Giant Structures and Metallic Bonding

In giant covalent structures every atom is covalently bonded into one huge network. Diamond (each C bonded to 4 others) is extremely hard with a very high melting point. Graphite has layers of C bonded to 3 others, with delocalised electrons, so it conducts and its layers slide. Silicon(IV) oxide is like diamond. In metals, positive ions sit in a sea of delocalised electrons: they conduct, are malleable and have high melting points.

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Section 1 of 14

Learning objectives

Describe the structures of diamond, graphite and silicon(IV) oxide.
Relate their properties to their uses.
Describe metallic bonding and explain the properties of metals.
Identify the structure of a substance from its properties.

Pause and think

Diamond and graphite are both pure carbon. Why is one the hardest natural material and the other soft enough to write with?

Section 2 of 14

Core idea

Start hereSame atoms, different worlds

A diamond ring and a pencil core are both pure carbon. One can cut glass; the other leaves marks on paper. The difference is entirely in how the atoms are joined.

The big idea

In giant structures, strong bonds run through the whole solid, so melting points are high. Whether they conduct, and whether they are hard or slippery, depends on how the bonds are arranged and whether any electrons are free to move.

1. Giant covalent structures

Diamond: each C bonded to 4 others in 3D. Very hard, very high mp, no conduction. Cutting and drilling tools.
Graphite: each C bonded to 3 others in flat hexagonal layers; one delocalised electron per atom. Conducts; layers slide. Electrodes, lubricants, pencils.
Silicon(IV) oxide: each Si bonded to 4 O, each O to 2 Si. Hard, high mp, no conduction. Sand, glass-making.

Worked exampleGraphite electrodes

Explain why graphite, not diamond, is used for electrodes in electrolysis.

Goal: Link structure to property to use.
Strategy: Compare free electrons.

$\text{graphite: 3 bonds per C}$
One outer electron per atom is delocalised.
$\text{delocalised }e^-\Rightarrow\text{conducts}$
They move through the layers carrying charge.
$\text{diamond: 4 bonds per C}$
No free electrons, so it doesn't conduct.

AnswerGraphite conducts electricity because it has delocalised electrons; it also has a high melting point.

✓ CheckBoth properties are needed for an electrode ✓.

2. Metallic bonding

Metal atoms release their outer electrons into a "sea" shared by the whole lattice. The metallic bond is the attraction between the positive ions and these delocalised electrons.

Explaining metal properties

Drawing diagram…

Every property traces back to the sea of electrons.

Common misconceptionGraphite is soft, not weak The bonds within each graphite layer are stronger than those in diamond. It is soft only because the forces between layers are weak.

Worked exampleIdentify the structure

Substance X melts at $1083\,^\circ\mathrm{C}$ and conducts as a solid and as a liquid.

Goal: Its structure.
Strategy: Use conduction in the solid state.

$\text{conducts as solid}$
Needs delocalised electrons: metal or graphite.
$\text{conducts as liquid, mp }1083\,^\circ\mathrm{C}$
Typical of a metal (it is copper).

AnswerMetallic.

✓ CheckAn ionic solid wouldn't conduct; a simple molecule would melt far lower ✓.

Pause and think

Why is a metal a good conductor of heat as well as electricity?

Summary

Diamond and $\mathrm{SiO_2}$ : hard, non-conducting.
Graphite: layers, conducts, slippery.
Metals: conduct, malleable, high mp.

Giant covalent: A huge network of atoms joined by covalent bonds.
Delocalised electrons: Electrons free to move through a structure.
Malleable: Can be hammered into shape without breaking.

HistoryFullerenes. In 1985 Kroto, Curl and Smalley discovered buckminsterfullerene,

\mathrm{C_{60}}

, a football-shaped molecule of carbon. They won the 1996 Nobel Prize in Chemistry.

Fun factPencils are graphite. "Lead" pencils have never contained lead: the core is graphite mixed with clay. Each line you draw leaves layers of graphite on the paper.

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