Covalent Compounds: Molecular and Network
Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions — start at whatever level suits you.
Carbon dioxide (CO₂) is a gas at room temperature and has a very low boiling point. Silicon dioxide (SiO₂, quartz) is a hard solid with an extremely high melting point. Both are formed from non-metal atoms joined by covalent bonds. Why do they behave so differently?
Key facts
- The difference between covalent molecular and covalent network (lattice) substances
- Examples of each type and their key properties
- The role of intermolecular forces vs covalent bonds in determining properties
Concepts
- Why covalent molecular substances have low MPs (break IMFs, not covalent bonds)
- Why covalent network solids have very high MPs (must break covalent bonds)
- Why molecular size/polarity affects boiling point within molecular substances
Skills
- Classify a substance as covalent molecular or covalent network from property data
- Explain any covalent property using the correct structural model
- Spot and correct reasoning errors about covalent substances
What changes on melting?
This is the single most important concept in covalent chemistry for IQ2:
Covalent network solids would melt by breaking covalent bonds throughout the lattice. These bonds are very strong → enormous energy needed → very high melting points.
| Feature | Covalent molecular | Covalent network |
|---|---|---|
| Structure | Discrete molecules with IMFs between them | Continuous covalent bond network throughout the crystal |
| What breaks on melting? | Intermolecular forces (IMFs) | Covalent bonds |
| Melting point | Low to moderate (<300°C typical) | Very high (>1000°C typical) |
| Hardness | Soft, easily deformed | Extremely hard (all covalent bonds) |
| Conductivity | None in any state (except some with delocalised π electrons) | None (except graphite — delocalised electrons in layers) |
| Solubility | Polar molecules → dissolve in water; non-polar → don't | Insoluble in all common solvents |
| Examples | H₂O, CO₂, CH₄, C₆H₁₂O₆, I₂ | Diamond, graphite, SiO₂, SiC |
Covalent molecular = discrete molecules held by weak IMFs between them; melting breaks only IMFs → low MP (<300°C). Covalent network = giant 3D lattice of covalent bonds throughout; melting requires breaking covalent bonds → very high MP (>1000°C). Neither conducts (exception: graphite, which has delocalised π-electrons in layers).
Pause — copy the highlighted distinction into your book before moving on.
Two truths and a lie: pick the false statement about molecular vs network covalent substances.
Trends in Boiling Point — IMF Strength
Within molecular substances, BP depends on the strength of IMFs. Stronger IMFs → higher BP.
- Larger molecules: More electrons → stronger dispersion forces → higher BP. e.g. I₂ (BP 184°C) > Br₂ (BP 59°C) > Cl₂ (BP −35°C) — same type of molecule, increasing molecular mass.
- Polar molecules: Dipole-dipole forces in addition to dispersion forces → higher BP than non-polar molecules of similar size.
- Hydrogen bonding: Strongest IMF. Molecules with N–H, O–H, or F–H bonds form hydrogen bonds → anomalously high BP. H₂O (BP 100°C) vs H₂S (BP −60°C): same structure, but O is more electronegative than S, so water forms hydrogen bonds while H₂S doesn't.
We just saw that molecular vs network covalent substances differ fundamentally in what breaks on melting. That raises a question: within the molecular category, why do some substances (like water) have much higher BPs than similar-sized molecules? This card answers it → the type and strength of IMFs between molecules determines BP.
Within molecular substances, BP depends on IMF strength: larger molecules have more electrons → stronger dispersion forces → higher BP (e.g. I₂ > Br₂ > Cl₂ > F₂). Polar molecules add dipole–dipole forces; N–H, O–H, or F–H bonds enable hydrogen bonding. When water boils, O–H···O intermolecular H-bonds break — O–H covalent bonds inside H₂O remain intact.
Add the highlighted IMF rule to your notes before the check below.
Fill the blanks: drag the correct word into each gap.
Within molecular substances, larger molecules have more ___, which produces stronger ___ forces. Water has an anomalously high boiling point because O–H bonds allow ___ bonding. When a molecular substance boils, only the ___ between molecules are overcome — the covalent bonds inside each molecule remain intact.
| Substance | Structure | MP (°C) | Hardness | Conductivity | Notable property |
|---|---|---|---|---|---|
| Diamond | Each C bonded to 4 others in 3D tetrahedral network | 3550 | Hardest natural substance (10 Mohs) | None | Transparent; used in cutting tools |
| Graphite | Each C bonded to 3 others in layers; 1 delocalised e⁻ per C within layers | ~3650 | Soft (layers slide — Mohs 1–2) | Yes (within layers) | Lubricant, electrode material, pencil lead |
| Silicon dioxide (SiO₂) | Each Si bonded to 4 O; each O bridges two Si — 3D network | 1713 | Very hard | None | Sand, quartz, glass (when amorphous) |
| Silicon carbide (SiC) | Similar to diamond — Si and C alternate in tetrahedral network | 2730 | Extremely hard (9.5 Mohs) | None (slightly semiconducting) | Abrasives, cutting discs |
We just saw how IMF type and molecular size control BP in molecular substances. That raises a question: what are the most important covalent network solids, and how do their specific structures explain their properties? This card answers it → diamond, graphite, and SiO₂ each illustrate a different facet of network covalent bonding.
Diamond: each C bonded to 4 others in 3D tetrahedral network → hardest natural substance, MP 3550°C, non-conducting. Graphite: each C bonded to 3 others in layers + 1 delocalised electron → conducts within layers, soft (weak dispersion between layers). SiO₂ (quartz): 3D Si–O–Si network → MP 1713°C, very hard, insoluble.
Pause — write the highlighted network solid examples into your book.
Odd one out: which substance is NOT a covalent network solid?
6. Explain why iodine (I₂, BP 184°C) has a much higher boiling point than fluorine (F₂, BP −188°C), even though both are non-polar covalent molecular substances of the same type. 3 MARKS
7. A student is given data on two unknown substances: Substance X (MP −22°C, no conductivity in any state, dissolves in water) and Substance Y (MP 1713°C, no conductivity in any state, insoluble in all solvents). Classify each substance and explain all of its listed properties in terms of structure and bonding. 5 MARKS
8. Using your knowledge of intermolecular forces, explain why water (H₂O, MW = 18) has a boiling point of 100°C, which is dramatically higher than propane (C₃H₈, MW = 44, BP −42°C), even though propane is a larger molecule. 4 MARKS
We just saw the specific structures of diamond, graphite, and SiO₂. That raises a question: how do you apply all of this in exam questions involving unknown substances or BP comparisons? This card answers it → always state what forces break on melting, and for BP comparisons, name every IMF present before ranking them.
Classify unknowns from data: low MP + non-conducting = covalent molecular; very high MP + non-conducting = covalent network. Always specify what breaks on melting: IMFs only (molecular) vs covalent bonds (network). For BP comparisons: name all IMFs in each molecule (dispersion is in every molecule), then identify which is strongest and why.
Pause — copy the highlighted classification rule into your book before moving on.
Two truths and a lie: which statement is wrong about answering the short-answer questions above?
Worked examples · reveal as you go
Carbon dioxide (CO₂) is a gas at room temperature (BP −78°C). Silicon dioxide (SiO₂) is a hard solid with a melting point of 1713°C. Both consist of a central atom bonded to oxygen atoms. Explain this dramatic difference in properties using your knowledge of covalent molecular and covalent network structures.
CO₂: discrete molecules. Each carbon forms two double bonds with oxygen (O=C=O). The molecules are separate — only weak dispersion forces act between them. CO₂ is a covalent molecular substance.
SiO₂: continuous 3D network. Each Si atom forms four single bonds with oxygen atoms; each O bridges two Si atoms. No discrete molecules exist — the entire crystal is one giant covalently bonded structure. SiO₂ is a covalent network solid.
CO₂: to convert solid CO₂ (dry ice) to liquid or gas, only the weak dispersion forces between CO₂ molecules need to be overcome. The strong C=O covalent bonds within each molecule are NOT broken. Very little energy needed → very low BP (−78°C).
SiO₂: to melt SiO₂, the strong Si–O covalent bonds throughout the entire network must be broken. These bonds have energies of ~450 kJ mol⁻¹ and there are enormous numbers of them — enormous energy required → very high MP (1713°C).
The difference in melting point is not due to C vs Si, or the type of bonds formed — it is due to the structural arrangement: CO₂ forms discrete molecules (weak IMFs between them) while SiO₂ forms a continuous covalent network (strong covalent bonds must be broken to melt).
Water (H₂O) boils at 100°C. Hydrogen sulfide (H₂S) boils at −60°C. Both are covalent molecular substances of similar structure (bent molecules, 2 X–H bonds). Explain the large difference in boiling points.
H₂O: oxygen is highly electronegative (χ = 3.5). The O–H bond is strongly polar. Water molecules can form hydrogen bonds — O–H···O interactions between adjacent molecules. Hydrogen bonding is the strongest type of IMF.
H₂S: sulfur is less electronegative (χ = 2.6). The S–H bond is weakly polar. H₂S cannot form hydrogen bonds (S is not electronegative enough). Only weak dispersion forces and weak dipole-dipole forces act between H₂S molecules.
To boil a liquid, the IMFs between molecules must be overcome. H₂O has strong hydrogen bonds → more energy needed to separate molecules → higher boiling point. H₂S has only weak dispersion and dipole-dipole forces → much less energy needed → much lower boiling point.
In both cases, the O–H covalent bonds within H₂O molecules and the S–H covalent bonds within H₂S molecules remain completely intact when boiling occurs. Only the intermolecular forces between molecules are overcome — H₂O stays as H₂O molecules in the gas phase, and H₂S stays as H₂S molecules. The covalent bonds inside molecules would only break in a chemical reaction (like combustion), not during a phase change.
Click two steps to swap them. Put the explanation of why CO₂ is a gas at room temperature while SiO₂ melts at 1713°C in the correct order.
- SiO₂ forms a continuous 3D covalent network — each Si bonds to 4 O atoms, each O bridges two Si atoms.
- CO₂ consists of discrete O=C=O molecules with only weak dispersion forces acting between them.
- To melt SiO₂, large numbers of strong Si–O covalent bonds throughout the lattice must be broken — requires enormous energy.
- To melt CO₂, only the weak IMFs between molecules need to be overcome — the C=O bonds inside molecules stay intact.
- Conclusion: the dramatic MP difference is due to structural type (molecular vs network), not the type of atoms involved.
Common errors · the 3 traps that cost marks
Misconception to fix
Wrong: All covalent substances have low melting points because covalent bonds are weak.
Misconception to fix
Right: Covalent bonds within molecules are strong, but intermolecular forces between molecules are weak. Covalent molecular substances have low melting points because only IMFs break. Covalent network solids like diamond have extremely high melting points because the entire covalent network must be broken.
Saying "covalent bonds are broken" when water boils
A classic exam-killer: students write that boiling water "breaks the O–H bonds". This would produce free H and O atoms, which is not what happens. Only the hydrogen bonds between molecules are overcome — each H₂O molecule survives intact into the gas phase.
Fix: When boiling a molecular substance, always write "IMFs are overcome / hydrogen bonds break" — never "covalent bonds break".
Quick-fire practice · 5 reps +2 XP per reveal
Classify each as covalent molecular or covalent network: CO₂, SiO₂, diamond, I₂.
When ice melts, what type of force is overcome?
Why is graphite a soft lubricant and an electrical conductor, while diamond is hard and an insulator? Both are pure carbon.
I₂ boils at 184 °C, F₂ at −188 °C. Same molecule type — why the huge difference?
Substance Z has MP 2730 °C, is extremely hard, does not conduct, and is insoluble. What type of substance is it, and explain its hardness and MP in one sentence each.
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Pick your answer, then rate your confidence — that tells the system what to drill next.
Complete the model answer for: "Explain why H₂O (BP 100°C) has a higher boiling point than H₂S (BP −60°C)." Type each missing word, then click Check.
H₂S has only weak forces (and very weak dipole-dipole) because sulfur is not electronegative enough for H-bonding.
Boiling overcomes the between molecules. Stronger IMFs → more energy needed → higher BP.
Therefore H₂O's BP is approximately °C higher than H₂S because hydrogen bonds are much stronger than the IMFs in H₂S.
Q1. 6. Explain why iodine (I₂, BP 184°C) has a much higher boiling point than fluorine (F₂, BP −188°C), even though both are non-polar covalent molecular substances of the same type.
Q2. 7. A student is given data on two unknown substances: Substance X (MP −22°C, no conductivity in any state, dissolves in water) and Substance Y (MP 1713°C, no conductivity in any state, insoluble in all solvents). Classify each substance and explain all of its listed properties in terms of structure and bonding.
Q3. 8. Using your knowledge of intermolecular forces, explain why water (H₂O, MW = 18) has a boiling point of 100°C, which is dramatically higher than propane (C₃H₈, MW = 44, BP −42°C), even though propane is a larger molecule.
📖 Comprehensive answers (click to reveal)
Activity 1
1. Covalent molecular compound. The very low MP of −85°C indicates only weak IMFs between discrete molecules (not covalent bonds, which are strong). Gaseous state at room temperature confirms very weak IMFs. No conductivity is consistent with no free electrons or ions. Slight water solubility suggests mild polarity. This substance could be HCl (BP −85°C).
2. Covalent network solid. MP 2730°C is extreme — only substances where strong covalent bonds extend throughout the crystal can reach this temperature. Extreme hardness confirms a continuous covalent network (all bonds must break to deform). No conductivity eliminates metals and graphite. Insolubility in all solvents is characteristic of network covalent solids. This substance is silicon carbide (SiC).
3. Covalent molecular compound (water, H₂O). A BP of 100°C is high for a small covalent molecule — this is explained by strong hydrogen bonding between H₂O molecules (O is highly electronegative → O–H bonds are strongly polar → O–H···O hydrogen bonds form). Classification is covalent molecular because it consists of discrete H₂O molecules with hydrogen bonds between them. Conducts only when ionised (pure water doesn't conduct well).
Activity 2
Response 1 — Error: Student A is incorrect. When water boils, the H–O–H covalent bonds do NOT break — the water molecules remain intact as individual H₂O molecules in the gas phase. What breaks is the intermolecular hydrogen bonds between adjacent water molecules. Water has a higher boiling point than many other molecules (such as H₂S) because it forms strong hydrogen bonds (O–H···O), not because of stronger intramolecular bonds. The strength of the H–O covalent bond is not relevant to the boiling point — it would only matter if you were breaking water molecules apart chemically.
Response 2 — Error: Student B incorrectly treated SiO₂ as a molecular substance. SiO₂ does NOT consist of discrete molecules — it is a covalent network solid where Si–O covalent bonds extend throughout the entire crystal. CO₂, by contrast, forms discrete O=C=O molecules with only weak dispersion forces between them. The higher MP of SiO₂ is because melting requires breaking strong Si–O covalent bonds (rather than just IMFs as in CO₂). The reason has nothing to do with molecular mass or IMF strength — it is entirely about structural type (network vs molecular).
❓ Multiple Choice
1. B — Boiling breaks intermolecular hydrogen bonds. Covalent O–H bonds remain intact. Gas-phase water is still H₂O molecules.
2. C — CO₂ (molecular, BP −78°C) vs SiO₂ (network, MP 1713°C) is the canonical example of this contrast. A = both ionic; B = both metallic; D = both molecular (comparing IMF type, not structural category).
3. D — Low BP in molecular substances reflects weak IMFs between molecules. The covalent bonds within molecules are not broken on boiling and their strength is irrelevant to the BP.
4. A — Dispersion forces increase with molecular size (more electrons → stronger temporary dipoles → stronger dispersion). These are the only IMFs in non-polar diatomic halogens. No hydrogen bonding; covalent bond strength is irrelevant to BP.
5. B — SiO₂ is a covalent network solid with a continuous Si–O bond network. It does not consist of discrete molecules; it does not dissolve in water; it does not conduct when molten.
Short Answer Model Answers
Q6 (3 marks): Both F₂ and I₂ are non-polar diatomic molecules — the only IMF acting between their molecules is dispersion forces (1 mark). Dispersion forces increase in strength with the number of electrons in a molecule. I₂ has 106 electrons compared to F₂'s 18 electrons — I₂ has many more electrons and a much larger, more easily polarised electron cloud (1 mark). The stronger dispersion forces between I₂ molecules require significantly more energy to overcome during boiling → much higher BP (184°C vs −188°C). No covalent bonds are broken in either case — only IMFs (1 mark).
Q7 (5 marks): Substance X is a covalent molecular compound (1 mark). Its low MP of −22°C indicates only weak IMFs between discrete molecules, which are easily overcome — covalent bonds within molecules are not broken on melting (1 mark). No conductivity in any state confirms no free electrons or mobile ions (1 mark). Water solubility indicates the substance is polar — like-dissolves-like; polar molecular substances can interact with polar water molecules and dissolve. Substance Y is a covalent network solid (1 mark). MP 1713°C requires breaking enormous numbers of strong covalent bonds extending throughout the crystal — consistent with a network solid (this is SiO₂, quartz). No conductivity confirms no free electrons or mobile ions — the strong, directional covalent bonds hold all electrons in place. Insolubility: the Si–O network is too stable and strongly bonded to be disrupted by water molecules (1 mark across Y properties).
Q8 (4 marks): Propane (C₃H₈) is a non-polar hydrocarbon molecule — it cannot form hydrogen bonds or dipole-dipole interactions, so only weak dispersion forces act between propane molecules (1 mark). Despite its larger size (MW 44 vs 18), propane's BP is −42°C because these dispersion forces are still relatively weak (1 mark). Water (H₂O) contains two polar O–H bonds where oxygen's high electronegativity creates a large partial negative charge on O and partial positive charge on H. This allows water molecules to form strong intermolecular hydrogen bonds (O–H···O) with each other (1 mark). Hydrogen bonds are much stronger than the dispersion forces in propane — approximately 5–10× stronger — so far more energy is required to separate water molecules, resulting in a boiling point 142°C higher than propane despite water being the smaller molecule (1 mark).
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