Why does graphite conduct electricity but diamond does not? Both are forms of pure carbon, but the atoms are connected differently. In diamond, each carbon atom makes four strong covalent bonds in a three-dimensional giant structure; there are no freely moving electrons available for ordinary electrical conduction. In graphite, each carbon bonds to three neighbouring carbon atoms in flat layers, leaving electrons delocalised within the layers. These mobile electrons allow graphite to conduct electricity. The layers can also slide relatively easily, making graphite soft and useful as a lubricant, while diamond’s three-dimensional bonding makes it exceptionally hard.
For Hougang parents and students searching for Secondary 3 Chemistry bonding, graphite versus diamond, giant covalent structure, electrical conductivity, O-Level Pure Chemistry, 2027 SEC G3 Chemistry K324 or properties of materials, the fastest exam answer is structure → particles and bonds → mobile charge carriers → observed property. Do not write “graphite conducts because it is a metal” or “diamond is hard because carbon is hard”. Carbon is a non-metal in both. The arrangement of atoms explains why the properties differ. The Hougang ionic versus covalent bonding guide covers the earlier bonding vocabulary; this article takes the next step into structure and properties.
Two materials, one element, completely different personalities
Imagine a student finishing school near Hougang Central and noticing the graphite in a pencil, then seeing diamond jewellery in a shop window. How could the dark, easily rubbed-off pencil material and a famously hard gemstone both be made from carbon? That question has a lovely answer because it reveals what Chemistry really studies: the same chemical elements can behave differently when their atoms are arranged differently.
The answer is not a mysterious extra ingredient secretly added to the diamond. Both diamond and graphite are allotropes of carbon—different structural forms of the same element. A diamond crystal is organised as a connected three-dimensional network. Graphite consists of extended covalently bonded sheets with weaker attractions between layers. Changing the architecture changes what the material can do.
This is a particularly good Secondary 3 learning moment. A student may memorise the words “giant covalent” and “strong covalent bonds” yet give the same explanation for every property. The exam question then asks why graphite conducts electricity or why it can leave a mark on paper. The words are familiar; the connections between them are missing. Let us build those connections calmly.
What is a giant covalent structure?
A giant covalent structure contains a very large network of atoms joined by covalent bonds rather than separate small molecules. The word giant refers to the extended network, not necessarily to the visible size of a rock or gemstone. Diamond, graphite and silicon dioxide are standard school examples, although their networks differ substantially.
Compare this with water, H₂O, or methane, CH₄. In a pure sample of water or methane, discrete molecules interact with one another; the bonds within each molecule are not the same thing as the forces between different molecules. When a simple molecular substance melts or boils, those intermolecular interactions are overcome while the ordinary intramolecular covalent bonds remain. In a giant covalent network, strong bonding extends throughout the solid.
| Structure type | Particle organisation | Typical inference |
|---|---|---|
| Simple molecular, e.g. methane | Individual covalent molecules with intermolecular attractions | Often lower melting and boiling temperatures |
| Giant covalent, e.g. diamond | Extended covalent network in three dimensions | Very hard; high-temperature resistance |
| Giant covalent, e.g. graphite | Strongly bonded carbon sheets with weaker interlayer attractions | Conductive along sheets; layers slide |
| Giant ionic, e.g. sodium chloride | Positive and negative ions in a lattice | Conducts when molten or suitably dissolved, not as a normal solid |
| Metallic, e.g. copper | Positive metal centres with delocalised electrons | Conducts electricity in solid form |
The important habit is to identify what is mobile and which forces must be overcome. The labels help us select an explanation, but the explanation must still describe the mechanism. A material is not electrically conductive merely because it has strong bonds.
Diamond: every carbon atom is connected in three dimensions
In the introductory model, every carbon atom in diamond makes four covalent bonds to neighbouring carbon atoms, arranged approximately tetrahedrally. That pattern continues through a giant network. Pressing or cutting a diamond requires disrupting a very strong structure in particular directions, which accounts for its exceptional hardness.
A student sometimes writes “diamond is strong because it has more carbon atoms”. That cannot be the answer: a chunk of graphite may contain many more carbon atoms than a tiny diamond. What matters is the connectivity of the atoms and how the network responds to forces. Four bonds per carbon in a three-dimensional structure is the school-level mechanism.
Diamond is also an important reminder that thermal conductivity and electrical conductivity are different properties. Pure diamond is an excellent conductor of heat in many circumstances because vibrations transmit energy effectively through its crystal, yet ordinary pure diamond is a poor conductor of electricity. If your child says every heat conductor must conduct current, this comparison exposes the mistake.
Graphite: strong sheets, easier sliding, mobile electrons
Each carbon atom in graphite forms three strong covalent bonds to neighbouring atoms within a roughly flat hexagonal sheet. The fourth valence electron contributes to delocalised electronic states extending within the structure. These mobile electrons support electrical conduction, especially along the layers.
The sheets are held together by much weaker interlayer forces than the strong covalent bonds inside a sheet. Under suitable shear forces the sheets can slide over one another or individual flakes can separate. This accounts for graphite’s lubricating action and why pencil graphite can transfer material to paper. It does not mean the covalent bonds inside each sheet are weak.
It is important to avoid saying “graphite has weak covalent bonds”. That line is both tempting and incorrect. Within a layer, the carbon atoms are strongly bonded; between layers, the interactions are comparatively weak. The correct answer names the location of each type of interaction. The material’s behaviour follows from their arrangement.
| Property | Diamond | Graphite |
|---|---|---|
| Carbon covalent bonding | Four neighbours per carbon | Three neighbours per carbon in sheets |
| Overall structure | Three-dimensional network | Layered giant covalent structure |
| Electrical current | Ordinary pure diamond is an electrical insulator | Mobile delocalised electrons allow conduction |
| Scratch / cutting behaviour | Exceptionally hard | Layers can separate and slide |
| Thermal behaviour | Transfers heat efficiently through its crystal | Thermal properties depend on direction and form |
| Composition | Carbon element | Carbon element |
Worked question 1: complete an exam explanation
Question: Explain why graphite conducts electricity but diamond normally does not, even though both are forms of carbon.
Model response: In graphite, each carbon atom bonds to three others, so electrons are delocalised and can move through the layers to carry electrical charge. In diamond, each carbon atom forms four covalent bonds in a three-dimensional network, leaving no freely mobile electrons for ordinary conduction. Therefore graphite conducts electricity, while pure diamond generally does not.
Notice the three parts: the bonding arrangement, the charge carrier, and the resulting property. An answer such as “graphite has electrons and diamond has none” is wrong because both materials contain electrons. The difference is whether electrons are sufficiently mobile to conduct electrical current.
Worked question 2: why does graphite leave a pencil mark?
Question: Graphite is used in the writing cores of ordinary pencils. Explain this use in terms of structure and bonding.
Model response: Graphite has layers of carbon atoms joined by strong covalent bonds within each sheet, with weaker attractions between layers. The layers can slide and small fragments can transfer to paper as the pencil moves, leaving a visible mark.
A useful follow-up is to ask why the pencil can leave material behind without becoming a puddle. The graphite is not melting during normal writing. Tiny fragments detach mechanically. That distinction connects Chemistry with everyday observations and avoids conflating separation of layers with a phase change.
Worked question 3: compare graphite with sodium chloride
Question: Both graphite and solid sodium chloride have extended structures. Why can graphite conduct in solid form while solid sodium chloride normally cannot?
Model response: Graphite has mobile delocalised electrons that carry charge along its layers. In solid sodium chloride, the sodium and chloride ions occupy fixed positions in an ionic lattice and cannot move freely. When molten or suitably dissolved, the ions become mobile and can carry current.
The phrase “giant structure” does not by itself predict electrical behaviour. We must ask which charged species can move. Graphite uses electrons, an ionic melt or aqueous solution uses ions, and a metal uses delocalised electrons. This is one of the most transferable structure-property questions in school Chemistry.
Worked question 4: the same carbon atom, a different structure
Question: A student says, “Graphite is soft because carbon atoms form weak covalent bonds, but diamond is hard because its covalent bonds are strong.” Correct the explanation.
Correction: Covalent bonds between carbon atoms are strong within both graphite sheets and diamond’s network. Diamond has strong bonding extending through a three-dimensional structure. Graphite has strong bonds within layers but much weaker attractions between those layers, allowing sliding. The relative ease of sliding does not mean the in-layer covalent bonds are weak.
This is a classic hidden-gap diagnosis. The student knows that “soft” must be linked to “weak” somehow, but assigns the weakness to the wrong interaction. Instead of telling the child to memorise a longer paragraph, ask them to mark which bonds or forces lie within the sheet and which act between layers.
Why this matters when a question gives an unfamiliar substance
Suppose a hypothetical material is described as containing an extended network with no mobile ions or electrons. It has a high melting point and is hard. The student might reasonably consider a giant covalent structure, but the evidence does not uniquely identify diamond. Silicon dioxide is also a giant covalent material with similar broad characteristics, although details differ.
A strong Chemistry answer stays within the provided evidence. If the question includes layers, weak interlayer attractions and electrical conduction through mobile electrons, graphite becomes a much stronger match. If it describes four bonds per carbon and no mobile charge carriers, diamond fits. Students should explain why a candidate is supported and which alternatives are less consistent.
Four small misconceptions to catch before the examinations
| Misconception | Why it fails | Correct connection |
|---|---|---|
| All forms of carbon have the same properties | Ignores different allotropes | Structure changes properties even with the same element |
| Graphite is a metal because it conducts | Electrical conduction is not proof of metallic identity | Graphite is non-metallic carbon with delocalised electrons |
| Diamond contains no electrons | Confuses absence of mobile carriers with absence of electrons | Electrons are present but not freely mobile |
| Every giant covalent solid is insulating | Overgeneralises diamond’s behaviour | Graphite is a notable conducting example |
| Weak forces are the same as weak covalent bonds | Conflates bond types and locations | Separate in-layer bonds from interlayer attractions |
A parent can turn these into five very short discussion questions. Ask your child to choose the wrong line and explain it in their own words. You do not need a crystal model at home to hear whether the learner understands how the structure produces its properties.
A short revision routine after a Secondary 3 test
- Draw two simple word models: “four bonds in 3D” and “three bonds in sheets”. These are memory aids rather than required examination-scale atomic diagrams.
- For each, name the relevant mobile charge carriers—or their absence.
- Explain why one is hard and why the other’s layers slide.
- Compare graphite with one giant ionic substance and one metallic substance.
- Close the notes and answer one changed question, such as “Which material could be used as an electrode, and why?”
A student who can answer these in a new order is beginning to build a usable structure-property model. A student who knows only the exact paragraph from the worksheet may fail when the question asks about conductivity first or changes the material name. Train the transferable link rather than the fixed position of each sentence.
What Singapore’s 2027 SEC G3 syllabus actually asks
In the official 2027 SEAB K324 Pure Chemistry syllabus, section 3.4 includes giant covalent substances such as diamond, graphite and silicon dioxide, and explicitly requires comparing diamond and graphite to explain properties including electrical conductivity, cutting action and lubrication. It also states that candidates are not required to draw the structures in that specific part of the syllabus. Understanding and comparison are the main jobs here.
This is why asking a student to reproduce elaborate three-dimensional crystal diagrams from memory may be an inefficient first revision step. Accurate schematic understanding can still support explanation, but the published scope is clear about the drawing requirement. Check the student’s own school subject and examination year, including differences between Pure Chemistry K324 and Combined Science K326/K328, before making an assessment-specific plan.
Chemistry in Hougang: useful examples without pretending they were tested
Pencil marks, jewellery, building materials and electrical devices offer everyday entry points. They do not prove the exact composition of a particular structure in Hougang. A visible black surface could be paint, soot, plastic or something else; a silver-looking fitting could be one of many alloys. Use these observations to ask questions, then consult evidence or a trusted reference before naming the material.
The broader How Materials Age in Hougang and eduKateSingapore materials guide connect atomic structure to engineering performance. Here the educational focus stays on the student’s ability to explain why a substance behaves as it does, using only claims supported by its structure.
What a Chemistry tutor should see in a three-student class
One learner may say “diamond is ionic”; another may correctly identify covalent bonding but cannot explain conductivity; a third may think graphite has weak covalent bonds in its layers. Their incorrect answers look similar if the teacher checks only the final grade. Effective small-group tutoring inspects the first wrong link and then asks each student to defend a different comparison.
eduKate’s immutable three-pax tutorial reference describes the small-group approach at its centre near Sixth Avenue MRT. This Hougang Chemistry page is a learning resource for local families, not a claim of a physical eduKate Chemistry classroom in Hougang. Parents can use the Hougang Chemistry tuition guide to decide whether travel and support format fit their child’s needs.
Frequently asked questions
Does diamond conduct electricity when heated?
Ordinary pure diamond is an electrical insulator in the school model. Specially doped forms and unusual conditions can change electrical properties, but those cases are not the basis of the familiar Secondary 3 comparison.
Why does diamond conduct heat well but not electrical current?
Heat can be transmitted efficiently by vibrations in the strongly bonded crystal, whereas electrical conduction requires mobile charged carriers that pure diamond does not readily provide.
Is graphite a compound?
No. Graphite is an allotrope of the element carbon. It contains carbon atoms arranged in a particular giant covalent structure.
Why is graphite soft if it contains strong covalent bonds?
The strong bonds are mainly within layers. Weaker interactions between layers allow sliding or separation, which explains the material’s soft, lubricating behaviour.
Must I draw the entire diamond and graphite lattices for SEC G3 K324?
The 2027 syllabus specifically says candidates will not be required to draw their giant covalent structures for the listed comparison. You still need to describe bonding, structure and the property being explained.
One strong explanation is worth more than a page of slogans
Ask tonight: “Where can charge move in graphite, and why can’t it move that way in diamond?” If the learner can answer by linking carbon bonding to delocalised electrons, they have learned something that will help with metals, ionic salts, polymers and future questions about materials. That is Chemistry becoming a connected system.
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