Giant Structures

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Notes

Giant Covalent Structures: Diamond & Graphite

Both diamond and graphite are **allotropes of carbon** with **giant covalent structures**.
Giant covalent structures contain billions of non-metal atoms joined by **covalent bonds** forming a giant lattice.
In **diamond**, each carbon atom bonds to **four others** in a tetrahedral arrangement.
In **graphite**, each carbon atom bonds to **three others**, forming layers of hexagons with one **delocalised electron** per atom.
Covalent bonds within graphite layers are **very strong**; layers are held by **weak intermolecular forces**.
Diamond has **no free electrons** – all outer electrons are in covalent bonds, so it does **not conduct electricity**.
Graphite conducts electricity because **delocalised electrons** can move between layers and carry charge.
Both have **very high melting points** due to strong covalent bonds requiring much energy to break.

Properties & Uses of Diamond

Diamond is **extremely hard and dense** because each carbon is bonded to four others in a rigid tetrahedral structure.
It is used in **cutting tools** (e.g., drills) due to its hardness.
Diamond is **brittle** – hard but can be smashed with a hammer.
It does **not conduct electricity** (no delocalised electrons).

Properties & Uses of Graphite

Graphite is **slippery** – layers can slide over each other due to weak intermolecular forces.
Used in **pencils** and as an **industrial lubricant** because of its slipperiness.
Graphite conducts electricity – used in electrodes.
Pencil 'lead' is actually graphite, not the metal lead.

Metallic Bonding

Metals consist of a **giant lattice** of positive metal ions surrounded by a **'sea of delocalised electrons'**.
Metallic bond is the **strong electrostatic attraction** between positive ions and delocalised electrons.
Delocalised electrons are **free to move** throughout the structure.
Metal alloys are mixtures of metals with the same bonding type.

Properties of Metals

Most metals have **high melting and boiling points** due to strong metallic bonds.
Metals are **good conductors of heat and electricity** because delocalised electrons carry charge and energy.
Metals are **malleable** (can be hammered into shape) and **ductile** (can be drawn into wires) because layers of ions can slide over each other.
Conduction in metals is by **free electrons**; in molten/ aqueous ionic compounds it is by **free ions**.

Diamond structure: each carbon atom is covalently bonded to four others in a tetrahedral arrangement, forming a giant covalent lattice.

Graphite structure: carbon atoms form hexagonal layers with delocalised electrons between layers; weak intermolecular forces allow layers to slide.

Metallic bonding: positive metal ions (grey) are surrounded by a 'sea' of delocalised electrons (red), which are free to move and carry charge.

Practice questions

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1.What type of structure does graphite have?
Easy
- AGiant ionic structure
- BGiant covalent structure
- CSimple molecular structure
- DMetallic structure
2.Why does graphite conduct electricity?
Medium
- AIt has free ions
- BIt has delocalised electrons
- CIt has a giant ionic lattice
- DIt has strong covalent bonds
3.State one use of diamond that relies on its hardness.
Easy
4.Diamond conducts electricity.
Easy
True or false?
5.Complete the sentence.
Medium
In graphite, each carbon atom is bonded $to ____$ other carbon atoms.
6.How many carbon atoms is each carbon atom bonded to in diamond?
Medium
- A3
- B4
- C6
- D5
7.Arrange the following substances in order of increasing electrical conductivity: diamond, graphite, copper.
Hard
- diamond
- graphite
- copper
8.Match each property to the correct carbon allotrope.
Medium
- Conducts electricity
- Very hard
- Slippery layers
- Diamond
- Graphite
- Both

Slides

Giant Structures

Giant structures contain billions of atoms joined by strong bonds in a lattice.
There are two types: giant covalent (e.g. diamond, graphite) and giant metallic.
Properties depend on bonding and structure.

Diamond – Structure

Each carbon atom bonds to four others in a tetrahedral arrangement.
All bonds are strong covalent bonds; no free electrons.
This makes diamond extremely hard and gives it a very high melting point.

How many carbon atoms is each carbon atom bonded to in diamond?

Graphite – Structure

Each carbon atom bonds to three others, forming hexagonal layers.
One electron per carbon is delocalised and free to move between layers.
Layers are held together by weak intermolecular forces, allowing them to slide.

Which of these are true about graphite?

Graphite conducts electricity.
Graphite has a low melting point.
The layers in graphite can slide over each other.
Each carbon atom in graphite is bonded to four others.

Properties of Diamond

Very high melting point due to strong covalent bonds.
Does not conduct electricity – no free electrons.
Extremely hard and dense – used in cutting tools.

Why does diamond not conduct electricity?

It has delocalised electrons.
All outer electrons are held in covalent bonds.
It has weak intermolecular forces.
It is a giant metallic structure.

Properties of Graphite

High melting point due to strong covalent bonds within layers.
Conducts electricity because delocalised electrons can move.
Slippery because layers slide over each other – used as lubricant and in pencils.

Graphite conducts electricity because it $has ____$ electrons that are free to move.

Metallic Bonding

Metals consist of a giant lattice of positive metal ions.
Outer electrons are delocalised and form a 'sea of electrons'.
The strong electrostatic attraction between ions and electrons is the metallic bond.

Which of these are properties of metals?

They have high melting points.
They are good conductors of electricity.
They are brittle.
They are malleable.

Why Metals Conduct

Delocalised electrons are free to move through the metal lattice.
They can carry electrical charge and thermal energy.
This makes metals good conductors of heat and electricity.

What is the force of attraction in metallic bonding?

Between positive metal ions and delocalised electrons
Between metal atoms and covalent bonds
Between layers of atoms
Between positive and negative ions

Malleability and Ductility

Metal atoms are arranged in layers that can slide over each other.
When force is applied, layers move without breaking bonds.
This allows metals to be hammered into shapes (malleable) or drawn into wires (ductile).

Metals are malleable because their atoms are arranged $in ____$ that can slide over each other.

Comparing Diamond and Graphite

Both are allotropes of carbon with giant covalent structures.
Diamond: each C bonded to 4 others; hard, non-conductor.
Graphite: each C bonded to 3 others; soft, conductor.

Which property is shared by diamond and graphite?

They both conduct electricity.
They both have a high melting point.
They are both soft.
They both have delocalised electrons.

Flashcards

What is a giant covalent structure?

A giant covalent structure contains billions of non-metal atoms, each joined to adjacent atoms by covalent bonds forming a giant lattice structure.

What are allotropes?

Allotropes are different structural forms of the same element, such as diamond and graphite for carbon.

Describe the structure of graphite.

Each carbon atom in graphite is bonded to three others forming layers of hexagons, with one free electron per carbon atom that becomes delocalised. The layers are held together by weak intermolecular forces.

Why does graphite conduct electricity?

Graphite conducts electricity because each carbon atom has one delocalised electron that is free to move between the layers and carry charge.

Why does graphite have a high melting point?

Graphite has a high melting point because it has a giant covalent structure with strong covalent bonds between carbon atoms that require a lot of energy to break.

Why is graphite slippery?

Graphite is slippery because its layers are held together by weak intermolecular forces, allowing them to slide over each other.

What are two uses of graphite based on its properties?

Graphite is used in pencils and as an industrial lubricant because its layers can slide over each other.

Describe the structure of diamond.

In diamond, each carbon atom bonds with four other carbons, forming a tetrahedron. All covalent bonds are identical and very strong, with no intermolecular forces.

Why does diamond not conduct electricity?

Diamond does not conduct electricity because all outer shell electrons are held in the four covalent bonds around each carbon atom, so there are no free-moving charged particles.

Why does diamond have a very high melting point?

Diamond has a very high melting point because it has a giant covalent structure with strong covalent bonds that require a lot of energy to break.

Why is diamond extremely hard?

Diamond is extremely hard because each carbon atom is bonded to four other carbon atoms by strong covalent bonds, forming a rigid tetrahedral structure.

What is a common use of diamond based on its hardness?

Diamond is used in cutting tools like drills because of its extreme hardness.

What is the difference between hard and strong?

Hard means resistant to scratching or indentation, while strong means able to withstand force without breaking. Diamond is hard but brittle.

What is metallic bonding?

Metallic bonding is the strong electrostatic attraction between positive metal ions and delocalised electrons in a giant metallic lattice.

Describe the structure of a metal.

Metals consist of a giant lattice of positive metal ions surrounded by a 'sea' of delocalised electrons that are free to move.

Why do most metals have high melting and boiling points?

Metals have high melting and boiling points because strong electrostatic forces between positive ions and delocalised electrons require a lot of energy to overcome.

Why are metals good conductors of electricity?

Metals conduct electricity because delocalised electrons are free to move and carry charge through the structure.

Why are metals malleable?

Metals are malleable because their atoms are arranged in layers that can slide over each other when force is applied.

Why are metals ductile?

Metals are ductile because their atoms are arranged in layers that can slide over each other, allowing them to be pulled into wires.

What is the difference between conduction in metals and in molten ionic compounds?

Metals conduct electricity via free electrons, while molten or aqueous ionic compounds conduct via free ions.

What are delocalised electrons?

Delocalised electrons are electrons that are not associated with a particular atom and are free to move throughout the structure, as seen in graphite and metals.

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