S2.2
The covalent model
10 SL · +8 HL (SL 2.2.1-2.2.10) (HL 2.2.11-2.2.16)
Unverified
← View on the mapParent topic: Models of bonding and structure
Guiding question
Guiding questionWhat determines the covalent nature and properties of a substance?
The one big idea
A covalent bond is a shared electron pair between two nuclei. VSEPR predicts molecular shape from electron-domain repulsion; intermolecular forces determine physical properties. At HL, resonance, formal charge, hybridization and sigma/pi bonding deepen the model.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS1.1 Introduction to the particulate nature of matterIntermolecular forces presuppose the particle model and states of matter.
- EditorialS1.3 Electron configurationsLewis formulas and the octet rule require valence-shell counting.
- EditorialS2.1 The ionic modelThe guide explicitly contrasts the ionic and covalent models (linking question S2.2.1).
What rests on this
Core concepts that must be mastered
- Covalent bonds and Lewis formulas(2.2.1–2.2.3)A covalent bond is electrostatic attraction between a shared electron pair and the positively charged nuclei on either side. The octet rule is a useful guide but not a law: some species have fewer than an octet (BF₃ is electron-deficient) and some have more (expanded octets at HL). You must deduce Lewis formulas for molecules and ions with up to four electron pairs on each atom. Single, double and triple bonds correspond to one, two and three shared pairs: more shared pairs mean shorter, stronger bonds. A coordination bond is a covalent bond where both shared electrons come from the same atom — the classic example is NH₃→BF₃ or the formation of the ammonium ion NH₄⁺. At HL, coordination bonds include transition element complexes.
- VSEPR and molecular geometry(2.2.4)VSEPR predicts molecular shape from the repulsion of electron domains. Count all electron domains around the central atom — bonding pairs and lone pairs — and arrange them to minimise repulsion. Two domains → linear (180°); three → trigonal planar (120°); four → tetrahedral (109.5°). Lone pairs repel more strongly than bonding pairs, so the presence of lone pairs compresses the bond angles: NH₃ (three bonding + one lone pair) is 107°, H₂O (two bonding + two lone pairs) is 104.5°. Multiple bonds count as one domain but repel more than a single bond. Electron domain geometry (arrangement of all domains) and molecular geometry (arrangement of atoms only) can differ — always state both.
- Bond and molecular polarity(2.2.5, 2.2.6)Bond polarity results from electronegativity difference: the more electronegative atom pulls the shared pair closer, acquiring a partial negative charge (δ−). Molecular polarity depends on bond polarity and molecular geometry: symmetric molecules cancel their bond dipoles. CO₂ has two polar C=O bonds but is non-polar because they point in opposite directions. CCl₄, BF₃ and SF₆ all have polar bonds and zero net dipole. A molecule is polar only if the bond dipoles do not cancel — check both the bonds and the shape.
- Network structures, allotropes and intermolecular forces(2.2.7–2.2.10)Carbon and silicon form covalent network structures — giant lattices of covalent bonds with no discrete molecules. Diamond: each carbon is tetrahedrally bonded to four others — hard, high-melting, non-conducting. Graphite: each carbon is bonded to three others in layers, with one delocalised electron per atom — conducts electricity, slippery because layers are held only by London forces. Intermolecular forces: London (dispersion) forces act between all particles and arise from instantaneous fluctuations in electron density — they strengthen with more electrons and greater polarisability. Dipole–induced dipole and dipole–dipole forces require a permanent dipole. Hydrogen bonding requires H covalently bonded to N, O or F, and an accessible lone pair on a neighbouring N, O or F. For comparable molar mass: London < dipole–dipole < hydrogen bonding. These forces determine volatility, electrical conductivity and solubility. Chromatography separates mixture components by their relative attractions (IMFs) to the mobile and stationary phases; R_F = distance travelled by spot ÷ distance travelled by solvent front.
- HL: Resonance, benzene, expanded octets, formal charge, sigma/pi bonds and hybridization(2.2.11–2.2.16)Resonance occurs when more than one valid Lewis formula exists — the real molecule is a single species with delocalised electrons, not a mixture that flips back and forth. Benzene is the key example: six identical C–C bonds of intermediate length, not alternating single and double. Expanded octets: atoms in period 3 onwards can use d orbitals to accommodate 5 or 6 electron domains (PCl₅, SF₆); deduce their geometry by VSEPR. Formal charge = valence electrons − (non-bonding electrons + ½ bonding electrons); the preferred Lewis formula minimises formal charge and places negative charge on the more electronegative atom. Sigma bonds form by head-on orbital overlap (electron density along the bond axis); pi bonds by lateral p-orbital overlap (density above and below the axis). A single bond is one σ; a double bond is σ + π; a triple bond is σ + 2π. Hybridization mixes atomic orbitals into hybrid orbitals: sp (linear, two electron domains), sp² (trigonal planar, three), sp³ (tetrahedral, four). Only sp, sp² and sp³ are assessed.
- How this sub-topic connectsThe covalent model is the largest sub-topic in Structure 2 (10 SL + 8 HL hours) and the second-heaviest HL block in the whole Structure strand. It connects to periodicity (S3.1: electronegativity trends explain bond polarity), to organic chemistry (S3.2: functional groups are covalent, and HL stereoisomerism, spectroscopy and NMR all rest on this model), and to thermodynamics (R1.1: bond enthalpies). At HL, hybridization and sigma/pi bonding are essential for understanding reaction mechanisms (R3.4). Resonance and benzene reactivity connect directly to electrophilic substitution (R3.4).
Quantitative non-negotiables
- Deduce Lewis formulas of molecules and ions for up to four electron pairs on each atom.
- Predict electron domain geometry and molecular geometry for up to four electron domains using VSEPR.
- Deduce net dipole moment from bond polarity and molecular geometry.
- Rank intermolecular forces: London < dipole–dipole < hydrogen bonding (for comparable molar mass).
- Calculate and interpret R_F values in chromatography.
- At HL: deduce resonance structures, expanded octet geometries (5 and 6 domains), formal charge, and hybridization (sp, sp², sp³).HL
Common failure modes
M-08 — Atoms bond because they want a full outer shell.
Confidence: verified
Why it’s wrong: Circular and causally backwards. This is Taber's 'full outer shells explanatory principle'.
Correction: Bonding is electrostatic attraction; full-shell configurations correlate with low energy, they don't cause bonding.
M-10 — Electrons remember which atom they came from.
Confidence: verified
Why it’s wrong: The 'ownership of electrons' conception.
Correction: Electrons are indistinguishable; once transferred or shared, the electronic history is chemically irrelevant.
M-11 — A compound is either ionic or covalent.
Confidence: verified
Why it’s wrong: Bonding is not a binary classification.
Correction: Bonding is a continuum; use Δχ and χ̄ and the bonding triangle. Most bonds are polar covalent to some degree.
M-12 — If a molecule has polar bonds, the molecule is polar.
Confidence: verified
Why it’s wrong: Molecular polarity depends on bond polarity and geometry.
Correction: Symmetric arrangements cancel the dipoles. CO₂, CCl₄, BF₃ and SF₆ have polar bonds and zero net dipole.
M-13 — Boiling water breaks the covalent bonds in H₂O.
Confidence: verified
Why it’s wrong: Boiling breaks intermolecular hydrogen bonds, not covalent bonds.
Correction: The O–H covalent bonds survive — the product is still water molecules, just gaseous.
M-14 — Non-polar molecules have no intermolecular forces.
Confidence: verified
Why it’s wrong: London dispersion forces act between all particles.
Correction: London forces arise from instantaneous fluctuations in electron density; they strengthen with more electrons / greater polarisability. That's why Br₂ is liquid and I₂ solid.
M-15 — Hydrogen bonding happens whenever there's a hydrogen atom.
Confidence: verified
Why it’s wrong: Hydrogen bonding has specific requirements.
Correction: Requires H covalently bonded to N, O or F, and an accessible lone pair on an N, O or F of a neighbouring molecule. HCl, CH₄ and CH₃F do not hydrogen bond.
M-16 — A hydrogen bond is a type of covalent bond.
Confidence: verified
Why it’s wrong: A hydrogen bond is an intermolecular force, not a covalent bond.
Correction: It is an intermolecular force, roughly 5–10% of the strength of a typical covalent bond.
M-17 — Bigger molecule = stronger bonds = higher boiling point.
Confidence: verified
Why it’s wrong: Boiling point depends on intermolecular forces, not bond strength.
Correction: Say 'stronger London dispersion forces because more electrons', never 'stronger covalent bonds'.
M-18 — Resonance means the molecule flips back and forth between two structures.
Confidence: verified
Why it’s wrong: The real molecule is a single species with delocalised electrons.
Correction: The resonance forms are limitations of Lewis notation, not states the molecule visits. Benzene has six identical C–C bonds of intermediate length.
M-19 — All molecules with four things around the central atom are 109.5°.
Confidence: verified
Why it’s wrong: Lone pairs repel more strongly than bonding pairs.
Correction: Count electron domains including lone pairs; lone pairs repel more, so NH₃ is 107° and H₂O 104.5°.
M-20 — Every atom must obey the octet rule.
Confidence: verified
Why it’s wrong: Period 3 onwards can expand the octet; some species are electron-deficient or odd-electron.
Correction: SF₆, PCl₅ expand the octet; BF₃ is electron-deficient; radicals have an odd electron.
M-21 — Giant covalent substances have high melting points because covalent bonds are strong, and simple molecular ones have low melting points because covalent bonds are weak.
Confidence: verified
Why it’s wrong: The covalent bonds are strong in both.
Correction: The difference is that melting a molecular solid only overcomes intermolecular forces, whereas melting diamond or SiO₂ requires breaking covalent bonds throughout the lattice.
M-22 — Graphite conducts because the layers slide.
Confidence: verified
Why it’s wrong: Two separate properties conflated.
Correction: It conducts because each carbon uses only three of four valence electrons in σ bonds, leaving one delocalised electron per atom; it is slippery because the layers are held only by London forces.
What "HL standard" actually looks like
HLHL adds resonance and delocalisation (benzene from physical and chemical evidence), expanded octets with VSEPR for 5 and 6 electron domains, formal charge to select the preferred Lewis formula, sigma vs pi bonds, and hybridization (sp, sp², sp³) linked to geometry. This is the second-heaviest HL block in Structure (+8 h).
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
How do intermolecular forces influence the type of mixture that forms between two substances?
- Official IB
Under comparable conditions, why do some gases deviate more from ideal behaviour than others?
- Official IB
How is formal charge used to predict the preferred structure of sulfate?
- Official IB
Why do noble gases form covalent bonds less readily than other elements?
- Official IB
Why do ionic bonds only form between different elements while covalent bonds can form between atoms of the same element?
- Official IB
How does the presence of double and triple bonds in molecules influence their reactivity?
- Official IB
Why do Lewis acid–base reactions lead to the formation of coordination bonds?
- Official IB
What properties of ionic compounds might be expected in compounds with polar covalent bonding?
- Official IB
What features of a molecule make it "infrared (IR) active"?
- Official IB
Why are silicon–silicon bonds generally weaker than carbon–carbon bonds?
- Official IB
To what extent can intermolecular forces explain the deviation of real gases from ideal behaviour?
- Official IB
How do the terms "bonds" and "forces" compare?
- Official IB
To what extent does a functional group determine the nature of the intermolecular forces?
- Official IB
Why are oxygen and ozone dissociated by different wavelengths of light?
- Official IB
How does the resonance energy in benzene explain its relative unreactivity?
- Official IB
What are the structural features of benzene that favour it undergoing electrophilic substitution reactions?
- Official IB
How does the ability of some atoms to expand their octet relate to their position in the periodic table?
- Official IB
What are the different assumptions made in the calculation of formal charge and of oxidation states for atoms in a species?
- Official IB
What is unique about carbon that enables it to form more compounds than the sum of all the other elements' compounds?
- Official IB
What are the advantages and disadvantages of different depictions of an organic compound (e.g. structural formula, stereochemical formula, skeletal formula, 3D models)?
- Official IB
What is the influence of the carbon chain length, branching and the nature of the functional groups on intermolecular forces?
- Official IB
How does the fact that there are only 3 isomers of dibromobenzene support the current model of benzene's structure?
- Official IB
What features of a molecule determine whether it is IR active or not?
- Official IB
Most combustion reactions are exothermic; how does the bonding in N₂ explain the fact that its combustion is endothermic?
- Official IB
How would you expect bond enthalpy data to relate to bond length and polarity?
- Official IB
Would you expect allotropes of an element, such as diamond and graphite, to have different ΔH⦵f values?
- Official IB
What is the reverse process of homolytic fission?
- Official IB
Chlorine radicals released from CFCs are able to break down ozone, O₃, but not oxygen, O₂, in the stratosphere. What does this suggest about the relative strengths of bonds in the two allotropes?
- Official IB
Do coordination bonds have any different properties from other covalent bonds?
- Official IB
What are the features of benzene, C₆H₆, that make it not prone to undergo addition reactions, despite being highly unsaturated?
5 further official linking questions target Tool, Inquiry or Nature of Science strands and are not drawn as edges on the map:
Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.
- Official IB
What are some of the limitations of the octet rule?
- Official IB
How useful is the VSEPR model at predicting molecular geometry?
- Official IB
How can advances in technology lead to changes in scientific definitions, e.g. the updated IUPAC definition of the hydrogen bond?
- Official IB
What experimental data demonstrate the physical properties of covalent substances?
- Official IB
How can a mixture be separated using paper chromatography or thin layer chromatography (TLC)?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.
- K-02 — Formulae of common ions and compoundsCarry-over
- K-09 — Giant vs simple molecular structures and the property consequencesCarry-over
- N-02 — Orbitals and quantum-derived structure — probability distributions, not pathsNew
- N-09 — Formal charge and Lewis-structure selection — choosing between competing structuresNew
- R-01 — Bonding stops being a set of boxes and becomes a continuumRedefined
- R-02 — "Atoms want a full outer shell" is retired as an explanationRedefined
- R-04 — Shells (2, 8, 8) become subshells and orbitalsRedefined
- R-12 — Intermolecular forces become a properly ranked, universal setRedefined
- R-14 — Dot-and-cross becomes Lewis + VSEPR + formal charge + resonance + hybridisationRedefined