Graphite: why it is soft and why it conducts

How much lead is in a pencil lead? It is grey, it leaves a mark, and it carries the name. There is none. Miners in Cumbria in the fifteen hundreds took a black mineral for lead ore. It was graphite, pure carbon, the same element as diamond. Yet graphite writes and conducts, and diamond does neither. The pencil gets its full answer from one carbon atom, counted carefully.

Start from what graphite shares with diamond: both are giant covalent structures, and both need a very high temperature to melt. The difference is in how the carbon atoms are arranged. Graphite is built in flat layers. Each layer is a sheet of rings, six carbon atoms to a ring, joined edge to edge like the cells of a honeycomb. Pick one carbon atom in the middle of a layer, well away from the edges, and follow the lines leaving it. How many covalent bonds does that middle carbon atom form? Three. Each carbon atom in graphite forms three covalent bonds, not the four it forms in diamond. The layers are stacked on top of each other. Within a layer, the atoms are held by strong covalent bonds. Between the layers, there are no covalent bonds, only weak forces. That last sentence carries most of this video, for Foundation and Higher students alike, so keep its exact words.

Graphite is soft and slippery because there are no covalent bonds between the layers. The layers can slide over each other. Drag a pencil across paper and layers slide off the stack and stay behind on the page. That grey line is layers of carbon atoms. A student writes: graphite is soft because there are no bonds between the layers. It is one word away from right, and that missing word changes everything. Which word is missing from the student's sentence about graphite? Covalent. There are no covalent bonds between the layers. Without that word the sentence is false, because there are weak forces between the layers. Check it for yourself. If there were nothing at all holding the layers to each other, graphite would fall apart into loose sheets. It would not be a solid. One examiner's report, on a graphite question, notes two more slips: "Examiners reported that many candidates are referring to the structure of graphite as being made up of layers of graphene, with the layers being held together in the giant structure by weak intermolecular bonds." Forces and bonds are different things: forces are overcome, bonds are broken. And a layer may be called graphene, but that word explains nothing. The fix is to describe the atoms: three bonds each, flat layers, no covalent bonds between the layers. Now a harder one. Graphite is soft enough to write with, but its melting point is over three thousand degrees C. Those two facts sound as if they cannot both be true. How can graphite be so soft and still melt so high? Say what is between the layers and what is inside them. Sliding happens between the layers, where there are no covalent bonds, only weak forces. Melting happens inside the layers, where an enormous number of strong covalent bonds must be broken. Two different parts of one structure answer the two questions.

Conduction comes from the electron count. A carbon atom has four outer electrons, and in graphite only three of them are used in bonds. The fourth electron from each atom is delocalised. Delocalised means it is not stuck to any one atom or bond, so it can move through the structure. That gives one spare, mobile electron for every carbon atom in every layer. Now, why does graphite conduct electricity, when diamond does not? Graphite's delocalised electrons are free to move through the structure, and they carry the charge. In diamond, all four outer electrons are used in bonds, so none is free. In that one way, graphite is like a metal, which also has delocalised electrons. The video on metals explains theirs. Those electrons come from the atoms in each layer, and their job is carrying charge. They are not what holds one layer to the next. On a Foundation paper, where students were given a diagram of graphite, one examiner's report says: "Students were unable to recall many details about the structure and bonding of graphite, even with the diagram as a prompt. Many did describe layers of carbon atoms, but few mentioned the giant structure or covalent bonding." The report goes on: "Many focussed on the use of graphite in pencils, which was not relevant in the context of the question. Most responses did not mention the number of bonds per atom, or the delocalised electrons, or the lack of bonds between layers." Pencils earned nothing there; the structure was the question. The fix is four things, in order: a giant covalent structure, three covalent bonds per carbon atom, no covalent bonds between the layers, and one delocalised electron per atom. Writing more is not the cure, though. On a Higher paper, a different report found graphite answered well, with one caution: "This question was very well answered, however, many candidates often gave too much information about the arrangement of the carbon atoms and details about delocalised electrons between the layers." Both reports point the same way: the four things and nothing more. A long tour of the hexagons buries them.

A table puts the two forms of carbon side by side, row by row: bonds per carbon atom, layers, delocalised electrons, hard or soft, and conducts or not. Diamond: four bonds, no layers, no delocalised electrons, very hard, does not conduct. Graphite: three bonds, layers, one delocalised electron per atom, soft, conducts. Which single row of that table explains the last two rows? Bonds per carbon atom. Three bonds leave flat layers with no covalent bonds between them, so graphite is soft. Three bonds also leave a spare electron, so graphite conducts. Four bonds give diamond neither. One layer of graphite, peeled away on its own, has its own name: graphene. It has its own video too, the next one in this chapter.

Quick retrieval on graphite now: say each answer before it appears. What is there between graphite's layers, in the exact words? No covalent bonds, only weak forces. Now something new: powdered graphite is used to help metal parts slide past each other. Which feature of its structure does that use? There are no covalent bonds between its layers, so the layers slide over each other readily. Here's another: why does graphite conduct electricity? Each carbon atom uses three electrons in bonds. The fourth is delocalised, free to move through the structure and carry the charge. One more: what is really in a pencil lead, and why does it leave a line? Graphite, which is pure carbon, with no lead at all. Its layers have no covalent bonds between them, so they slide off onto the paper.

If this one's yours now, give it a thumbs-up. Over a whole course, those thumbs show you which videos you can stop worrying about. If the layers are still hazy, skip the thumb and save it. Give it a day and look again, with a pencil in your hand.

Next in the chapter: graphene, fullerenes and carbon nanotubes. Hopefully that's next in your feed, but if not, search for ScholaFly C H oh five, oh seven.

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