Chemistry · 2 Atoms, elements and compounds

2.6 Giant covalent structures

Before you start

A covalent bond is a shared pair of electrons between two non-metal atoms, and it is strong.
Simple molecular substances such as water or carbon dioxide have low melting points because only the weak intermolecular forces between molecules are overcome; the covalent bonds inside each molecule stay intact.
Electronic configurations: carbon is 2,4, silicon is 2,8,4 and oxygen is 2,6. You can read the group number from the Periodic Table printed in the paper.
To conduct electricity a substance must contain charged particles that are free to move — electrons or ions.

Syllabus objectives covered (4)

Describe the giant covalent structures of graphite and diamond
Relate the structures and bonding of graphite and diamond to their uses, limited to: (a) graphite as a lubricant and as an electrode (b) diamond in cutting tools
Describe the giant covalent structure of silicon(IV) oxide, SiO2
Describe the similarity in properties between diamond and silicon(IV) oxide, related to their structures

Part 1 of 6

Giant means there are no molecules

In 2.5 you met covalent substances made of small, separate molecules. Each molecule is held together by strong covalent bonds, but the molecules are attracted to each other only weakly. That is why water and carbon dioxide melt or boil easily.
A giant covalent structure has no separate molecules at all. Every atom is joined by covalent bonds to its neighbours, and those neighbours to theirs, on and on through the whole crystal. The entire lump is one continuous network of covalent bonds, called a lattice.
Notice what that does to melting. In a simple molecular substance, melting only pulls molecules apart. In a giant covalent substance there are no molecules to pull apart, so melting means breaking strong covalent bonds — a very large number of them. That needs a lot of energy, so the melting point is very high.
The trap here is the phrase 'intermolecular forces'. Intermolecular means between molecules. Diamond and silicon(IV) oxide have no molecules, so if you write that melting them overcomes intermolecular forces, you have contradicted yourself and lost the mark.
You need three giant covalent structures for this syllabus: diamond and graphite, which are both pure carbon, and silicon(IV) oxide, $SiO_{2}$ . Diamond and graphite are the same element with its atoms arranged differently, and that difference in arrangement is what gives them such different properties.

Part 2 of 6

Diamond: four bonds on every carbon

Start from the electrons. Carbon is 2,4, so it has four outer electrons. Each can be shared with a different neighbour, so each carbon atom can form four covalent bonds.
In diamond every carbon atom is covalently bonded to four other carbon atoms. The four bonds point to the corners of a tetrahedron, so the network spreads out in all three dimensions — up, down and sideways equally.
Count the bonds on any one atom: always four, all strong covalent bonds, running in every direction.
Very hard. To scratch or break diamond you have to break strong covalent bonds, and because the network of strong bonds runs in all three directions there are no weak forces anywhere for layers to slide along, as there are in graphite. Diamond is the hardest natural substance.
Very high melting point. Melting means breaking a huge number of strong covalent bonds, which takes a great deal of energy.
Does not conduct electricity. All four outer electrons of every carbon are being used in covalent bonds, so there are no free electrons. There are no ions either. Nothing charged can move, so no current flows.
Insoluble in water. Water molecules cannot break the covalent bonds of the lattice.
When you are asked to describe the structure of diamond, give three things: each carbon is covalently bonded to four other carbons; the arrangement is tetrahedral; and it is a giant three-dimensional structure. 'Carbon atoms joined together' is not a description — it is true of graphite too.

Part 3 of 6

Graphite: layers, and one spare electron

In graphite each carbon atom is covalently bonded to only three other carbon atoms. The three bonds lie flat, so the atoms form flat sheets — layers — made of hexagonal rings, like a honeycomb.
Strong covalent bonds inside each layer; only weak forces between one layer and the next.
Now do the electron arithmetic. Carbon has four outer electrons and uses three of them in bonds. That leaves one outer electron per carbon atom that is not in a bond.
$4 outer electrons - 3 used in bonds = 1 delocalised electron per atom$
These spare electrons are delocalised: they do not belong to any one atom and are free to move along the layers. Moving charged particles are exactly what a current needs, so graphite conducts electricity — unusual for a non-metal.
Between one layer and the next there are no covalent bonds, only weak forces of attraction. Very little energy is needed to overcome them, so the layers can slide over each other. That makes graphite soft and slippery to the touch.
Here is the awkward case. Graphite is soft, yet its melting point is very high. Both are true because they depend on different things: sliding only overcomes the weak forces between layers, but melting means breaking the strong covalent bonds within the layers.
Name the weak forces carefully. They are weak forces between the layers. Never write 'weak covalent bonds between the layers' — the covalent bonds in graphite are strong, and the ones you are describing are not covalent bonds at all.

Part 4 of 6

From structure to use

The syllabus limits you to three uses, and every one is answered with the same three-link chain: use → the property it needs → the structure that causes the property. A mark scheme usually has one mark for the property and one for the structural reason.
Graphite as a lubricant. A lubricant has to be slippery so that surfaces slide past each other. Graphite is slippery because its layers are held together only by weak forces, so the layers slide over each other easily.
Graphite as an electrode. An electrode has to conduct electricity. Graphite conducts because each carbon atom has one delocalised electron, and these electrons are free to move through the structure. Graphite is also unreactive and has a high melting point, which is why it is used as an inert electrode in electrolysis.
Diamond in cutting tools. Drill bits and saw blades for cutting rock or glass are tipped with diamond because it is very hard. It is hard because every carbon atom is held by four strong covalent bonds in a giant three-dimensional lattice, so the tool can cut a material without being worn away itself.
The word that scores for the cutting tool is hard. 'Strong' is not the same property, and 'high melting point' is true but is not why diamond cuts. Give the property the use actually needs.
Reverse the logic to explain why the two forms of carbon are not swapped. Diamond could never be an electrode, because it has no free electrons. Graphite could never be a cutting tool, because its layers would slide apart under the load.

Part 5 of 6

Silicon(IV) oxide: diamond's pattern with oxygen in between

Silicon is in Group IV with carbon, so its atoms also have four outer electrons (2,8,4) and each silicon atom forms four covalent bonds.
Oxygen has six outer electrons (2,6) and needs two more, so each oxygen atom forms two covalent bonds.
$8 - 6 = 2 bonds per oxygen atom$
In silicon(IV) oxide each silicon atom is covalently bonded to four oxygen atoms, arranged tetrahedrally, and each oxygen atom is bonded to two silicon atoms. Every oxygen is a bridge between two silicons, and the network continues in three dimensions: a giant covalent structure.
Four bonds on every silicon, two on every oxygen. It is diamond's tetrahedral network with an oxygen bridge in every link.
The formula $SiO_{2}$ comes from those bond counts. Each silicon has four bonds to oxygen, but each oxygen is shared between two silicons, so each silicon effectively owns half of four oxygens — two of them.
$4 \times \frac{1}{2} = 2 \Rightarrow Si : O = 1 : 2$
So the formula of a giant structure is a ratio of atoms, not a molecule. There is no such thing as 'one $SiO_{2}$ molecule'; a grain of sand is one enormous network with silicon and oxygen atoms in the ratio 1 : 2. Sand and quartz are mostly silicon(IV) oxide.
Do not let carbon dioxide mislead you. $CO_{2}$ and $SiO_{2}$ have formulae that look alike, but carbon dioxide is made of small separate molecules and is a gas at room temperature. Silicon(IV) oxide is a giant structure and a hard solid with a very high melting point. Same pattern of formula, completely different structure.

Part 6 of 6

Why diamond and silicon(IV) oxide behave alike

Diamond and silicon(IV) oxide contain different elements, yet their properties are very similar. The reason is that their structures are the same kind: in both, every atom is held in a giant three-dimensional lattice by strong covalent bonds, and no electrons are free to move.
Both have very high melting and boiling points, because many strong covalent bonds must be broken, which needs a lot of energy.
Both are hard, because the strong covalent bonds run in all three directions and must be broken to scratch the solid.
Neither conducts electricity, because all the outer electrons are used in covalent bonds — there are no delocalised electrons and no ions.
Both are insoluble in water.
The structure of an answer is the thing that scores: name the shared property, then attach the shared structural reason. 'They are both hard and have high melting points' earns the property marks; adding 'because both are giant covalent structures with many strong covalent bonds that need a lot of energy to break' earns the explanation.
Graphite is the one that breaks the pattern. It is giant covalent too, but its layered structure and delocalised electrons make it soft and conducting. If a question asks which substance is like silicon(IV) oxide, the answer is diamond, not 'carbon'.

Worked examples

Example 1. Diamond and graphite are both made only of carbon atoms. Explain why graphite conducts electricity but diamond does not, and state which of the two is used as an electrode.

Show the working

Start with what conduction needs: charged particles that are free to move. Neither substance has ions, so the question is about electrons.
Diamond: each carbon is bonded to four other carbons. Carbon has four outer electrons, so all four are used in covalent bonds. No electrons are free, so there is nothing to carry a current.
Graphite: each carbon is bonded to only three others, so one outer electron per atom is not used in bonding.
$4 - 3 = 1$
That electron is delocalised and free to move along the layers, so graphite conducts. Say 'electrons' explicitly — 'charges move' or 'ions move' does not score.
Because it conducts, graphite is the one used as an electrode.

Answer: In diamond each carbon atom uses all four outer electrons in covalent bonds to four other carbons, so there are no free electrons. In graphite each carbon is bonded to only three others, leaving one delocalised electron per atom that is free to move through the structure, so graphite conducts. Graphite is used as the electrode.

Example 2. Oil cannot be used to lubricate a machine part that becomes very hot, because the oil would decompose. Suggest why graphite powder is a suitable lubricant for this part.

Show the working

Two properties are needed here: it must be slippery, and it must survive high temperature. Deal with each in turn.
Slippery: graphite has layers held together only by weak forces, so the layers slide over each other easily.
Survives the heat: within each layer every carbon is held by strong covalent bonds, so a very large amount of energy is needed to melt graphite. It stays solid at temperatures that would destroy the oil.
Notice that the second point uses the bonds within the layers and the first uses the forces between them. Mixing them up is the commonest way to lose these marks.

Answer: Graphite's layers are held together by weak forces, so they slide over each other, making it slippery. The carbon atoms within each layer are joined by strong covalent bonds, so graphite has a very high melting point and remains a solid lubricant at high temperature.

Example 3. In silicon(IV) oxide every silicon atom is covalently bonded to four oxygen atoms and every oxygen atom is covalently bonded to two silicon atoms. Use this to show that the formula is $SiO_{2}$ , and explain why the formula does not describe a molecule.

Show the working

Take one silicon atom. It has four bonds, each to a different oxygen atom.
Each of those oxygens is also bonded to a second silicon, so it is shared equally between two silicon atoms. One silicon can claim only half of each oxygen.
Multiply: four oxygens, each half-owned, gives two oxygen atoms per silicon atom.
$4 \times \frac{1}{2} = 2$
So the ratio of silicon to oxygen atoms is 1 : 2, written $SiO_{2}$ .
The network has no edges inside the crystal — every atom is bonded on to others in all directions — so there is no small group of three atoms that stands alone. The formula is the ratio of atoms in a giant lattice.

Answer: Each silicon is bonded to 4 oxygens, each shared with one other silicon, so each silicon has $4 \times \frac{1}{2} = 2$ oxygens: ratio Si : O = 1 : 2, giving $SiO_{2}$ . It is a giant covalent structure with no separate molecules, so the formula gives only the ratio of atoms.

Example 4. Silicon carbide has a giant structure in which every silicon atom is covalently bonded to four carbon atoms and every carbon atom is covalently bonded to four silicon atoms, arranged tetrahedrally in three dimensions. Predict two physical properties of silicon carbide, and suggest a use for it.

Show the working

You have not been taught silicon carbide, and you do not need to have been. Read the structure: tetrahedral, four bonds on every atom, giant, three-dimensional. That is the diamond and silicon(IV) oxide pattern.
Same structure type, so predict the same properties: very high melting point, very hard, does not conduct electricity, insoluble in water. Any two will do.
Back each one with the structural reason: many strong covalent bonds in a three-dimensional lattice, and all outer electrons used in bonding so none are free.
For a use, ask which property is most unusual. Hardness is, and the use you know that depends on hardness is cutting. So suggest cutting or grinding tools, exactly as for diamond.
As a check on the structure, the bond counts give the formula: four bonds per silicon, four per carbon, so the ratio is 1 : 1.
$\frac{4}{4} = 1 \Rightarrow Si : C = 1 : 1$

Answer: Any two of: very high melting point; very hard; does not conduct electricity; insoluble in water — because it is a giant covalent structure with many strong covalent bonds and no free electrons. A suitable use is in cutting or grinding tools, because it is very hard.

Where people go wrong

Mistake: Writing that diamond or silicon(IV) oxide has a high melting point because of strong intermolecular forces.

Instead: There are no molecules in a giant covalent structure. Melting breaks many strong covalent bonds.

Intermolecular forces belong to simple molecular substances. Using the term here contradicts the correct idea, so the mark is lost even if you also mention covalent bonds.

Mistake: Saying graphite is soft because the covalent bonds between the layers are weak.

Instead: There are no covalent bonds between the layers — only weak forces of attraction. The covalent bonds within the layers are strong.

Calling them weak covalent bonds also makes it impossible to explain why graphite has a high melting point.

Mistake: Saying graphite conducts because it has free ions, or because 'charges' move.

Instead: Graphite conducts because it has delocalised electrons that are free to move along the layers.

Graphite contains no ions at all. The word 'electrons' is the marking point.

Mistake: Saying diamond is used in cutting tools because it is strong or has a high melting point.

Instead: Diamond is used because it is very hard, and it is hard because of its giant three-dimensional network of strong covalent bonds.

Each use is paired with one property. Hardness is the property a cutting tool needs, so it is the one that scores.

Mistake: Treating $SiO_{2}$ as a small molecule like $CO_{2}$ .

Instead: Silicon(IV) oxide is a giant covalent structure; $SiO_{2}$ is only the 1 : 2 ratio of atoms in the lattice.

A similar-looking formula says nothing about structure. Carbon dioxide is a gas; silicon(IV) oxide is a hard solid with a very high melting point.

Check yourself

State how many other carbon atoms each carbon atom is covalently bonded to in (i) diamond and (ii) graphite.
Answer
(i) four, (ii) three.
Graphite is soft but has a very high melting point. Explain how both can be true.
Answer
Softness depends on the weak forces between the layers, which are easily overcome so the layers slide. Melting requires breaking the strong covalent bonds within the layers, which needs a lot of energy.
State the property of diamond that makes it suitable for cutting tools, and give the structural reason for it.
Answer
It is very hard, because each carbon atom is held by four strong covalent bonds in a giant three-dimensional lattice.
Suggest why graphite, not diamond, is used to make electrodes.
Answer
Graphite conducts electricity because it has delocalised electrons that are free to move; diamond has no free electrons, because all four outer electrons of each carbon are used in bonding, so it does not conduct.
Describe the structure of silicon(IV) oxide.
Answer
A giant covalent structure in three dimensions: each silicon atom is covalently bonded to four oxygen atoms (tetrahedrally) and each oxygen atom is covalently bonded to two silicon atoms.
Carbon dioxide is a gas at room temperature, but silicon(IV) oxide is a solid with a very high melting point. Explain the difference.
Answer
Carbon dioxide is made of simple molecules with weak intermolecular forces between them, which need little energy to overcome. Silicon(IV) oxide is a giant covalent structure, so melting it breaks many strong covalent bonds, which needs a lot of energy.
Give three properties shared by diamond and silicon(IV) oxide.
Answer
Any three of: very high melting and boiling points; hard; do not conduct electricity; insoluble in water.

Exam questions

Try each part first. Open the mark scheme after you have written your answer.

Question 1. Diamond and graphite are two forms of the element carbon. Both have giant covalent structures.

(a) Describe the structure of diamond. 2 marks No calculator

Mark scheme

B1 each carbon atom is (covalently) bonded to four other carbon atoms
B1 giant / three-dimensional / tetrahedral structure or lattice

(b) Explain why graphite conducts electricity. 2 marks No calculator

Mark scheme

B1 each carbon atom is bonded to three others, so one outer electron per atom is not used in bonding / graphite has delocalised electrons
B1 these electrons are free to move (through the structure / along the layers); any mention of ions moving is CON

(c) Graphite is used as a lubricant. Explain, in terms of its structure, why graphite is a good lubricant. 2 marks No calculator

Mark scheme

B1 (graphite has) layers held together by weak forces (of attraction); weak covalent bonds between layers is CON
B1 the layers can slide over each other, owtte

(d) The tips of drill bits used to cut through rock are coated with diamond. State the property of diamond that makes it suitable for this use, and explain this property in terms of its structure and bonding. 3 marks No calculator

Mark scheme

B1 (very) hard; strong alone is insufficient; ignore high melting point
B1 each carbon atom forms four strong covalent bonds
B1 in a giant three-dimensional lattice, so many bonds must be broken to scratch or break it

Question 2. Sand is mainly silicon(IV) oxide, $SiO_{2}$ . Silicon(IV) oxide has properties similar to those of diamond. Carbon dioxide, $CO_{2}$ , is a gas at room temperature.

(a) Describe the structure of silicon(IV) oxide. 3 marks No calculator

Mark scheme

B1 each silicon atom is (covalently) bonded to four oxygen atoms
B1 each oxygen atom is (covalently) bonded to two silicon atoms
B1 giant covalent / three-dimensional / tetrahedral structure or lattice; molecules of silicon(IV) oxide is CON

(b) State two physical properties that silicon(IV) oxide and diamond have in common. 2 marks No calculator

Mark scheme

B1 any one from: very high melting point / very high boiling point (melting and boiling point together count as one property); hard; does not conduct electricity; insoluble in water
B1 a second property from the same list

(c) Explain, in terms of structure and bonding, why silicon(IV) oxide and diamond both have very high melting points. 3 marks No calculator

Mark scheme

B1 both have giant covalent structures (no separate molecules); intermolecular forces is CON
B1 many strong covalent bonds
B1 must be broken, which needs a lot of energy

(d) The formulae $CO_{2}$ and $SiO_{2}$ look similar. Explain why carbon dioxide is a gas at room temperature but silicon(IV) oxide is a solid. 2 marks No calculator

Mark scheme

B1 carbon dioxide consists of simple molecules with weak intermolecular forces (between them), which need little energy to overcome
B1 silicon(IV) oxide is a giant structure, so (many) strong covalent bonds must be broken to melt it, ORA

How this is examined

Why diamond and silicon(IV) oxide both have very high melting points and are hard: a two- or three-mark Explain, often after you have been told the two substances have similar properties, sometimes alongside a data table of melting points Three separate ideas: both are giant covalent structures (or 'giant lattice', 'macromolecular'); there are many strong covalent bonds; these must be broken, which needs a lot of energy. Any mention of intermolecular forces being overcome contradicts the answer and costs the mark.
Describing a giant covalent structure precisely enough to tell diamond, graphite and silicon(IV) oxide apart: a two- or three-mark Describe, sometimes paired with a diagram to label or a sketch to complete The bond count on each atom plus 'giant': diamond — each carbon covalently bonded to four other carbons, tetrahedral, giant three-dimensional; graphite — each carbon bonded to three others in flat hexagonal layers with weak forces between layers; silicon(IV) oxide — each silicon bonded to four oxygens and each oxygen to two silicons, giant three-dimensional. 'Lots of atoms joined together' scores nothing.
Why graphite conducts electricity and diamond does not: a two-mark Explain, often written as a contrast between the two forms of carbon Graphite: one outer electron per carbon is not used in bonding, so there are delocalised (free, mobile) electrons that can move. Diamond: all four outer electrons are used in bonds, so there are no free electrons. The word 'electrons' must appear; 'ions' is wrong.
Why graphite works as a lubricant: a two-mark Explain, usually 'in terms of structure' Two marks, two ideas: the layers are held together by weak forces; so the layers can slide over each other. 'Weak covalent bonds between layers' contradicts the correct science.
Why diamond is used in cutting tools: a one- or two-mark State or Explain, often a choice between diamond and graphite for a given use The property 'very hard' — not 'strong' — then the reason: each carbon atom has four strong covalent bonds in a giant three-dimensional structure.

Words to know

Giant covalent structure: A structure in which a very large number of atoms are joined by covalent bonds in a continuous network, with no separate molecules. Diamond, graphite and silicon(IV) oxide are examples.
Lattice: A regular, repeating arrangement of particles that extends through the whole crystal.
Tetrahedral: Describes four bonds from one atom pointing to the four corners of a tetrahedron, as in diamond and around each silicon atom in silicon(IV) oxide.
Delocalised electron: An electron that is not held in a bond between two particular atoms and is free to move through the structure, as in graphite.
Lubricant: A substance that reduces friction by letting surfaces slide over each other easily.
Electrode: A conductor through which electricity enters or leaves a liquid or solution during electrolysis.
Silicon(IV) oxide: $SiO_{2}$ , the main compound in sand; a giant covalent structure in which each silicon is bonded to four oxygens and each oxygen to two silicons.

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