One element, two opposite materials
Put a diamond next to the tip of a pencil. The diamond is clear and sparkling. It is the hardest natural material, so hard that diamond-tipped blades are used to cut stone. The pencil "lead" is mostly graphite. It is grey, soft and slippery. It rubs off onto paper at the lightest touch.
Now test them with a battery and a bulb. Diamond does not conduct electricity at all. Graphite does.
Here is the surprise. Both are pure carbon. The same atoms, with the same six protons and the same six electrons, make both materials. So the difference cannot be in the atoms. It must be in how the atoms are joined together.
The lectures on metals said that diamond and graphite are giant lattices of a single element, like a metal, but held together by covalent bonds. They also said that diamond has no mobile electrons, while graphite, surprisingly, does. This lecture explains both statements, and every property that follows from them.
Giant covalent structures
You have met covalent bonds in small molecules, such as water and methane. Each molecule is a separate unit. Between molecules there are only weak intermolecular forces, such as London forces. That is why so many molecular substances melt and boil at low temperatures: heating only has to pull the molecules apart, not break them.
Carbon can do something different. A carbon atom has four outer electrons, so it can form up to four covalent bonds. If each carbon bonds to other carbons, and those bond to more carbons, the bonding never has to stop. The result is not a small molecule. It is one huge network of atoms joined by covalent bonds, as large as the crystal itself.
A giant covalent structure is a solid in which covalent bonds join every atom to its neighbours in a continuous network, with no separate molecules.
Some books call these "macromolecular" or "network covalent" structures. The idea is the same. A whole diamond is, in a sense, one molecule. A diamond of one carat, which is 0.2 g, contains about carbon atoms, all bonded into a single network.
Diamond and graphite are both giant covalent structures of carbon. Different forms of the same element, in the same physical state, are called allotropes. Diamond and graphite are allotropes of carbon.
Diamond: four bonds in four directions
In diamond, every carbon atom forms four single covalent bonds, one to each of four other carbon atoms. You know from VSEPR theory what four bonding regions do: they spread as far apart as possible, into a tetrahedral arrangement with angles of 109.5°.
Each of those four neighbours is itself bonded to four carbon atoms, and so on. The tetrahedra link up in all three directions of space. There is no weak direction and no weak layer. Every atom is held in place by four strong bonds pointing four different ways.
Every property of diamond follows from that picture.
- Extremely hard. To scratch diamond, you must break strong C–C covalent bonds. Every direction you push in is braced by bonds. This is why diamond is used on the edges of saws and drills that cut stone and glass.
- Does not conduct electricity. Each carbon uses all four of its outer electrons in its four bonds. Every electron is held in a bond between two particular atoms. There are no delocalised electrons and no ions, so nothing carries a charge through the crystal.
- Insoluble in water and in other common solvents. Dissolving would mean separating the atoms, and that would mean breaking covalent bonds. A solvent's attractions are far too weak to do that.
- Very high melting point. To turn the solid into a liquid, you must break covalent bonds throughout the crystal. Textbooks often give about 3550 °C. In fact, at normal pressure, strongly heated diamond turns into graphite rather than melting. Carbon then turns directly into a gas at about 3640 °C. Either way, the temperatures are far higher than for almost any other substance.
Diamond has one more surprising property. It conducts heat better than any other natural material. The lecture on metals showed that heat can travel through a solid in two ways: by moving electrons, and by vibrations passed from atom to atom. Diamond has no moving electrons, but its stiff, light, tightly bonded lattice passes vibrations along extremely quickly. So a good thermal conductor is not always a good electrical conductor.
Graphite: three bonds and one free electron
In graphite, each carbon atom forms only three covalent bonds, to three neighbours in the same flat plane. Three bonding regions spread out to 120°, a trigonal planar arrangement. The atoms join into flat rings of six, and the rings join into an endless flat sheet, like a honeycomb.
Count the electrons. Each carbon has four outer electrons, but it uses only three of them in its three bonds. The fourth electron is not held between any two atoms. It is delocalised: free to move anywhere within its layer. In this one way, a graphite layer behaves like a metal.
The sheets stack on top of each other. Between the sheets there are no covalent bonds at all, only London forces. Look at the two distances in the figure. Within a layer, neighbouring atoms are 142 pm apart. Between layers, the gap is 335 pm, more than twice as far. Such a large gap is the sign of a weak attraction.
So graphite has two kinds of bonding at once: very strong covalent bonds in two directions, inside each layer, and very weak forces in the third direction, between layers. Every property of graphite follows from that combination.
- Soft and slippery. The layers can slide over each other, because only weak London forces hold them together. When you write, layers of graphite slide off the pencil and stay on the paper. Pencil leads are made of powdered graphite mixed with clay, and more clay makes a harder pencil. Graphite powder is also used as a dry lubricant.
- Conducts electricity along the layers. The delocalised electrons can move along a layer when a potential difference is applied. Graphite is used for electrodes in electrolysis for this reason. Conduction across the layers is much poorer, about 1000 times lower, because the electrons cannot easily jump the large gap between sheets.
- Grey, shiny and opaque. Like the electrons in a metal, the delocalised electrons absorb and re-emit visible light.
- Insoluble in water. For the same reason as diamond: dissolving would need covalent bonds to break.
Graphite is soft enough to rub off onto paper. Does that mean the bonds in graphite are weaker than the bonds in diamond?
Show the answer
No. The covalent bonds inside a graphite layer are actually a little shorter than those in diamond (142 pm against 154 pm), and shorter bonds between the same atoms are usually stronger. Graphite is soft because of the weak forces between the layers, not because of weak bonds. Softness tells you about the weakest link in a structure. The layers slide, but each layer stays whole.
The melting point puzzle
If the layers are held together so weakly, you might expect graphite to melt easily. Make a prediction before you read on.
The layers in graphite are held together only by weak London forces. What does this tell you about the temperature needed to melt graphite? Choose before you read on.
Graphite does not melt until well above 3000 °C. At normal pressure, it turns directly into a gas at about 3640 °C. Separating the layers is easy, but that does not melt the solid. To turn graphite into a liquid or a gas, the strong covalent bonds inside every layer must break.
This is one of the most common exam traps, so state the rule carefully. A property depends on which bonds the change has to break. Sliding the layers over each other breaks only the weak forces between layers, so graphite is soft. Melting breaks the whole structure apart, including the strong covalent bonds within layers, so the melting point is very high. Softness and melting point measure different things.
That is it. That is the whole idea of this lecture. For any property, ask which bonds or forces must break, and how strong they are. The answer explains the property.
Worked examples
Explain why graphite conducts electricity but diamond does not.
Start with what conduction needs: charged particles that are free to move. Then count electrons in each structure.
In diamond, each carbon atom forms four covalent bonds and uses all four outer electrons in them. Every electron is fixed in a bond, so there are no free charged particles. Diamond does not conduct.
In graphite, each carbon atom forms three covalent bonds. The fourth outer electron is delocalised and can move along the layer. When a potential difference is applied, these electrons flow. Graphite conducts, mainly along the layers.
Check the answer against the rule from the metals lectures: conduction always needs mobile charge. Metals have delocalised electrons. Molten ionic compounds have mobile ions. Graphite has delocalised electrons. Diamond has neither.
How many moles of C–C bonds are there in 12 g of diamond? Estimate the energy needed to break all of them, using the average C–C bond energy of 346 kJ/mol.
12 g of carbon is 1 mol of atoms, because of carbon is 12. Each atom has four bonds. But each bond joins two atoms, so counting "four per atom" counts every bond twice. The number of bonds is therefore per atom.
So 1 mol of diamond contains 2 mol of C–C bonds. The energy to break them all is about kJ.
Compare this with a molecular substance. Boiling separates molecules without breaking any covalent bonds, so it needs far less energy. That difference is the whole reason giant covalent structures have such high melting points. The figure 692 kJ is an estimate, because 346 kJ/mol is an average over many compounds.
Diamond has a density of about 3.52 g/cm³. Graphite has a density of about 2.1 to 2.2 g/cm³. Both are pure carbon. Explain the difference.
Density is mass divided by volume. The atoms are identical, so the difference must come from how much space each atom takes up.
In diamond, the atoms are packed in a network with bonds in all directions, each neighbour 154 pm away. In graphite, the atoms in a layer are close together, but the layers are 335 pm apart. That large, empty gap between layers wastes space. So the same number of atoms fills a larger volume.
Check with numbers: 12 g of diamond fills cm³. 12 g of graphite, at 2.2 g/cm³, fills about cm³.
Now sort the properties. For each one, decide whether it belongs to diamond, to graphite, or to both, and say which feature of the structure explains it.
Put each card in a box. For each one, think about the bonds: four per atom, or three per atom with one delocalised electron, and the layers. Every card explains itself when you place it.
Ten statements sorted by structure.
- Diamond only: each carbon forms four covalent bonds; bond angles of 109.5°; the hardest natural material; does not conduct electricity; transparent when pure.
- Graphite only: each carbon forms three covalent bonds at 120°; has delocalised electrons; layers slide, so it is used as a lubricant.
- Both: a giant covalent structure of carbon; insoluble in water; does not melt until well above 3000 °C.
Three classic mistakes
The weak forces between layers explain softness, not melting. Melting graphite means breaking the strong covalent bonds within the layers. Always name the bonds that the change must break.
Graphite contains no ions, and it is not a metal. It conducts because each carbon has one delocalised electron, left over after three covalent bonds. Write the reason in those words: "one delocalised electron per carbon atom, free to move along the layers".
Students sometimes write that "weak covalent bonds" hold the layers together. Covalent bonds join the atoms within a layer. The forces between layers are intermolecular forces, London forces. Mixing the two words makes the explanation of softness and melting fall apart.
Work each one on paper first. Open a hint only when you are stuck.
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An exam answer says: "Diamond is hard because its covalent bonds are strong." Explain why this answer is incomplete, using graphite as your evidence.
Hint 1
Graphite also contains strong covalent bonds. Is graphite hard?
Hint 2
What is different about the directions of the bonds in the two structures?
Full answer
Graphite has strong covalent bonds too, even slightly shorter ones, yet it is soft. So strong bonds alone cannot explain hardness. The full answer is that diamond's strong bonds extend in all three directions in a continuous network. There is no weak direction, so any scratch must break covalent bonds. In graphite, the strong bonds lie only in the layers, and the weak forces between layers let them slide.
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Graphite conducts electricity about 1000 times better along its layers than across them. Explain why, using the structure.
Hint 1
Where are the delocalised electrons free to move?
Hint 2
Compare the 142 pm distance within a layer with the 335 pm gap between layers.
Full answer
The delocalised electrons are spread across each layer, so they move easily along it. To go across, an electron must cross the 335 pm gap to the next layer, where there are no bonds to carry it. So conduction across the layers is much poorer. A material whose properties differ by direction like this is said to be "anisotropic". Diamond, with bonds in all directions, behaves the same in every direction.
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Four substances are tested. A: melts at 80 °C and does not conduct, solid or molten. B: melts at 900 °C, does not conduct when solid, conducts when molten. C: melts above 3000 °C and does not conduct, solid or molten. D: melts at 1000 °C and conducts when solid. (These are invented data.) Classify each as simple molecular, ionic, metallic or giant covalent, with a reason.
Hint 1
Start with conduction. Which structures have mobile charge in the solid? Which only when molten?
Hint 2
Then use the melting point to separate the two non-conductors.
Full answer
A is simple molecular: a low melting point means only weak intermolecular forces break, and there are no mobile charges. B is ionic: the ions are fixed in the solid but free to move when molten. C is giant covalent: covalent bonds must break to melt it, and it has no delocalised electrons or ions, like diamond. D is metallic: it conducts as a solid because of delocalised electrons. Graphite would be the one giant covalent structure that breaks this pattern, because it conducts as a solid.
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Diamonds can be made from graphite by squeezing it at very high pressure and temperature. A company first did this in a way others could repeat in 1954. (a) Use the densities in Worked Example 3 to find the volume of 12 g of each allotrope. (b) Explain why very high pressure helps graphite turn into diamond.
Hint 1
Volume = mass ÷ density. Use 3.52 g/cm³ for diamond and 2.2 g/cm³ for graphite.
Hint 2
Pressure pushes a substance towards whichever arrangement takes up less space.
Full answer
(a) Diamond: cm³. Graphite: cm³. Diamond takes up only about of the volume. (b) Squeezing a substance favours the arrangement that fits the same atoms into less space. Diamond is that arrangement: it has no wide gaps between layers. At normal pressure graphite is the more stable form, and diamond turns into graphite only extremely slowly, which is why diamonds last. At very high pressure, the balance shifts towards diamond. The high temperature lets the atoms rearrange quickly enough.
For a long time, no one knew what diamond was made of. In 1772 the French chemist Antoine Lavoisier heated a diamond in a sealed jar of oxygen, using sunlight focused by a large lens. The diamond burned, and the gas it produced was carbon dioxide. This showed that diamond is made of carbon, and the result was confirmed by other chemists in 1797. A precious gem and the black graphite of a pencil were the same element.
That discovery holds the lesson of this whole unit. Properties do not come from atoms alone. They come from how the atoms are arranged and bonded. Carbon has other arrangements too: single sheets of graphite, rolled tubes and hollow cages. The next lecture looks at graphene, a single layer of graphite, and at other giant covalent structures.
- A giant covalent structure is a continuous network of covalent bonds with no separate molecules, and diamond and graphite are two such structures, allotropes of carbon.
- In diamond each carbon forms four covalent bonds in a tetrahedral arrangement, so the network has no weak direction and diamond is extremely hard.
- Diamond does not conduct electricity, because all four outer electrons of each carbon are held in bonds.
- In graphite each carbon forms three bonds in flat layers and has one delocalised electron, so graphite conducts electricity along the layers.
- Weak London forces between graphite's layers let the layers slide, which makes graphite soft and slippery.
- Both allotropes have very high melting points, because melting either one requires strong covalent bonds to break.