and a triple bond three.
Show every lone pair. Check each atom has the correct number of electrons: 8 for period 2 atoms (or 2 for H), but some period 3 atoms can have more (PCl₅ has 10, SF₆ has 12).
For ions, use square brackets with the charge outside.
For a dative bond, use the same symbol for both electrons in the pair (both dots) and mention that both electrons come from the donor.

Sigma and pi bonds. A σ bond is the head-on overlap of orbitals along the bond axis. A π bond is the side-on overlap of two p orbitals, above and below the axis. A single bond is one σ; a double bond is one σ + one π; a triple bond is one σ + two π.

In the double bond of ethene, the C atoms use sp² hybrid orbitals for the σ bonds (trigonal planar, 120°) and the unused p orbitals overlap side-on to form the π bond. A π bond is weaker than a σ bond because side-on overlap is less effective, and it stops rotation about the C=C bond.

Hybridisation link: the carbon atoms in ethane use sp³ (four σ bonds, 109.5°), in ethene sp² (three σ bonds, 120°, plus one π) and in ethyne sp (two σ bonds, 180°, plus two π).
The π bond is weaker than the σ bond and it prevents rotation about the double bond, which is why alkenes show cis–trans isomerism (Topic 13).

Bond length and bond energy. Bond energy is the energy needed to break one mole of a bond in the gas phase. As the number of shared pairs between two atoms increases, the bond gets shorter and stronger: C–C (154 pm, 347 kJ mol⁻¹), C=C (134 pm, 612 kJ mol⁻¹), C≡C (120 pm, 838 kJ mol⁻¹). Note that C=C is less than twice C–C because the π component is weaker than a σ bond.

Triple bond strength explains why N₂ is unreactive (N≡N is very strong, and the molecule is non-polar).

⚠️ Examiner Trap: Counting σ and π bonds. Count every bond, including C–H. For CH₃CH=CHCN: nine σ bonds and three π bonds. A double bond is 1 σ + 1 π, and a triple bond is 1 σ + 2 π; never count a triple bond as three σ bonds.

Electron-deficient molecules. BF₃ and AlCl₃ have only six electrons around the central atom and an empty p orbital, so they accept a lone pair. That is why NH₃ forms H₃N→BF₃, and why aluminium chloride dimerises to Al₂Cl₆ (each Al accepts a lone pair from a Cl atom of the other unit, forming dative bonds).

5. Shapes of Molecules and Ions

VSEPR rule: electron pairs around the central atom repel and get as far apart as possible. Lone-pair–lone-pair repulsion > lone-pair–bonding-pair > bonding-pair–bonding-pair, so lone pairs squeeze the bond angle.

VSEPR: electron pairs repel and spread as far apart as possible. Count the electron pairs (bonding + lone) around the central atom, then name the shape from the atoms only. Lone pairs repel more than bonding pairs, so each lone pair squeezes the bond angle by about 2.5° (109.5° → 107° → 104.5°).

Electron pairs	Lone pairs	Shape	Angle	Example
2	0	Linear	180°	BeCl₂, CO₂
3	0	Trigonal planar	120°	BF₃, AlCl₃
4	0	Tetrahedral	109.5°	CH₄, NH₄⁺, PCl₄⁺
4	1	Pyramidal	107°	NH₃, PCl₃
4	2	Bent (V-shaped)	104.5°	H₂O
5	0	Trigonal bipyramidal	90° and 120°	PCl₅
6	0	Octahedral	90°	SF₆, PCl₆⁻

Method: (1) find the number of outer electrons on the central atom, add one for each bond it makes and adjust for the ion charge; (2) divide by two to get pairs; (3) subtract bonding pairs to get lone pairs; (4) match the table. A double bond counts as one region of electron density.

⚠️ Examiner Trap: Name the shape, quote the angle. “Tetrahedral” for NH₃ loses the mark: the arrangement of pairs is tetrahedral but the shape is pyramidal. Also give the angle as a number, and say why it is less than 109.5° (lone pairs repel more than bonding pairs).

6. Intermolecular Forces

Intermolecular forces (IMFs) act between molecules. They are far weaker than covalent bonds. There are three types.

Instantaneous dipole–induced dipole (id–id): in every molecule. Electrons move to one side to form a temporary dipole, which induces a dipole in a neighbour. The force increases with the number of electrons (bigger electron cloud, easier to distort). It explains the rise in boiling point down Group 17 and along the alkanes.
Permanent dipole–dipole (pd–pd): between polar molecules, in addition to id–id.
Hydrogen bonding: the strongest IMF. It needs H bonded to N, O or F, and a lone pair on a different N, O or F. It is an attraction between the δ+ H and the lone pair.

Ranking the forces. When a simple molecular substance melts or boils, only the intermolecular forces break; the covalent bonds inside the molecules stay intact. Never write “covalent bonds break” for a simple molecular substance melting.

Hydrogen bonding between two NH₃ molecules, as Cambridge asks it. Show the lone pair on N, the partial charges (δ+ on H, δ− on N), and the dotted line running N–H···N in a straight line.

Water is special. Each water molecule can form up to four hydrogen bonds (two through its two H atoms and two through the two lone pairs on O). This gives water a high melting and boiling point, high surface tension and a high enthalpy of vaporisation. In ice the molecules are held in an open lattice with more space between them, so ice is less dense than liquid water.

Boiling points of the hydrides rise from period 3 to period 5 because id–id forces grow with the number of electrons. H₂O and HF are ‘out of line’ because hydrogen bonds between their molecules take extra energy to overcome (this is Paper 1 style reasoning: H₂O is a liquid and H₂S a gas at room temperature).

Why H₂O melts higher than NH₃: each H₂O molecule can form two hydrogen bonds on average, but NH₃ has only one lone pair so forms one, and the O–H···O hydrogen bond is stronger because O is more electronegative than N.

⚠️ Examiner Trap: Is it really hydrogen bonding? CH₃F and CH₃OCH₃ have F or O but no H bonded to it, so no hydrogen bonds form between their own molecules. CH₃OH and CH₃CH₂OH do. Check that the H is attached to N, O or F, not to carbon.

7. Structure and Properties

Type	Particles and force	Melting point	Conducts?	Examples
Giant ionic	Ions; strong ionic bonding in a lattice	High	Only molten or aqueous	NaCl, MgO
Giant metallic	Positive ions in a sea of delocalised electrons	High (varies)	Solid and liquid	Na, Mg, Cu
Giant covalent	Atoms; strong covalent bonds throughout	Very high	No (except graphite)	Diamond, SiO₂, graphite
Simple molecular	Small molecules; weak IMFs between them	Low	No	I₂, SiCl₄, H₂O, CO₂

Diamond: each C bonded to four others (tetrahedral), very hard, no mobile electrons. Graphite: each C bonded to three others in layers; the fourth electron is delocalised so it conducts along the layers; the layers are held by weak id–id forces so they slide. SiO₂ is giant covalent; SiCl₄ is simple molecular.
Iodine is a simple molecular solid with weak id–id forces, so it has a low melting point and sublimes on gentle heating. Only the forces between molecules break.
Period 3 chlorides: NaCl and MgCl₂ are giant ionic (solutions pH 7 and about 6); SiCl₄ is simple molecular and is hydrolysed by water to give an acidic solution (pH 1–4).

8. Quick Sheet and Checklist

Definitions: electronegativity, ionic, metallic, covalent, dative.
Trend in electronegativity and the reason for it; polar bond versus polar molecule.
Diagrams: metallic lattice, dot-and-cross (with all lone pairs), NH₃ hydrogen bond, σ and π overlap.
Seven shapes with angles; lone-pair repulsion explanation.
Three IMFs and the H-bond conditions; four hydrogen bonds per water; ice density.
Structure table and the properties explained by particles and forces.

Before you leave this topic, can you:

Draw two NH₃ molecules with the hydrogen bond, including lone pairs and partial charges?
Explain why MgO has a much higher melting point than NaCl?
Give the shape and bond angle of H₂O and explain why the angle is 104.5°?
Say why graphite conducts electricity but diamond does not?

If yes to all four, attempt the ten exam-style questions with the mark schemes covered, then open the flashcards.