This quiz covers Lessons 6–10: Physical Properties and Classification, Ionic Compounds, Metallic Bonding, Covalent Molecular and Network Solids, and Intermolecular Forces. 20 multiple choice + 3 short answer questions.
Lesson Summaries — Quick Review
Four structural types: ionic (high MP, conducts molten/dissolved, hard+brittle), covalent molecular (low MP, no conductivity, soft), covalent network (very high MP, no conductivity except graphite, extremely hard except graphite), metallic (variable MP, always conducts, malleable). Classification requires at least three properties — no single property is uniquely diagnostic.
Ionic compounds form regular 3D lattices of cations and anions. Lattice energy depends on ionic charge (higher charge → stronger attraction) and ion size (smaller → stronger). Key conductivity rule: ions fixed in solid (no conductivity) → mobile when molten or dissolved (conducts). Hard but brittle due to layer-shift causing like-charge repulsion.
Metallic bonding: cation lattice + sea of delocalised valence electrons. Non-directional → malleable/ductile. Mobile electrons → electrical and thermal conductivity. Metallic bond strength increases with more delocalised electrons per atom and higher cation charge → higher MP. Alloys: foreign atoms (different size) disrupt regular lattice → harder, less malleable.
Covalent molecular: discrete molecules, IMFs between them → low MP (break IMFs, not covalent bonds). Covalent network: continuous covalent bonds throughout crystal → very high MP (break covalent bonds). Classic comparison: CO₂ (molecular, gas) vs SiO₂ (network, MP 1713°C). Graphite: network but conducts and is soft (layer structure, delocalised electrons within layers).
Three IMFs: dispersion forces (all molecules, increases with size), dipole-dipole (polar molecules), hydrogen bonding (N–H, O–H, F–H only — strongest IMF). BP prediction: check for H-bonding first, then polarity, then size. Larger non-polar molecules can exceed small H-bonding molecules in BP. BP anomalies (e.g. H₂O vs H₂S) explained by H-bonding.
Updates as you answer. Short answer marks are self-assessed.
1. Which combination of properties is uniquely consistent with a covalent network solid? L06
2. A substance melts at 801°C, does not conduct as a solid, but conducts well when melted. It is hard and brittle. Its structural type is: L06
3. Graphite is a covalent network solid, yet it is used as a lubricant and as electrode material. Which properties make this possible? L06
4. Why are ionic solids described as "hard but brittle"? L06/L07
5. Magnesium oxide (MgO) has a much higher melting point than sodium chloride (NaCl). The best explanation is: L07
6. A student dissolves potassium iodide (KI) in water and measures excellent electrical conductivity. Which correctly explains this? L07
7. Which pair of ionic compounds would have the most similar melting points? L07
8. In the electron sea model of metallic bonding, what are the charge carriers that conduct electricity? L08
9. Tungsten (W, Group 6, transition metal) has a much higher melting point than potassium (K, Group 1). Using the electron sea model, the best explanation is: L08
10. Which correctly explains why adding tin to copper (to make bronze) increases hardness? L08
11. When liquid iodine (I₂) vaporises, what type of interaction is broken? L09
12. CO₂ is a gas at room temperature (BP −78°C), while SiO₂ is a solid with MP 1713°C. The best explanation for this dramatic difference is: L09
13. Which substance is correctly classified as a covalent network solid? L09
14. Which of the following molecules has ONLY dispersion forces as its IMF? L10
15. Which correctly ranks these substances from lowest to highest boiling point? L10
16. The boiling points of the hydrogen halides are: HF 19°C, HCl −85°C, HBr −67°C, HI −35°C. Which statement correctly explains why HF has a higher BP than HCl, while HI has a higher BP than HBr? L10 HOT
17. A student claims that ethanol (C₂H₅OH, MW 46, BP 78°C) has a higher boiling point than propane (C₃H₈, MW 44, BP −42°C) because "ethanol is a larger molecule with more electrons." Evaluate this claim. L10 HOT
18. A substance has a sharp MP of 1085°C, excellent electrical conductivity in solid and liquid states, and is malleable. When melted, the conductivity remains excellent. What is the charge carrier in the molten state? L06/L08
19. Two substances both have the formula XY₂: one is a gas at room temperature (BP −78°C), the other is a hard solid that melts at 1713°C. The most likely explanation is: L09 HOT
20. A student is given an unknown substance and tests: sharp MP 660°C, excellent conductivity as solid and liquid, malleable, not soluble in water. A second unknown has: BP 100°C, no conductivity in any state, dissolves in water, forms hydrogen bonds. Which pair of structural types correctly identifies both substances? L06–L10 HOT
Attempt all three before checking model answers.
21. A data table lists the following properties of three substances: X (MP 2852°C, no solid conductivity, no liquid conductivity, extremely hard, insoluble); Y (MP 650°C, excellent solid conductivity, malleable, insoluble in water); Z (MP −117°C, no conductivity in any state, soft, slightly soluble in water). Identify the structural type of each substance, classify each as a named structural category, and justify your classification for each with reference to at least two properties. 6 MARKS
22. Explain why the boiling point of ammonia (NH₃, BP −33°C) is much higher than the boiling point of phosphine (PH₃, BP −88°C), even though phosphine is a larger, heavier molecule. In your answer, identify all IMFs present in each substance and explain which type is responsible for the difference. 4 MARKS
23. Compare the electrical conductivity of solid copper (Cu), solid sodium chloride (NaCl), and molten sodium chloride. For each, state whether it conducts, identify the charge carrier (if any), and explain why. 5 MARKS
24. (L06/L07 — 3 marks) Magnesium oxide (MgO) has a melting point of 2852°C, while magnesium chloride (MgCl₂) melts at 714°C. Both are ionic compounds. Explain this difference in melting point using Coulomb's law and ionic charge. 3 MARKS
25. (L07 — 3 marks) Sodium chloride (NaCl) does not conduct electricity as a solid but does conduct when melted or dissolved in water. Explain this behaviour at the particle level for each state. 3 MARKS
26. (L08 — 3 marks) Explain why metals are: (a) good electrical conductors, (b) malleable (can be shaped by hammering), and (c) have relatively high melting points. Refer to the metallic bonding model (electron sea) in each answer. 3 MARKS
27. (L08/L09 — 3 marks) Graphite conducts electricity while diamond does not, even though both are made entirely of carbon atoms. Explain this difference using electron configuration and bonding. 3 MARKS
28. (L09 — 4 marks) Compare the structure and properties of silicon dioxide (SiO₂) and carbon dioxide (CO₂). Both contain a Group 14 element bonded to oxygen, yet their properties are dramatically different. In your answer, address: bonding type, structure, melting point, and electrical conductivity. 4 MARKS
29. (L10 — 4 marks) The boiling points of the noble gases are: He (−269°C), Ne (−246°C), Ar (−186°C), Kr (−153°C), Xe (−108°C). (a) Describe and explain the trend. (b) Noble gases are monatomic and non-polar — what type of IMF exists between them? (c) Explain why Xe has a higher boiling point than He despite both being monatomic and non-polar. 4 MARKS
30. (L10/Synthesis — 4 marks) Hydrogen fluoride (HF, BP 19.5°C), water (H₂O, BP 100°C), and ammonia (NH₃, BP −33°C) all form hydrogen bonds. Explain why their boiling points differ despite all three forming H-bonds. In your answer, compare the number of H-bonds each molecule can form and relate this to boiling point. 4 MARKS
1. D — Very high MP + no molten conductivity + extremely hard = covalent network. A = metal. B = ionic. C = covalent molecular.
2. B — 801°C + no solid conduct + conducts molten + hard/brittle = ionic compound (NaCl).
3. C — Graphite's unique layered structure: delocalised electrons within layers (conducts); weak interlayer van der Waals (soft/lubricant).
4. A — Hard: rigid lattice. Brittle: layer shift → like-charge alignment → repulsion → fracture. Metals avoid this with non-directional bonding and electron sea.
5. C — Higher ionic charge (±2 vs ±1) → stronger electrostatic attraction → higher lattice energy → higher MP. The charge effect dominates.
6. B — Water hydrates and separates K⁺ and I⁻; mobile hydrated ions carry charge. No free electrons involved.
7. D — NaCl and KBr: both ±1 charges, similar ion sizes → similar lattice energies → similar MPs (~800–730°C). All other pairs have large charge or size differences.
8. B — Delocalised electrons are always the charge carriers in metallic conductors. Cations don't move significantly in normal conductivity.
9. C — W contributes ~6 valence electrons per atom (Group 6); K contributes only 1. Higher cation charge + denser electron sea = much stronger metallic bonding = much higher MP.
10. A — Different-sized Sn atoms disrupt regular Cu lattice → impede layer sliding → harder. Not ionic bonds, not electron reduction, not covalent network.
11. C — Boiling iodine breaks dispersion forces between I₂ molecules. The covalent I–I bond within each molecule stays intact. Gas-phase iodine is still I₂.
12. D — Structural type is the key: CO₂ = discrete molecules (weak IMFs between them); SiO₂ = continuous covalent network (must break covalent bonds to melt). Not about bond strength or polarity per se.
13. A — SiC: MP 2730°C + extremely hard + non-conducting = covalent network solid. B = covalent molecular (gas). C = covalent molecular (sugar). D = metal.
14. B — CCl₄ is non-polar (symmetric tetrahedral) → only dispersion. HF: H-bonding. HCl: dipole-dipole + dispersion. NH₃: H-bonding.
15. C — CH₄: dispersion only, very small → lowest (−161°C). HCl: dispersion + dipole-dipole → moderate (−85°C). H₂O: H-bonding → highest (100°C).
16. D — HF is anomalously HIGH (H-bonding). HCl→HBr→HI follows normal dispersion trend (increasing electrons → increasing BP: −85, −67, −35°C).
17. A — Ethanol is higher, but the reason is H-bonding (O–H), not molecular size. The MW difference is tiny (46 vs 44) — dispersion forces are almost identical. The ~120°C BP difference is almost entirely due to H-bonding in ethanol.
18. C — MP 1085°C + conducts solid + malleable = metal (copper, Cu). In molten metal, the electron sea is maintained → delocalised electrons remain the charge carriers (not ions, unlike molten ionic compounds).
19. B — CO₂ (molecular, XY₂ gas) vs SiO₂ (network, XY₂ solid) — the classic same-formula, different-structural-type pair.
20. D — First: MP 660°C (Al's exact MP) + excellent conductivity + malleable → metallic (aluminium). Second: BP 100°C + no conductivity + dissolves in water + H-bonding → covalent molecular (water). No other combination fits all data points.
Q24 (3 marks): MgO contains Mg²⁺ and O²⁻ (both charge magnitude 2). MgCl₂ contains Mg²⁺ and Cl⁻ (charges 2 and 1). By Coulomb's law, electrostatic force ∝ q₁×q₂/r². For MgO: charge product = 2×2 = 4. For MgCl₂: charge product = 2×1 = 2. MgO has twice the charge product → electrostatic attraction between ions is roughly twice as strong per ion pair → lattice energy is much higher → far more thermal energy needed to overcome the lattice and melt → much higher MP (2852°C vs 714°C). Additionally, O²⁻ has a smaller ionic radius than Cl⁻ (O²⁻ ≈ 140 pm; Cl⁻ ≈ 181 pm) → smaller r → even stronger force (force ∝ 1/r²) → further elevating MgO's lattice energy above MgCl₂.
Q25 (3 marks): Solid: Na⁺ and Cl⁻ ions are held in fixed positions in the face-centred cubic lattice by strong electrostatic forces. Ions cannot move — they can only vibrate about their lattice positions. No mobile charge carriers → no current → insulator. Molten: heating supplies enough thermal energy to overcome the lattice energy. Ions break free from their fixed positions and move freely as a liquid. Under an applied electric field, Na⁺ migrates toward the negative electrode and Cl⁻ toward the positive electrode → charge flows → conducts. Dissolved in water: the δ− oxygen of water molecules is attracted to Na⁺ (ion-dipole); the δ+ hydrogen is attracted to Cl⁻. These ion-dipole forces overcome lattice energy — ions are hydrated and dispersed throughout the solution. Hydrated Na⁺(aq) and Cl⁻(aq) ions are free to migrate under an applied field → conducts.
Q26 (3 marks): (a) Conductivity: valence electrons leave their parent atoms and become delocalised throughout the metallic lattice — the "electron sea." These electrons are not fixed to any atom and move freely. Under a potential difference, electrons flow from negative to positive terminal — electrical current. No new bonds need to be broken; the electron sea simply flows. (b) Malleability: metal cations sit in the electron sea. When a mechanical force causes one layer of cations to slide relative to another, the electron sea instantly surrounds the cations in their new positions and continues to hold them together electrostatically. No bonds are permanently broken or formed — the structure is maintained in any shape. Compare to ionic crystals where slip aligns same-charge ions → repulsion → fracture. (c) High melting points: all metal cations are simultaneously attracted to the entire surrounding electron sea (non-directional, 3D electrostatic bonding). Melting requires enough thermal energy so that all cation-electron sea interactions are overcome simultaneously across the lattice — significant energy, giving moderate-to-high melting points. More delocalised electrons (e.g. transition metals with d-electrons) → stronger binding → higher melting points.
Q27 (3 marks): Carbon has 4 valence electrons. Diamond: each C uses all 4 in localised sp³ C–C sigma bonds (tetrahedral). Every electron is paired in a covalent bond — no free or delocalised electrons exist anywhere in the crystal. No mobile charge carriers → electrical insulator. Graphite: each C uses 3 valence electrons in sp² C–C sigma bonds within hexagonal layers. The remaining 1 electron per C is not bonded — it occupies a p orbital perpendicular to the layer and overlaps with p orbitals on all neighbouring C atoms → completely delocalised π-electron system across the entire layer. These π-electrons are mobile (analogous to conduction electrons in a metal) and can carry charge parallel to the layers under an applied field → graphite conducts electricity along its layers.
Q28 (4 marks): CO₂ structure: C (Period 2, small atom) can form stable C=O double bonds. Each CO₂ molecule is discrete: O=C=O, linear, non-polar. Between molecules: only weak London dispersion forces. MP = −78°C (sublimes) — only very small energy to overcome weak dispersion forces. Non-conducting in all states: all electrons localised in C=O bonds; no ions; no free electrons. SiO₂ structure: Si (Period 3, larger atom with available d-orbitals) cannot form stable Si=O double bonds (pπ-pπ overlap is inefficient for larger atoms). Instead, Si forms 4 Si–O single bonds (sp³); each O bridges two Si atoms — infinite 3D covalent network. MP ≈ 1710°C — must break strong Si–O covalent bonds (≈460 kJ/mol) throughout the entire network. Non-conducting: all electrons in Si–O bonds; no free electrons or ions. Key difference: molecular vs network structure arises from the inability of Si to form double bonds, forcing it into an extended covalent network.
Q29 (4 marks): (a) BP increases from He to Xe as atomic number increases. Each successive noble gas has more electrons in a larger, more diffuse electron cloud — the electron cloud is more easily distorted (more polarisable). (b) London (instantaneous) dispersion forces — the only IMF possible for non-polar, monatomic species. An instantaneous fluctuation in electron distribution creates a temporary dipole, which induces a dipole in neighbouring atoms. (c) Xe (Z=54, 54 electrons, electron cloud extending to n=5 shell) vs He (Z=2, 2 electrons, only n=1 shell). Xe's much larger, more polarisable electron cloud creates larger instantaneous dipoles → stronger induced dipoles in neighbours → stronger dispersion forces → more energy needed to overcome them in the liquid → higher BP. He's tiny, barely-polarisable electron cloud creates only extremely weak, fleeting dipoles → almost no IMF → BP just above absolute zero.
Q30 (4 marks): All three molecules contain H bonded to a highly electronegative atom (F, O, N) — H-bond donor. All three have lone pairs — H-bond acceptors. The BP differences arise primarily from the number and strength of H-bonds each molecule can form. H₂O (BP 100°C): O has 2 lone pairs (accept 2 H-bonds) + 2 O–H bonds (donate 2 H-bonds) → up to 4 H-bonds per molecule. Forms an extensive, highly directional 3D H-bond network → strongest and most numerous H-bonds → highest BP. NH₃ (BP −33°C): N has 1 lone pair (accept 1 H-bond) + 3 N–H bonds (donate 3) — can form up to 4 H-bonds geometrically, but lone pair limitation means each molecule can only accept 1. Also N is less electronegative than O (χ: N=3.0 vs O=3.4) → weaker H-bonds (less δ+ on H). Network is less complete than water → intermediate BP. HF (BP 19.5°C): F has 3 lone pairs (accept up to 3 H-bonds) but only 1 H–F (donate 1) → maximum 2 H-bonds per molecule. Despite F being the most electronegative element (strongest H-bonds individually), the HF chain can only form linear H-bond chains (not a 3D network) due to the limited number of donors → fewer H-bonds per molecule than H₂O → lower BP than water but much higher than HCl (which has no H-bonding).
Q21 (6 marks): X is a covalent network solid (1 mark). MP 2852°C is extremely high — only covalent network solids with continuous strong covalent bonds throughout can require this much energy to melt (1 mark). No conductivity in solid or molten state confirms absence of free electrons or mobile ions (excludes metals and ionic compounds). Extremely hard and insoluble further confirms — this substance is MgO... wait, actually MgO is ionic (conducts when molten). Re-check: no liquid conductivity + MP 2852°C + extremely hard + insoluble = covalent network solid. Could be SiC or similar. Correct answer: X = covalent network solid — justified by very high MP AND no conductivity in both states (2 criteria required). Y is a metallic element (1 mark). MP 650°C is consistent with a metal (Mg); excellent solid conductivity indicates delocalised electrons; malleability is exclusive to metals with non-directional bonding. Identified as magnesium (Mg) (1 mark). Z is a covalent molecular compound (1 mark). MP −117°C is very low — only weak IMFs between discrete molecules can be overcome at this temperature; no conductivity and softness confirm molecular structure; slight water solubility suggests polarity. Justified by low MP AND soft/no conductivity (1 mark).
Q22 (4 marks): NH₃ IMFs: dispersion forces AND hydrogen bonding (N–H bonds present; N is electronegative enough to form N–H···N hydrogen bonds with adjacent molecules) (1 mark). PH₃ IMFs: dispersion forces and weak dipole-dipole forces only (P–H bonds present but P is not electronegative enough to form hydrogen bonds; χ(P) ≈ 2.2 vs χ(N) = 3.0) (1 mark). NH₃ has a higher BP because hydrogen bonding between NH₃ molecules is far stronger than the dipole-dipole forces in PH₃ — despite PH₃ being larger and heavier (and therefore having stronger dispersion forces), the hydrogen bonding energy in NH₃ dominates (1 mark). To boil NH₃, the strong N–H···N hydrogen bonds must be overcome — this requires more energy than overcoming the weaker dispersion + dipole-dipole forces in PH₃ → BP of −33°C for NH₃ vs −88°C for PH₃, a difference of 55°C (1 mark).
Q23 (5 marks): Solid copper (Cu): conducts electricity (1 mark). Charge carrier: delocalised electrons (1 mark). Reason: Cu metal has a sea of delocalised valence electrons that are always mobile — they can flow under a voltage without any structural disruption. Solid NaCl: does not conduct electricity. Charge carrier: none mobile in solid state. Reason: Na⁺ and Cl⁻ ions are fixed in the rigid ionic lattice at defined positions — they cannot move to carry charge (1 mark). Molten NaCl: conducts electricity excellently. Charge carrier: mobile Na⁺ and Cl⁻ ions (1 mark). Reason: at 801°C the ionic lattice breaks down — Na⁺ and Cl⁻ ions become free to move independently in the liquid. Under a voltage, Na⁺ migrates to the negative electrode and Cl⁻ to the positive electrode — this ion movement constitutes a current (1 mark).
Tick when you've checked all answers and noted areas to review.