IB DP Chemistry Structure 2.2: The covalent model: bonds, Lewis formulas, shapes and intermolecular forces – Revision Notes

Use these notes for fast recall of Structure 2.2.1 to 2.2.10, the first block of the covalent model in the IB’s February 2023 Chemistry guide, which applies to Diploma Programme courses first assessed in 2025. Both SL and HL candidates need all of it, apart from transition element complexes, which are HL only. The full teaching, with worked examples, is in the Structure 2.2 study guide.

Test yourself afterwards with the Structure 2.2 practice set. Browse further units from the Diploma Chemistry overview, keep track with the checklist of Diploma Chemistry understandings, or take a free diagnostic test to locate weak areas.

Key definitions

Term	Meaning
Covalent bond	Two positive nuclei both attracting the same electron pair that they share
Octet rule	Atoms tend to bond until their outer (valence) shell holds eight electrons
Lewis formula	Diagram showing every valence electron, bonding and non-bonding; pairs as dots, crosses or dashes
Coordination bond	Covalent bond where both shared electrons come from the same atom
Electron domain	A lone pair, a single bond or a whole multiple bond around a central atom
Electronegativity	Ability of an atom to attract a bonding pair of electrons
Net dipole moment	Overall polarity of a species once bond dipoles are added as vectors
Allotropes	Different structural forms of the same element
van der Waals forces	Umbrella term for dipole–dipole, dipole-induced dipole and London forces
RF	Distance moved by a component ÷ distance moved by the solvent front

Structure 2.2.1: Lewis formulas in steps

1. Count valence electrons (adjust for charge: + for negative, − for positive). Halve for pairs.
2. Central atom = least electronegative (never H). Single bonds to each outer atom.
3. Octets on outer atoms first, then leftover pairs on the centre.
4. Centre still short? Convert an outer lone pair into a multiple bond.
5. Ions: brackets and charge.

Worked reminder: phosgene, COCl₂, has 4 + 6 + 2(7) = 24 electrons, 12 pairs. C is central with C=O and two C–Cl bonds (4 bonding pairs). O carries two lone pairs and each Cl three, so 8 lone pairs; C has none.

Incomplete octets: boron in BCl₃ or BF₃ has six electrons; beryllium in BeH₂ has four. Use organic and inorganic species when practising.

Structure 2.2.2: Bonds, length and strength

Bond type	Shared pairs	Length	Strength
Single	1	longest	weakest
Double	2	shorter	stronger
Triple	3	shortest	strongest

The comparison only works for bonds between the same two elements. More shared pairs mean more electron density between the nuclei, so a stronger pull and a shorter bond.

Structure 2.2.3: Coordination bonds

• Need a donor (lone pair) and an acceptor (empty orbital, such as H⁺ or an electron-deficient B or Al atom).
• Examples: PH₄⁺ (P lone pair to H⁺), H₃O⁺, Al₂Cl₆ (a Cl lone pair to each Al), CO, and adducts such as H₃N–BF₃.
• Once formed, a coordination bond is indistinguishable from the other bonds of that type: all four P–H bonds in PH₄⁺ are identical.
• HL only: in transition element complexes, every ligand forms a coordination bond to the central metal ion.

Structure 2.2.4: VSEPR at a glance

Domains	Lone pairs	Molecular geometry	Angle	Example
2	0	linear	180°	C₂H₂ (each C)
3	0	trigonal planar	120°	BF₃
3	1	bent	below 120°	NOCl
4	0	tetrahedral	109.5°	SiF₄
4	1	trigonal pyramidal	about 107°	NH₃
4	2	bent	about 104.5°	OF₂

Rules of thumb:

• Repulsion: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
• Each lone pair closes the angle by roughly 2–3°.
• A double bond counts as one domain but repels more than a single bond.
• Electron domain geometry follows the total number of domains; molecular geometry ignores lone pairs.

Structures 2.2.5 and 2.2.6: Polarity

• Bond polarity: subtract the electronegativity values. The higher value is δ−. Show with δ+/δ− or a vector arrow towards δ−.
• Molecular polarity: polar bonds + an asymmetric arrangement = net dipole.
• Symmetric shapes with identical outer atoms (linear, trigonal planar, tetrahedral) give no net dipole.
• Lone pairs on the central atom (pyramidal, bent) usually prevent cancelling.
• Different outer atoms on a symmetric shape (CH₂F₂, CHCl₃) give a net dipole.
• Apply the same reasoning to ions: tetrahedral NH₄⁺ has no net dipole; bent NH₂⁻ has one.

Structure 2.2.7: Network structures

Substance	Bonds per atom	Conducts?	Notes
Diamond	C: 4 (tetrahedral)	no	hardest; very high melting point
Graphite	C: 3 (layers, 120°)	yes, along layers	delocalised electrons; layers slide
Graphene	C: 3 (single sheet)	yes, very well	very strong, light, nearly transparent
Fullerene C₆₀	C: 3 (closed cage)	poorly	molecular; London forces between cages
Silicon	Si: 4 (tetrahedral)	semiconductor	diamond-like lattice
SiO₂	Si: 4 O; O: 2 Si	no	very high melting point; insoluble

Si–Si is weaker than C–C because the larger Si atoms form longer bonds.

Structures 2.2.8 and 2.2.9: Intermolecular forces

Force	Present when	Relative strength (comparable molar mass)
London (dispersion)	always	weakest
Dipole-induced dipole	polar next to non-polar	between
Dipole–dipole	polar molecules	stronger
Hydrogen bond	H on N, O or F, plus a lone pair on N, O or F nearby	strongest

Method for deducing forces:

1. Draw or picture the molecule and decide its shape.
2. Every molecule has London forces; more electrons and more surface contact make them stronger.
3. Net dipole? Add dipole–dipole forces.
4. H bonded to N, O or F? Add hydrogen bonding.

Properties:

• Volatility: weaker intermolecular forces mean lower boiling point.
• Conductivity: no ions or delocalised electrons, no conduction (graphite and graphene excepted).
• Solubility: like dissolves like; hydrogen bonding with water helps; long non-polar chains hinder.

Structure 2.2.10: Chromatography

• Stationary phase: water in paper fibres (paper) or silica/alumina (TLC). Mobile phase: solvent.
• Stronger attraction to the stationary phase gives a lower RF.
• RF = component distance ÷ solvent front distance, both from the base line. Range 0 to 1, no unit.
• Same RF as a reference, same plate, same solvent: probably the same substance.
• Locating agents and GC/HPLC operating details are not required.

Linking questions in brief

• Limits of the octet rule: boron and beryllium compounds have fewer than eight electrons; odd-electron species such as NO cannot pair every electron; some period 3 atoms exceed eight (an HL idea, covered in Structure 2.2.11 to 2.2.17).
• Noble gases: their valence shells are already full and their ionization energies are high, so they rarely share electrons; xenon does form a few compounds, such as XeF₂.
• Same-element bonds: identical atoms have equal electronegativity, so neither can take an electron from the other; they can only share.
• Multiple bonds and reactivity: the extra electron density of C=C attracts electrophiles, so alkenes undergo addition; N≡N, by contrast, is so strong that N₂ is very unreactive.
• HL only, Lewis acids and bases: a Lewis base donates a lone pair to a Lewis acid, so the bond formed is a coordination bond.
• How useful is VSEPR? It predicts shapes of main-group molecules reliably but gives only approximate angles.
• Polar bonds and ionic character: a large electronegativity difference gives partial ionic character, so HCl ionizes in water.
• HL only, IR activity: only vibrations that alter a molecule’s dipole moment absorb infrared, so CO₂ absorbs through its asymmetric stretch but not its symmetric one.
• Real gases: attractions between molecules explain why real gases depart from ideal behaviour, especially when compressed or cooled, and why NH₃ departs more than N₂.
• Evidence and functional groups: melting and boiling points, conductivity and solubility tests supply the data; –OH and –NH₂ groups bring hydrogen bonding, C=O brings dipole–dipole forces.

Must-know distinctions

• Bonds vs forces: boiling a simple molecular substance overcomes intermolecular forces, not covalent bonds.
• Electron domain vs molecular geometry: PH₃ has tetrahedral electron domains but a trigonal pyramidal shape.
• Polar bond vs polar molecule: SiF₄ has polar bonds but no net dipole.
• Graphite vs graphene: many stacked layers vs one layer.
• Fullerene vs diamond: molecular vs network.

Quick self-test

1. How many lone pairs are on the central atom of COCl₂, and what is its shape?
2. Predict the molecular geometry and bond angle of OF₂.
3. Explain how the PH₄⁺ ion contains a coordination bond.
4. Using N 3.0, F 4.0, S 2.6, H 2.2, Si 1.9 and O 3.4, rank N–F, S–H and Si–O in order of increasing polarity.
5. Explain why SiF₄ is non-polar.
6. State the H–C–C angle in ethyne, C₂H₂.
7. Name the strongest intermolecular force between hydrazine, N₂H₄, molecules.
8. Name the forces between propanone and tetrachloromethane molecules in a mixture.
9. Explain why graphite conducts electricity but diamond does not.
10. A spot travels 2.6 cm while the solvent front travels 6.5 cm. Calculate RF.
11. A component has RF 0.30. The solvent front travels 8.5 cm. How far does the spot move?

Answers

1. None; trigonal planar (three domains, no lone pairs).
2. Bent; less than 109.5°, about 104.5° (four domains, two lone pairs).
3. The lone pair on P in PH₃ is donated to an H⁺ ion, so both electrons of that P–H bond come from P.
4. S–H (0.4) < N–F (1.0) < Si–O (1.5).
5. Each Si–F bond is polar, but the four dipoles point to the corners of a regular tetrahedron and cancel.
6. 180°: each carbon has two domains (one C–H, one C≡C).
7. Hydrogen bonding (each N carries H atoms and a lone pair).
8. London forces and dipole-induced dipole forces (propanone is polar, CCl₄ is not).
9. Graphite: each C uses three electrons in bonds, and the fourth is delocalised and mobile within its layer. Diamond: every outer electron is fixed in one of four localised bonds.
10. RF = 2.6 ÷ 6.5 = 0.40.
11. 0.30 × 8.5 = 2.55 cm.

Where marks are usually lost

• Drawing outer atoms in a Lewis formula without their lone pairs.
• Forgetting to adjust the electron count for an ion’s charge, or omitting brackets and charge.
• Counting a double or triple bond as more than one electron domain.
• Answering with the arrangement of domains when the question wants the shape of the atoms.
• Stating “lone pairs repel more” without saying what happens to the bond angle.
• Judging molecular polarity from the bonds alone and ignoring the geometry.
• Comparing bond lengths of bonds between different pairs of elements as if bond order alone decided.
• Claiming fullerene C₆₀ has a very high melting point like diamond.
• Writing that covalent bonds break when a simple molecular substance boils.
• Measuring RF to the top of the paper instead of the solvent front.

Official syllabus

Source document: the International Baccalaureate Organization’s Chemistry guide for the Diploma Programme, first assessment 2025, published in February 2023. Sections used: Structure 2.2.1 to 2.2.10.