--- title: "Giant Covalent Structures: Diamond, Graphite and Graphene" description: "Some covalent substances never stop bonding. Explain diamond, graphite, graphene and silicon dioxide from how many bonds each atom makes and what is left over." canonical: https://lightmysky.com/learn/science/giant-covalent-structures-diamond-graphite-and-graphene-mt_YTEj9srwPG source: https://lightmysky.com/learn/science/giant-covalent-structures-diamond-graphite-and-graphene-mt_YTEj9srwPG.md retrieved: 2026-09-02 --- > **Agent view.** This is the Markdown twin of the page, for tools and assistants. > When to use this site, and the call that answers each job: https://lightmysky.com/agent-instructions.md > API description (OpenAPI 3.1): https://lightmysky.com/openapi.json · Authentication: https://lightmysky.com/auth.md > Pricing: https://lightmysky.com/pricing.md · Catalog: https://lightmysky.com/llms.txt · Full catalog: https://lightmysky.com/llms-full.txt > Every machine-readable file on this domain: https://lightmysky.com/.well-known/ai-catalog.json > Ask for Markdown with `Accept: text/markdown`, a `.md` address, or `?mode=agent`. # Giant Covalent Structures: Diamond, Graphite and Graphene Some covalent substances never stop bonding. Explain diamond, graphite, graphene and silicon dioxide from how many bonds each atom makes and what is left over. Subject: Science · Area: Matter & Materials · Ages 15 to 16 Page: https://lightmysky.com/learn/science/giant-covalent-structures-diamond-graphite-and-graphene-mt_YTEj9srwPG ## Ready when they can - Explains diamond's hardness and very high melting point from four covalent bonds per carbon atom - Explains why graphite conducts and why its layers slide, using the fourth electron and the weak forces between layers - Compares graphene and a single graphite layer and gives one use that follows from the structure - Explains why giant covalent substances do not dissolve in water and, apart from graphite, do not conduct ## Lesson: Giant covalent structures: diamond, graphite and graphene Last week the useful split was strong bonds inside a molecule and weak pulls between molecules. That worked because the molecules were small and separate. Three things on the bench break the pattern. The phone's glass screen, the black powder inside its battery, and the cutting disc in the tool drawer are all held together by covalent bonds, and not one of them melts anywhere near a low temperature. Two of them are the same element, carbon, arranged two different ways, and they could hardly be less alike. Carbon makes four bonds. In diamond every carbon atom spends all four on four other carbon atoms, and each of those does the same, and so on with no edge and no end. There is no molecule to point at. Breaking a diamond means breaking covalent bonds rather than pulling molecules apart, so it stays solid past 3500 degrees and is the hardest natural substance known. That is why the cutting disc has diamond grit along its rim. And with all four outer electrons tied into bonds, none is free to move, so diamond does not conduct. *(drawing: Four bonds each, repeated in every direction. There is nothing here that counts as a molecule.)* Graphite is the same element doing something else. Each carbon spends only three of its four bonds, joining into flat sheets of hexagons. The fourth outer electron is not tied into any bond. It is delocalised, meaning it is free to move along the sheet, and that is what lets graphite conduct when diamond and silicon dioxide cannot. The sheets themselves are only weakly attracted to each other, so they slide apart under a fingertip. That sliding is what a pencil line is, and it is why graphite works as a dry lubricant. *(drawing: Three bonds used, one electron spare, and sheets that barely hold on to each other.)* **Example.** Graphene is a single sheet of graphite, one atom thick. It keeps the strong hexagonal bonding and the free electrons that carry a current. What it loses is the stack, so there are no weak layers left to slide. That leaves something very strong for its weight and nearly transparent, which is why it turns up in flexible screens. Silicon dioxide is the other case on the bench. Each silicon bonds to four oxygens and each oxygen to two silicons, giving another endless network. That is sand, and the glass on the phone. Line them up and the family shows. Every one of them stands up to very high temperatures, because getting them apart has to break covalent bonds. None of them dissolves in water: there are no separate molecules for water to surround and no ions for it to pull away, only one enormous connected structure. And apart from graphite and the single sheet of it that is graphene, none of them conducts, because no electron in them is free to move. Those two are the exceptions for one reason: the spare fourth electron on each carbon is free to travel along the sheet. **Recap.** Some covalent substances never stop bonding. Diamond spends all four of carbon's bonds in every direction, which makes it hard and keeps it solid past 3500 degrees. Graphite spends three, leaving a free electron that conducts and sheets that slide. Graphene is one of those sheets on its own. Silicon dioxide is the same idea with two elements. All of them stand up to high temperatures and none dissolves, and only graphite and graphene conduct. ## Practice 24 questions on this page, each with its working shown. ## Needs first - [Covalent Bonding: Shared Pairs and Simple Molecules](https://lightmysky.com/learn/science/covalent-bonding-shared-pairs-and-simple-molecules-mt_fn5Dq6qUjP) ## Opens up - [Metallic Bonding and Alloys](https://lightmysky.com/learn/science/metallic-bonding-and-alloys-mt_yXOGOwHmZ4)