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

Here you will learn Structure 2.2.1 to 2.2.10, the opening understandings of the covalent model, following the Chemistry guide that the IB published in February 2023 for Diploma Programme courses first assessed in 2025. Everything in this range is taught at both SL and HL. The single exception is the extension on transition element complexes, which belongs to the additional higher level material and is marked HL only below.

Once you finish, move on to the short-form Structure 2.2 revision notes and then the marked Structure 2.2 practice set. For the wider course, open the Diploma Chemistry overview; to record progress understanding by understanding, print the checklist of Diploma Chemistry understandings; and to find gaps quickly, try a free diagnostic test.

What this unit covers

Understanding	What you must do	Level
Structure 2.2.1	Deduce Lewis formulas (up to four pairs per atom), including incomplete octets	SL and HL
Structure 2.2.2	Link number of bonds to bond length and strength	SL and HL
Structure 2.2.3	Identify coordination bonds; complexes	SL and HL; complexes HL only
Structure 2.2.4	Predict electron domain and molecular geometry (up to four domains) and bond angles	SL and HL
Structure 2.2.5	Deduce bond polarity from electronegativity	SL and HL
Structure 2.2.6	Deduce net dipole moment from polarity and shape	SL and HL
Structure 2.2.7	Describe and explain Si, SiO₂, diamond, graphite, fullerenes, graphene	SL and HL
Structure 2.2.8	Deduce the intermolecular forces present	SL and HL
Structure 2.2.9	Explain volatility, conductivity and solubility	SL and HL
Structure 2.2.10	Explain, calculate and interpret RF values	SL and HL

Structure 2.2.1: Covalent bonds and Lewis formulas

A covalent bond holds two atoms together because both of their nuclei, which are positive, attract the same pair of electrons sitting between them. The octet rule describes the tendency of atoms to reach a valence shell of eight electrons (hydrogen is stable with two).

A Lewis formula shows all the valence electrons in a species, bonding and non-bonding. Pairs may be drawn as dots, crosses or dashes. Below, a dash is a bonding pair and •• is a lone pair.

Method

1. Total the valence electrons. A negative ion gains one electron per unit of charge; a positive ion loses one. Halve the total to get pairs.
2. Put the least electronegative atom (never H) in the centre and join each outer atom with a single bond.
3. Complete the octets of the outer atoms with lone pairs; put any left-over pairs on the central atom.
4. If the central atom is still short of eight, turn an outer lone pair into a double bond.
5. For an ion, add square brackets and the charge.

Worked example 1: methanal, H₂CO. Valence electrons = 2 + 4 + 6 = 12, so 6 pairs. Three single bonds use 3 pairs and three lone pairs on O use the rest. Carbon has only six electrons, so one O lone pair becomes a second C–O bond.

             ••
   H — C  =  O
       |     ••
       H

Worked example 2: nitrosyl chloride, NOCl. 5 + 6 + 7 = 18 electrons, 9 pairs. N is central: Cl–N–O uses 2 pairs, octets on Cl and O use 6, and the last pair goes on N. N then has six electrons, so an O lone pair forms a double bond: Cl–N=O, with one lone pair on N, three on Cl, two on O.

Worked example 3: the amide ion, NH₂⁻. 5 + 2 + 1 = 8 electrons, 4 pairs: two N–H bonds and two lone pairs on N, inside brackets with a 1− charge.

Fewer than an octet. In BCl₃ (24 electrons, 12 pairs), three bonds and nine chlorine lone pairs use every pair, so boron has only six electrons. Beryllium in gaseous BeH₂ has four. Such atoms readily accept a lone pair.

Structure 2.2.2: Bond order, length and strength

Single, double and triple bonds have one, two and three shared pairs. More shared pairs put more electron density between the nuclei, pulling them closer and holding them more firmly.

Worked example 4. Use these typical values:

Bond	Length / pm	Bond enthalpy / kJ mol⁻¹
C–N	147	305
C=N	128	615
C≡N	116	891

More bonds between the same two atoms: shorter and stronger. The triple bond is not three times as strong as the single (891 ÷ 305 = 2.92).

Structure 2.2.3: Coordination bonds

A coordination bond is a covalent bond whose two electrons were both supplied by one of the two atoms. After it forms, you cannot tell it apart from the other bonds of its kind. Look for a donor with a lone pair and an acceptor with a vacancy.

• H₃N–BH₃ (ammonia borane): N donates its lone pair to electron-deficient B.
• CH₃NH₃⁺: the N lone pair of methylamine bonds to H⁺.
• C≡O: in one of the three shared pairs, both electrons come from O.

HL only: transition element complexes. Each ligand (such as H₂O, NH₃, Cl⁻, CN⁻) donates a lone pair to the metal ion. In [Co(NH₃)₆]³⁺ there are six coordination bonds.

Structure 2.2.4: VSEPR shapes

The VSEPR model says electron domains around a central atom repel and get as far apart as possible. An electron domain is a lone pair, a single bond, or a whole multiple bond (a double or triple bond counts as one domain). Electron domain geometry includes lone pairs; molecular geometry describes only the atoms.

Domains	Electron domain geometry	Lone pairs	Molecular geometry	Angle
2	linear	0	linear	180°
3	trigonal planar	0	trigonal planar	120°
3	trigonal planar	1	bent	below 120°
4	tetrahedral	0	tetrahedral	109.5°
4	tetrahedral	1	trigonal pyramidal	about 107°
4	tetrahedral	2	bent	about 104.5°

Lone pairs repel more strongly than bonding pairs, squeezing bond angles. A multiple bond holds more electron density, so it also repels more and opens the angles next to it.

Worked example 5.

• HCN: two domains, no lone pairs: linear, 180°.
• NCl₃: four domains, one lone pair: tetrahedral electron domains, trigonal pyramidal, about 107°.
• NOCl: three domains (single, double, lone pair): bent, below 120°.
• Methanal: three domains on C, trigonal planar. The C=O domain repels more, so H–C–H is a little under 120° and each H–C=O a little over.

Structure 2.2.5: Bond polarity

Unequal sharing makes a bond polar: the more electronegative atom becomes δ−, the other δ+. Show a dipole with partial charges (δ+C–Clδ−) or a vector arrow pointing to the δ− end. The guide states that electronegativity values are given in the data booklet; here they are supplied.

Worked example 6. Values: H 2.2, B 2.0, C 2.6, O 3.4, F 4.0, Cl 3.2. Differences: C–H 0.4, C–Cl 0.6, O–H 1.2, B–F 2.0. So polarity increases C–H < C–Cl < O–H < B–F, with δ− on C, Cl, O and F respectively.

Structure 2.2.6: Molecular polarity

A species has a net dipole moment only if its bond dipoles do not cancel. Find the polar bonds, then use the shape.

Worked example 7. trans-1,2-Dichloroethene has two C–Cl dipoles pointing in opposite directions: they cancel, so it is non-polar. In the cis isomer both chlorines are on one side, so the dipoles add: polar.

• BCl₃: trigonal planar, three equal dipoles cancel: non-polar.
• NCl₃: pyramidal, dipoles cannot cancel: weakly polar.
• CH₂F₂: tetrahedral but two different bond types: polar.
• NH₂⁻ (bent) has a net dipole; NH₄⁺ (tetrahedral) has none.

Structure 2.2.7: Covalent network structures

In a covalent network, covalent bonds run through the whole solid. Allotropes of an element differ in bonding and structure, so they differ in properties.

Substance	Structure	Properties
Diamond	each C bonded to 4, tetrahedral	very hard, very high melting point; no conduction (electrons localised)
Graphite	each C bonded to 3 in hexagonal layers; one delocalised electron per C; London forces between layers	conducts along layers; soft, layers slide
Graphene	one layer of graphite	very strong, excellent conductor, almost transparent
Fullerene C₆₀	separate spherical molecules	much lower melting point (London forces only); poor conductor
Silicon	each Si bonded tetrahedrally to 4	hard, high melting point, semiconductor
SiO₂	each Si bonded to 4 O, each O to 2 Si	hard, very high melting point, insoluble, non-conducting

Si–Si bonds are generally weaker than C–C bonds: Si atoms are larger, so the bonding pair is further from each nucleus.

Structure 2.2.8: Intermolecular forces

• London (dispersion): between all molecules; an instantaneous dipole induces one in a neighbour. Stronger with more electrons and more contact area.
• Dipole-induced dipole: a polar molecule induces a dipole in a non-polar one.
• Dipole–dipole: between polar molecules.
• Hydrogen bonding: when H is covalently bonded to a strongly electronegative atom (usually N, O or F), it is drawn towards a lone pair on an electronegative atom in a nearby molecule.

Use van der Waals forces as the umbrella term for dipole–dipole, dipole-induced dipole and London forces. IUPAC broadened its definition of the hydrogen bond in 2011.

Worked example 8. Pentane and 2,2-dimethylpropane are both C₅H₁₂ (42 electrons). Pentane boils at 36 °C, its branched isomer at about 10 °C: the longer molecules have more contact, so stronger London forces.

Structure 2.2.9: Strength of forces and properties

At comparable molar mass: London < dipole–dipole < hydrogen bonding.

Worked example 9. Butane (58.14 g mol⁻¹) boils at −0.5 °C, propanone (58.09 g mol⁻¹) at 56 °C and propan-1-ol (60.11 g mol⁻¹) at 97 °C. London forces are similar because the molar masses are close. Butane is non-polar (London only). Propanone’s polar C=O adds dipole–dipole forces. Propan-1-ol’s O–H allows hydrogen bonding, the strongest, so it needs most energy to separate its molecules.

• Volatility: simple molecular substances are volatile because only intermolecular forces are overcome; network solids are not, because covalent bonds must break.
• Conductivity: most covalent substances have no mobile charged particles, so they do not conduct; graphite and graphene are exceptions. HCl forms ions in water, so its solution conducts.
• Solubility: like dissolves like. Methanol hydrogen bonds with water and dissolves; iodine dissolves in hexane. Longer non-polar chains lower water solubility. Network solids are insoluble.

Structure 2.2.10: Chromatography

Components separate by their relative attraction, through intermolecular forces, to a stationary phase (water held in paper, or a thin silica or alumina layer in TLC) and a mobile phase (the solvent). Stronger attraction to the stationary phase means a shorter distance travelled.

RF = distance moved by component ÷ distance moved by solvent front (both from the base line). RF has no unit and lies between 0 and 1.

Worked example 10. On a silica TLC plate the solvent front moves 7.6 cm; spots are at 1.9 cm and 5.7 cm. RF = 1.9 ÷ 7.6 = 0.25 and 5.7 ÷ 7.6 = 0.75. Silica is polar, so the low-RF component is probably the more polar one. Identify spots by comparing with known substances on the same plate and solvent, because RF depends on the solvent, stationary phase and temperature. Locating agents, and the operating details of gas chromatography and HPLC, are not assessed.

Common errors

• Counting a double bond as two domains.
• Giving electron domain geometry when molecular geometry is asked.
• Forgetting an ion’s charge in the electron count.
• Calling a molecule polar without checking whether dipoles cancel.
• Writing that boiling a simple molecular liquid breaks its covalent bonds.
• 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.