Giant covalent structures: diamond, graphite and silica

Diamond and graphite are both made only of carbon, yet one is very hard and the other is soft and slippery. Their different structures explain their different properties. In this lesson, connect each property to the particular bonds, forces or electrons responsible for it.

A network of covalent bonds

A giant covalent structure contains many atoms joined by strong covalent bonds in a continuous network. It does not consist of separate small molecules.

Diamond and graphite are forms of the element carbon. Silicon(IV) oxide, also called silicon dioxide or silica, is a compound with the formula SiO₂.

The syllabus groups giant covalent structures and poly(ethene) under “macromolecules”, but their structures differ. Poly(ethene) consists of long-chain molecules. We compare these with giant networks in Comparing covalent structures.

Diamond: a rigid three-dimensional network

Each carbon atom in diamond is covalently bonded to four other carbon atoms in a tetrahedral arrangement. The bonds extend throughout the solid in three dimensions.

• Very hard: moving atoms out of this rigid network requires breaking strong covalent bonds. This makes diamond useful in cutting tools.
• Does not conduct electricity: all four valence electrons per carbon are involved in covalent bonding. There are no mobile charged particles to carry current.
• Strongly bonded network: a large energy input is needed to break many of its covalent bonds. Do not treat diamond as a collection of molecules held together by weak intermolecular forces.

Graphite: strong layers with mobile electrons

Each carbon atom in graphite is covalently bonded to three other carbon atoms, making layers of linked hexagons. One electron per carbon is delocalised: it can move through the layer rather than belonging to one bond.

• Soft and slippery: relatively weak forces act between the layers, so layers can slide over one another without breaking the strong covalent bonds within them. This explains its use in pencil leads and as a lubricant.
• Conducts electricity: delocalised electrons can move along the layers and carry charge. This makes graphite useful as an electrode when it is suitable for the reaction conditions.
• Resists high temperatures: destroying the layers requires breaking many strong covalent bonds. Sliding a layer is a different change from breaking it apart.

Silicon dioxide: a network with two kinds of atom

In silicon dioxide, each silicon atom is covalently bonded to four oxygen atoms, and each oxygen bridges two silicon atoms. The formula SiO₂ gives the overall ratio of atoms, rather than describing a separate molecule in the solid.

• Hard: strong Si–O bonds hold a rigid three-dimensional network together.
• High melting point: melting requires disrupting the network and breaking many strong covalent bonds, so much energy is needed.
• Does not conduct electricity: the structure has no mobile charged particles.

Quartz, a form of silicon dioxide, is common in many sands. Some sands have other compositions, such as calcium carbonate from shells. Silica is a raw material used to make glass and ceramics.

Compare the three giant structures

All three have many strong covalent bonds. That shared feature does not make all their properties identical: graphite’s layers and delocalised electrons account for its important differences.

The comparison with methane, iodine and poly(ethene) now has its own lesson: Comparing covalent structures.

Check that each explanation uses the right feature

• Graphite is soft because layers slide, not because its covalent bonds are weak.
• Being made of carbon does not guarantee conductivity. Compare the availability of mobile electrons in graphite and diamond.
• For silicon dioxide’s melting point, explain the energy needed to break many strong covalent bonds. “It has strong bonds” alone does not connect the structure to the temperature.
• Silicon and silicon dioxide are different substances. Use the correct name and formula.

Worked examples

Modelled example 1

Why graphite can act as a lubricant

Core

Problem

Explain how the structure of graphite allows it to act as a solid lubricant.

Study the worked solution

1. Explain graphite's slipperiness

   Method

   Identify its layered giant structure.

   Reason

   Weak forces between layers allow the layers to slide over one another.

   Working

   Layered structure → sliding layers → lubricating action.

2. Connect sliding to use

   Method

   Allow surfaces to move past one another.

   Reason

   Layers can slide while strong covalent bonds hold each layer together.

   Working

   Sliding layers make graphite useful for lubrication under suitable operating conditions.

Guided practice 2

Why silicon dioxide has a high melting point

About 5 min

Problem

Silicon dioxide has a high melting point. Identify its structure and explain why melting requires a large energy input.

Link structure, bond and energy

Silicon dioxide structure

Simple molecularGiant covalent networkGiant ionic lattice

What melting must overcome

Only weak intermolecular forcesMany strong covalent bondsMobile electrons

Check my attempt

Hints

Hint 1: decide whether molecules exist

Solid silicon dioxide does not consist of separate small molecules.

Hint 2: identify the repeated connection

Strong covalent bonds extend throughout the silicon dioxide network.

View solution step by step

1. State the structure

   Method

   Identify silicon dioxide as a giant covalent structure.

   Reason

   Atoms are joined by covalent bonds through a continuous network.

   Working

   SiO₂: a giant covalent network.

2. Link bonding to melting point

   Method

   Break many strong covalent bonds.

   Reason

   A large energy input is required to disrupt enough of the network for melting.

   Working

   Many strong bonds → much energy → very high melting point.

Common misconception 3

Does easy layer sliding mean easy bond breaking?

Find and correct the mistake

Learner response

A student writes: “Graphite layers slide easily, so little energy is needed to break the layers apart.” Explain why the conclusion does not follow.

Match each force to the correct property

Energy needed to break up graphite layers

LowVery high: many strong covalent bonds must breakNone: sliding and bond breaking are the same change

Weak forces explain

Layer slidingEasy covalent-bond breakingElectron delocalisation

Check my attempt

View solution step by step

1. Locate the property error

   Method

   Correct the claim that breaking the layers apart requires little energy.

   Reason

   Breaking up a layer requires many strong covalent bonds between carbon atoms to be broken.

   Working

   Many strong covalent bonds → much energy needed to destroy the layer.

2. Assign the weak forces correctly

   Method

   Use weak forces between layers to explain sliding.

   Reason

   The layers can move relative to one another without breaking the strong bonds within a layer.

   Working

   Weak interlayer forces → softness and slipperiness.

3. Write the corrected distinction

   Working

   Strong covalent bonds within layers explain the high energy needed to destroy them; weak forces between layers allow sliding.

Examiner practice 4

Why graphite conducts and diamond does not

4 marks

Examination question

Explain why graphite conducts electricity but diamond does not. [4 marks]

Compare bonding and charge carriers

Write a linked graphite–diamond comparison

View solution step by step

1. Explain graphite bonding

   1 mark

   Method

   State that each carbon bonds to three other carbon atoms.

   Reason

   Only three of carbon’s four valence electrons are used in these bonds.

   Working

   One electron per carbon remains delocalised.

2. Explain graphite conduction

   1 mark

   Method

   Make the delocalised electrons mobile along the layers.

   Reason

   Moving electrons carry electrical charge.

   Working

   Graphite conducts.

3. Explain diamond bonding

   1 mark

   Method

   State that each carbon bonds to four other carbon atoms.

   Reason

   All four valence electrons are used in covalent bonds.

   Working

   Diamond has no delocalised electrons.

4. Explain diamond's result

   1 mark

   Method

   Identify the absence of mobile charged particles.

   Reason

   No electrons are available to move through the structure and carry charge.

   Working

   Diamond does not conduct electricity.

Self-mark with the mark scheme

Compare your response with each mark point. Select a point only when your response contains that evidence.

Self-mark electron bonding and mobility for each carbon allotrope.

Mark points

1. 1 mark: States that each graphite carbon forms three bonds, leaving one delocalised electron.
2. 1 mark: States that mobile delocalised electrons carry charge in graphite.
3. 1 mark: States that each diamond carbon uses all four valence electrons in bonds.
4. 1 mark: Links the absence of mobile charge carriers to non-conduction in diamond.

Calculate my mark

The property-evidence example now appears in Comparing covalent structures.

Try it yourself

Mind stretcher 1: Choosing an electrode materialExtension

Question: An electrode must conduct electricity and remain solid at a high operating temperature. Tests show that graphite remains solid and does not react with the electrolyte under these conditions. Choose graphite or diamond and explain the choice.

Show Answer

Choose graphite:

• conducts electricity due to delocalised electrons
• many strong covalent bonds hold its layers together, consistent with the supplied observation that it remains solid at the operating temperature

Diamond does not conduct electricity.

Mind stretcher 2: Same element, different propertiesExtension

Question: Diamond and graphite are both C. Why do they have different properties?

Show Answer

They have different structures (different bonding and arrangement of atoms):

• diamond: 3D tetrahedral network, no delocalised electrons
• graphite: layered structure with delocalised electrons

Different structure → different properties.

Practise and check