A diamond and graphite are both pure carbon — same element, zero difference in chemical composition. Yet a diamond cuts glass while graphite crumbles into your fingers. The difference is entirely structural. This is IQ2 in a nutshell: structure determines properties.
Core Content
Wrong: Hardness and strength mean the same thing when describing materials.
Right: Hardness refers to resistance to scratching or indentation. Strength refers to resistance to breaking under force. Diamond is extremely hard but brittle — it can be scratched by almost nothing yet shatters under impact. These are independent properties.
In IQ1, physical properties told you pure vs mixture. In IQ2, the same properties — measured more carefully — reveal which type of bonding holds the substance together. The five properties are: melting point, electrical conductivity (solid and molten/dissolved), solubility in water, hardness/malleability, and thermal conductivity.
| Structural type | MP/BP | Conductivity (solid) | Conductivity (molten/aq) | Hardness/Malleability | Solubility in water |
|---|---|---|---|---|---|
| Ionic compound | High (hundreds–thousands °C) | None (ions fixed in lattice) | Excellent (mobile ions) | Hard, brittle | Often soluble |
| Covalent molecular | Low (often <200°C) | None | None | Soft, crumbly | Variable (polar dissolves in water; non-polar doesn't) |
| Covalent network/lattice | Very high (>1000°C typically) | None (except graphite) | None (except graphite) | Extremely hard (except graphite) | Insoluble |
| Metallic (element) | Variable (Hg −39°C to W 3422°C) | Excellent | Excellent | Malleable, ductile | Generally insoluble (or reacts) |
Melting point reflects the strength of attractive forces that must be overcome to separate particles. Ionic bonds and covalent network bonds are very strong → very high MPs. Intermolecular forces in covalent molecular substances are weak (dispersion forces, dipole-dipole, hydrogen bonding) → low MPs. Metallic bond strength varies with metal type and structure.
Conduction requires particles that can move and carry charge. Metals have a sea of delocalised electrons — always free to move → conduct in all states. Ionic compounds have ions, but only mobile when lattice is broken (molten or dissolved) → conduct then, not as solid. Covalent substances have no free electrons or ions → don't conduct in any state.
Ionic solids are hard because the lattice is rigid, but brittle — when force is applied, like-charged ions align and repel, shattering the lattice. Metals are malleable because metal atom layers can slide past each other while the electron sea maintains bonding. Covalent network solids are hard because every bond must be broken to deform the structure.
Worked Examples
Activities
1 Substance A: MP = 98°C, conducts well as solid and as liquid, malleable and ductile, insoluble in water.
2 Substance B: MP = 3550°C, does not conduct as solid or molten, extremely hard, insoluble in water.
3 Substance C: MP = −78°C, does not conduct in any state, soft, slightly soluble in water.
Question: "Does sodium chloride conduct electricity?"
Question: "Graphite is a covalent network solid — why is it used as a lubricant and as electrode material?"
Question: "An unknown solid has a very high MP and is extremely hard. What is its most likely structural type?"
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Multiple Choice
5 random questions from a replayable lesson bank — feedback shown immediately
Short Answer
6. Describe the relationship between electrical conductivity and structural type for each of the following: ionic compound, metallic element, and covalent molecular compound. In each case, explain why the substance does or does not conduct electricity. 3 MARKS
7. A student is given three unknown solid substances (X, Y, Z) and measures: X — MP 801°C, no solid conductivity, conducts when dissolved; Y — MP 1085°C, excellent solid conductivity, malleable; Z — MP −115°C, no conductivity in any state, soft. Classify each structural type and justify. 4 MARKS
8. Using your knowledge of structure and bonding, explain why diamond is extremely hard while graphite is soft enough to be used as a pencil lead, even though both consist entirely of carbon atoms. 4 MARKS
1. Metallic element (sodium, Na). MP 98°C is within the metallic range; excellent conductivity as both solid and liquid indicates free delocalised electrons (metallic bonding); malleability and ductility are exclusive to metals.
2. Covalent network solid (diamond, C). MP 3550°C is extremely high — only covalent network solids reach this level. No conductivity in any state confirms absence of free electrons or mobile ions. Extreme hardness is characteristic of a continuous covalent bond network.
3. Covalent molecular compound (dry ice/CO₂ or similar). MP −78°C is very low — characteristic of weak intermolecular forces in a covalent molecular substance. No conductivity in any state confirms no free electrons or ions. Softness is consistent with weak van der Waals forces between discrete molecules.
Response 1 — Error: Student A's answer is incomplete — NaCl does not conduct as a solid, but the claim that it "does not conduct electricity" is incorrect as a general statement. Correct answer: NaCl does not conduct as a solid (ions are fixed in the lattice and cannot move). However, when melted or dissolved in water, the ions become mobile and NaCl conducts electricity well. The reason for non-conduction as a solid is immobile ions, not the absence of free electrons — that reasoning applies to covalent substances, not ionic ones.
Response 2 — Error: Student B incorrectly applied the general rules for covalent network solids to graphite. Graphite is a special case: within each carbon layer, one electron per carbon atom is delocalised and free to move → graphite conducts electricity. Between layers, only weak dispersion forces act → layers slide easily → graphite is soft. Graphite is used as an electrode precisely because it conducts; as a lubricant because its layers slide. Student B should have noted graphite as an exception to the general covalent network rules.
Response 3 — Error: Student C is incorrect to say "definitely ionic". Both ionic compounds AND covalent network solids can have very high MPs and be very hard. The key distinguishing property is conductivity: ionic compounds conduct when molten or dissolved; covalent network solids do not. Without conductivity data, the classification cannot be definitive — it is "ionic compound or covalent network solid".
1. C — Metallic: variable MP + conducts as solid and liquid + malleable/ductile. A = ionic (but "dissolved" is wrong word for metals). B = ionic compound. D = covalent molecular.
2. B — No solid conductivity but conducts dissolved = ionic compound. Metals always conduct as solid; covalent substances never conduct; covalent network doesn't conduct in any state.
3. A — Layer shift in ionic → like charges align → repulsion → fracture. Metal layers slide with electron sea intact → no fracture → malleable.
4. D — S: MP 2030°C (very high) + no conductivity in solid or molten state + extremely hard = covalent network. P = covalent molecular. Q = ionic. R = metal.
5. B — Graphite has delocalised electrons within layers (→ conducts) and weak interlayer forces (→ soft). It is a covalent network solid, not a metal; it doesn't dissolve in water; its melting point is very high.
Q6 (3 marks): Ionic compound: does not conduct as a solid (ions are fixed in the rigid lattice and cannot move to carry charge), but conducts when molten or dissolved (ions become mobile and free to carry charge) (1 mark). Metallic element: conducts in all states because delocalised electrons are always free to move throughout the metal structure, carrying charge regardless of state (1 mark). Covalent molecular compound: does not conduct in any state because there are no free electrons and no ions — all electrons are localised in covalent bonds between specific atoms (1 mark).
Q7 (4 marks): X is an ionic compound — MP 801°C is high but not extreme; no conductivity as a solid indicates fixed ions in a lattice; conducts when dissolved confirms ionic character (mobile ions in solution) (1 mark + 1 justification). Y is a metallic element — MP 1085°C (consistent with copper); excellent conductivity as a solid indicates delocalised electrons; malleability is exclusively metallic (1 mark + 1 justification). Z is a covalent molecular compound — MP −115°C is very low, indicating only weak intermolecular forces between discrete molecules; no conductivity in any state confirms no free electrons or ions; softness is consistent with weak van der Waals forces (1 mark). [Note: 4 marks allocated across the three classifications and their justifications]
Q8 (4 marks): In diamond, each carbon atom forms four strong covalent bonds to four other carbon atoms in a continuous 3D tetrahedral network (1 mark). Every electron is localised in a covalent bond — there are no free electrons or weak points in the structure. To scratch or deform diamond, covalent bonds must be broken; the energy required is extremely high, making diamond the hardest natural substance (1 mark). In graphite, each carbon atom forms three covalent bonds within a flat hexagonal layer, with one remaining electron delocalised within the layer (1 mark). Between layers, only weak dispersion (van der Waals) forces act. These weak interlayer forces are easily overcome — layers slide past each other under small forces. The layers detach and transfer to surfaces (paper), which is why graphite writes. The low force needed to slide layers apart is what makes graphite soft (1 mark).
Return to your Think First response. You should now be able to explain diamond and graphite using bonding and structure:
Climb platforms, hit checkpoints, and answer questions on metallic, ionic, covalent molecular and covalent network solids. Quick recall from lessons 1–6.
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Use the PDF for classwork, homework or revision. It includes key ideas, activities, questions, an extend task and success-criteria proof.
Diamond is one of the hardest known substances and does not conduct electricity. Graphite is soft and slippery and conducts electricity well. Both are made of pure carbon. How can two substances with the same element have such different properties?
Before reading on, write your best answer. What must be different about the way carbon atoms are arranged in diamond compared to graphite?