Graphite is the kind of material that shows up in an exam question and quietly dares you to overcomplicate it.
It looks like a non-metal. It writes in your pencil. It feels soft. And yet, in IB Chemistry, it’s the famous exception that conducts electricity. That surprise is exactly why examiners love it: graphite forces you to connect bonding, structure, and properties in one clean explanation.
By the end of this article, you’ll be able to explain why graphite conducts electricity in the language IB Chemistry marks reward: layers, sp2, delocalised electrons, and directionality.
The exam checklist (learn this once)
When a question asks why graphite conducts, run this quick checklist (it’s pure IB Chemistry structure-to-property logic):
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Graphite is a giant covalent structure made of layers
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Each carbon forms three sigma bonds in a hexagonal sheet
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Carbon is sp2 hybridised, leaving one electron not used in sigma bonding
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That electron becomes delocalised in a pi system across the layer
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Delocalised electrons can move and carry charge (so graphite conducts)
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Conductivity is strong within layers, weak between layers
If you want the broader bonding context, review Bonding as a continuum (S2.4.1 notes) and the Covalent network structures notes (S2.2.7).
Graphite’s layered structure (what you must describe)
In graphite, carbon atoms are arranged in flat, hexagonal layers. Within a layer, each carbon atom is covalently bonded to three neighbouring carbon atoms. Those bonds are strong and form an extended network, which is why graphite has a high melting point (typical giant covalent behaviour).
But the layers do not connect with strong covalent bonds. They are held together by weak London dispersion forces, so layers can slide. That’s why graphite is soft enough for pencils and slippery enough to work as a lubricant.
This “strong in-plane, weak between planes” structure is the stage. The real plot twist is the electron that doesn’t get assigned a seat.
For a fast giant-structure comparison that often appears in IB Chemistry, see Allotropes of carbon explained.
The real reason graphite conducts electricity: delocalised electrons
Each carbon has four valence electrons. In graphite, three are used to form three sigma (σ) bonds. That leaves one electron per carbon.
Instead of being trapped in a single bond, that remaining electron becomes part of a delocalised pi (π) electron system spread across the sheet. Because these electrons are mobile, they can move when a potential difference is applied. Mobile charged particles are the requirement for electrical conductivity in IB Chemistry, and graphite has them.
This is the sentence that gets marks:
Graphite conducts electricity because it has delocalised electrons that are free to move along its layers and carry charge.
To reinforce this structure-property pattern, it helps to contrast with metals. Metals conduct due to a sea of delocalised electrons in metallic bonding. Graphite is different (it’s a non-metal), but the key similarity is still delocalised electrons. Read Metallic bonding explained for IB Chemistry for that comparison.
sp2 hybridisation: the IB Chemistry explanation behind the scenes
To make the delocalisation feel less like magic, connect it to hybridisation.
In IB Chemistry, graphite is described as sp2 hybridised:
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One s orbital mixes with two p orbitals to make three sp2 orbitals
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These three sp2 orbitals form three sigma bonds in a trigonal planar arrangement (120°)
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One unhybridised p orbital remains on each carbon
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The p orbitals overlap sideways to form a delocalised π system across the layer
If hybridisation still feels slippery, review Hybridization explained for IB Chemistry and the syllabus-linked 14.2 Hybridization notes.
Why diamond does not conduct (your comparison paragraph)
Diamond is also carbon, so the comparison is irresistible in IB Chemistry.
In diamond, each carbon forms four sigma bonds in a 3D tetrahedral network (sp3 hybridisation). Every valence electron is locked into a covalent bond. There are no delocalised electrons and no mobile ions. With no mobile charge carriers, diamond does not conduct electricity.
This contrast is a gift in exam questions: same element, different bonding arrangement, completely different properties.
Conductivity is directional (the detail that wins extra marks)
Graphite conducts within layers much better than between layers.
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Within a layer, the delocalised π electrons can move relatively freely
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Between layers, weak intermolecular forces keep the layers close but do not create an efficient path for electron movement
So if a question mentions “conducts along layers but not between,” you now know exactly which words to use.
For practice in the exact style IB Chemistry expects, use the S2.4 From models to materials Questionbank.
Conclusion: turn graphite into easy marks
Graphite conducts electricity for one beautifully testable reason: in IB Chemistry, each carbon in graphite forms three sigma bonds (sp2), leaving one electron to join a delocalised π system that can move along the layers and carry charge. Add the layered structure and weak interlayer forces, and you also explain softness and directional conductivity.
To lock this in before exams, use RevisionDojo’s Chemical bonding and structure topic hub, drill the Structure 2 flashcards, and test yourself in the Questionbank with instant feedback from AI Chat, grading tools, and modelled explanations. When graphite appears again, it won’t feel like trivia. It’ll feel like a predictable consequence of structure--and that’s the mindset that wins IB Chemistry marks.