0620 · 2.6Giant covalent structures

Giant covalent bonding and diamond

0:00Diamond can form the cutting edge of a tool, while graphite helps surfaces slide past each other. Yet both substances contain only carbon atoms. Their different behaviour comes from how those atoms are connected. By the end of this video, you will understand how giant covalent structures explain hardness, electrical conductivity and these useful applications. We need one starting idea about bonding. A covalent bond is a shared pair of electrons between two atoms.

0:25This bond strongly holds the atoms together. In a simple molecular substance, only a small group of atoms is joined together within each molecule. A giant covalent structure is different because covalent bonds extend through a huge network of atoms. There are no separate small molecules making up the substance. The word giant describes the connected structure, not the size of each atom. We can therefore understand its properties by examining the bonds and the arrangement throughout that network.

0:51In diamond, each carbon atom forms four strong covalent bonds to four other carbon atoms. These neighbours are arranged tetrahedrally around it. This means they point towards the corners of a tetrahedron, rather than lying in one flat square. The same bonding pattern continues throughout a giant three-dimensional structure. The important feature is that strong bonds hold the structure together in every direction. There are no layers that can easily slide over one another.

1:15Moving atoms out of position requires breaking strong covalent bonds. This makes diamond very hard, meaning that it strongly resists scratching. That hardness explains why diamond is used in cutting tools. A diamond cutting edge can scratch or cut into a softer material while resisting wear itself. The explanation begins with four bonds per carbon and a rigid three-dimensional network. Simply saying that diamond has strong bonds misses the importance of how those bonds are arranged.

1:42Diamond also has a very high melting point. Melting diamond requires breaking many strong covalent bonds. A large amount of energy is needed, so a very high temperature is required. The key connection is between the extensive bonded network and the energy needed to disrupt it. Despite containing electrons, diamond does not conduct electricity. Its outer electrons are involved in covalent bonds and cannot move freely through the structure. Electrical conduction requires mobile charged particles.

2:09Having electrons inside a substance is not enough if those electrons cannot carry charge through it.

Graphite’s layers and conduction

2:14Graphite has a different arrangement of the same element. Each carbon atom forms strong covalent bonds to three other carbon atoms. These connections produce flat layers containing hexagonal rings. The layers are stacked, with weak forces between them rather than strong covalent bonds joining neighbouring layers. Separate the bonding within a layer from the forces between layers. Within each layer, the atoms are strongly held together. The attractions between neighbouring layers are much weaker.

2:41Sliding one layer over another does not require breaking the strong covalent bonds within it. Imagine a stack of sheets that can slide across each other while each sheet remains intact. Graphite behaves similarly because its strongly bonded layers can move past neighbouring layers. This makes graphite soft and slippery. The analogy describes the layered arrangement, although real graphite layers consist of bonded carbon atoms rather than paper. A lubricant reduces friction between moving surfaces.

3:07Graphite can act as a lubricant because its layers slide over one another easily. The moving surfaces can therefore pass each other more smoothly. For this use, the crucial feature is the weak forces between layers, not weak covalent bonds within them. Graphite also differs from diamond in its electrical behaviour. Carbon has four outer electrons, but each carbon atom in graphite forms only three covalent bonds. One electron per carbon atom is delocalised.

3:32Delocalised means that it is not confined to a particular atom or covalent bond. These delocalised electrons can move along the layers and carry electrical charge. Graphite therefore conducts electricity, unlike diamond. This makes graphite suitable for use as an electrode. An electrode provides an electrical connection, so the material must allow charge to move through it. Graphite therefore has a separate structural explanation for each of its uses. Weak forces between layers explain lubrication because the layers can slide.

4:01Its use as an electrode depends instead on mobile delocalised electrons, which conduct electricity. Sliding layers do not cause conduction, and moving electrons do not explain why graphite is slippery. Graphite also has a very high melting point because many strong covalent bonds must be broken to disrupt its giant structure. This does not contradict its softness. Sliding intact layers requires much less energy than breaking up those layers. Hardness and melting point describe different changes, so a soft substance can still be very difficult to melt.

Silicon(IV) oxide structure

4:30This next part is for the Extended paper. Silicon(IV) oxide also has a giant covalent structure, but it contains two different elements. Each silicon atom is covalently bonded to four oxygen atoms. Each oxygen atom is covalently bonded to two silicon atoms, linking the network together. The four oxygen neighbours around each silicon have a tetrahedral arrangement. Each oxygen acts as a link between two silicon atoms. These connections continue throughout a giant three-dimensional network.

4:58As with diamond, the diagram shows only a small part of a structure that extends much further. The formula on screen gives the ratio of silicon atoms to oxygen atoms. There are two oxygen atoms for every silicon atom overall. Although each silicon touches four oxygens, each oxygen is shared between two silicon centres. The formula describes the simplest atom ratio, not a separate molecule containing only three atoms. Silicon(IV) oxide is hard because strong covalent bonds hold its three-dimensional network together.

5:28It has a very high melting point because many of those bonds must be broken. Breaking them requires a large amount of energy. These properties are similar to diamond because both have extensive, strongly bonded networks. Silicon(IV) oxide also does not conduct electricity. It has no free electrons or other mobile charged particles to carry charge through its structure. Diamond has the same electrical limitation. Their strong covalent bonds explain hardness and high melting points, while the absence of mobile charge carriers explains non-conduction.

Identifying structures in exams

5:58Let us work through an identification example rather than just recall a name. An unknown solid contains only carbon, conducts electricity and has layers that slide easily. First, use its composition to identify it as a form of carbon, rather than silicon(IV) oxide. Then use its arrangement and conductivity to distinguish graphite from diamond. The sliding layers match graphite because weak forces act between neighbouring layers. Its conductivity provides a second piece of evidence because graphite has mobile delocalised electrons.

6:27Diamond would not fit because it has a rigid three-dimensional network and no free electrons. The answer is graphite, supported by two separate structure and property connections. A common exam mistake is to describe diamond or graphite as collections of small molecules. Their atoms form giant covalent structures instead. Another mistake is to say that graphite contains weak covalent bonds. Its covalent bonds within the layers are strong; only the forces between layers are weak.

6:53Also avoid saying that every giant covalent substance cannot conduct electricity. Graphite is the important exception because it contains delocalised electrons. Nor should you explain non-conduction simply by mentioning strong bonds. You must identify whether mobile charged particles are available to carry charge. An exam question may ask you to explain graphite’s uses in terms of structure and bonding. Start by separating the lubricant explanation from the electrode explanation. For each use, identify the relevant structural feature and connect it to a property.

7:25Then show how that property makes the material suitable. Graphite has carbon atoms arranged in layers, with each carbon covalently bonded to three others. There are weak forces between the layers, so they slide over one another easily. This makes graphite suitable as a lubricant. Graphite has delocalised electrons that move through its layers and carry charge. It therefore conducts electricity and can be used as an electrode. An Extended question may ask you to describe similarities between diamond and silicon(IV) oxide.

7:54Both are hard because their atoms are held in rigid giant networks by strong covalent bonds. Both have very high melting points because many strong covalent bonds require much energy to break. Neither conducts electricity because neither has mobile charged particles. The central idea is that arrangement matters as much as the type of atom. Diamond has four bonds per carbon in a tetrahedral three-dimensional network, making it suitable for cutting tools.

8:17Graphite has three bonds per carbon in flat hexagonal layers. Weak forces between those layers explain lubrication, and delocalised electrons allow conduction. In silicon(IV) oxide, each silicon bonds to four oxygen atoms, and each oxygen links two silicon atoms. Its giant network explains its similarities to diamond in hardness, melting point and electrical behaviour. Always build an explanation from structure to property to use. That chain of reasoning turns a list of facts into an understanding of the material.