Structure and Bonding - Diamond and Graphite

Giant Covalent Structures

Carbon atoms can join together using strong covalent bonds. When these bonds connect throughout an entire solid, they create what we call a giant covalent structure. Diamond and graphite are both made entirely of carbon atoms, but they have completely different properties because their atoms are arranged and bonded in different ways.

Think of it like building with the same LEGO bricks - you can make very different structures depending on how you connect them together.

Key features of giant covalent structures

• Each atom bonds to several neighbouring atoms by sharing pairs of electrons

• The covalent bonds are very strong, so you need lots of energy to break them - this gives very high melting points

• How the atoms are arranged controls properties like hardness and electrical conductivity

Diamond

In diamond, every carbon atom forms four single covalent bonds arranged in a tetrahedral shape (like a pyramid with a triangular base). This creates a rigid three-dimensional network that extends throughout the entire crystal.

• Each C–C bond is identical and extremely strong

• There are no free electrons that can move around

• All bonds point in different directions, creating a stiff, interlocking structure

Properties of diamond explained

Hardness: Diamond is the hardest natural substance because to break or scratch it, you must break many strong covalent bonds in all directions at once.

Very high melting point: To melt diamond, you need to break apart the entire three-dimensional network, which requires enormous amounts of energy.

Does not conduct electricity: All four outer electrons from each carbon atom are used up in the four covalent bonds, so no electrons are free to carry electric current.

Conducts heat well: Vibrations (heat energy) travel quickly through the rigid, tightly-connected structure.

Graphite

Graphite has a completely different structure. It contains layers of carbon atoms arranged in flat sheets. Each carbon atom forms only three covalent bonds, creating hexagonal patterns that look like chicken wire or honeycomb.

• The bond angle is 120°120°, giving each sheet its hexagonal pattern

• Only three of carbon's four outer electrons are used in bonds - the fourth electron becomes delocalised, meaning it can move freely across the whole layer

• The layers are held together by weak intermolecular forces (van der Waals forces), not covalent bonds

Properties of graphite explained

Soft and slippery: The layers can slide over each other easily because only weak forces hold them together. This makes graphite useful as a dry lubricant and explains why pencils work - layers rub off onto paper.

High melting point: Even though the layers slide easily, the strong covalent bonds within each layer still need lots of energy to break.

Conducts electricity: The delocalised electrons can move freely within the layers, allowing electric current to flow - similar to how metals conduct.

Lower density than diamond: The layers are spaced further apart than the tightly-packed atoms in diamond's structure.

Comparing diamond and graphite

Diamond is a three-dimensional network with all electrons locked in bonds. Graphite is made of two-dimensional layers with mobile electrons. These structural differences explain why two materials made of the same element can have such different properties.

Key terms

Covalent bond - A chemical bond formed when two atoms share a pair of electrons

Delocalised electron - An electron that is not attached to one specific atom or bond and can move freely through a structure

Giant covalent structure - A large network where atoms are connected by covalent bonds throughout the entire material

Van der Waals forces - Weak attractive forces between molecules or layers of atoms

Worked example

Question: Diamond does not conduct electricity. Explain why.

Solution:

1. Identify the bonding: Each carbon atom in diamond forms four covalent bonds with neighbouring carbon atoms

2. Account for all electrons: All four outer electrons from each carbon atom are used up in forming these four covalent bonds

3. Link to electrical conduction: For a material to conduct electricity, it needs mobile charge carriers such as delocalised electrons or ions that can move freely

4. Draw conclusion: Since no electrons are free to move in diamond's structure, it cannot conduct electricity

Graphite in pencils: When you write with a pencil, thin layers of graphite break off and stick to the paper. The delocalised electrons in these graphite layers absorb light, which is why pencil marks appear dark grey or black. The "lead" in pencils is actually graphite mixed with clay - more clay makes a harder pencil (H grades), while more graphite makes a softer, darker pencil (B grades).

Aim: To investigate the electrical conductivity of different carbon structures

Apparatus:

• Low voltage power supply (6-12V)

• Small bulb and holder

• Crocodile clips and connecting wires

• Diamond sample (or cubic zirconia as substitute)

• Graphite rod (from inside a pencil)

• Ammeter (optional)

Method:

1. Set up a simple circuit with the power supply, bulb, and crocodile clips

2. Place the diamond sample between the crocodile clips and observe the bulb

3. Replace the diamond with a graphite rod and observe any changes

4. If available, use an ammeter to measure the current flowing through each material

Safety:

• Use low voltage (less than 12V) to prevent overheating

• Handle materials carefully to avoid breakage

Observations:

• Diamond: No current flows, bulb remains off

• Graphite: Current flows, bulb lights up

Conclusion: Graphite conducts electricity due to delocalised electrons, while diamond does not conduct because all electrons are fixed in covalent bonds

Comparison table

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Advanced structural details: In graphite, the delocalised electrons form what chemists call a "sea of electrons" above and below each layer. These electrons can move freely within the plane of each layer but cannot easily jump between layers. This explains why graphite conducts electricity much better along the layers than perpendicular to them.

The weak van der Waals forces between graphite layers (about 0.3 nm apart) are approximately 100 times weaker than the covalent bonds within layers. This huge difference in bond strength explains the dramatic difference in properties within layers versus between layers.

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