3.1.1 (e) Giant Covalent Lattices

Syllabus

(e) explanation of the solid giant covalent lattices of Carbon (Diamond, Graphite and Graphene) and Silicon as networks of atoms bonded by strong covalent bonds

{Use of ideas about bonding to explain the strength and conductive properties of graphene, and its potential applications and benefits.}

What does this mean?

A few elements (and also a few compounds that aren't on the syllabus) create Giant Covalent Lattices.

That is, they bond covalently but do not form small molecules.

Instead they form enormous structures containing billions of atoms in a regularly repeating pattern.

The two most familiar will be allotropes of Carbon.

Diamond

Carbon has four unpaired valence (outer) electrons.

So it can form a maximum of 4 covalent bonds.

With no lone-pairs to worry about these bonds are perfectly tetrahedral in arrangement at 109.5^o

And because Carbon is a small atom these will be short bonds.

Short bonds are strong - making Diamond the hardest natural substance and giving it a high melting point.

This means it can scratch away any other substance without melting - hence its use in drills and saws.

It does not mean it can put up with being hit with a hammer!

Since all the electrons are localised in bonds there is nothing to carry charge through the solid.

Hence, it is a terrible conductor of electricity.

But its very regular lay out makes it transparent.

In theory, compressing any sample of Carbon should create a diamond.

But it's very expensive to create gem quality diamonds artificially.

And imperfect diamonds are fairly common - which is why they are cheap enough for industrial use.

Graphite

Despite Carbon having 4 valance electrons the Graphite structure uses only 3 of them in covalent bonds.

This forms flat layers of hexagonal carbon atoms.

The spare electron has no fixed place - it is delocalised between the layers.

This accounts for Graphite's conductivity.

The layers are held together by forces that we can think of as van der Waals.

These are strong enough to hold the solid together but weak enough to allow layers to slide over each other.

This makes Graphite a useful lubricant in places where oil cannot be used.

It also makes them useful for pencils.

The arrangement is less regular than diamond so it is opaque.

You might think that the melting point would be a lot lower than diamond's but this isn't true.

It is still necessary to break many strong covalent bonds before graphite melts so the melting point is still very high.

Graphene

Graphene is just a single sheet of hexagonally bonded carbon atoms.

Because it is so thin it is basically transparent - which must make working with it very difficult.

So its shares a lot with Graphite - from which it was first made.

A single sheet of Graphene has an incredible strength to weight ratio, close to 100 timers higher than that of Steel.

It also has interesting potential for electrical uses as the electron not used in bonding is still delocalised.

Silicon

Like Carbon, Silicon is in Group 4.

So it also has four bonding electrons.

Which makes its structure very similar to diamond.

So it is still hard and difficult to melt but because the atoms are bigger the bonds are longer and weaker.

Like diamond, pure Silicon is an electrical insulator because all its valence electrons are used up in bonding.

To make semi-conductors for computer chips Silicon is "doped" with Boron/Gallium that have only 3 valence electrons and creates a "positive hole", or Arsenic/Phosphorus that has a spare electron.

Silicon Oxide

Not to be confused with Silicon is Silica, or Silicon Oxide.

This isn't actually on the syllabus any more but examiners sometimes like to ask questions about substances that are similar to ones that you should know about.

And since Silica used to be a regular question it wouldn't be surprising to see it again.

In the structure you can see that the Silicon atoms again form 4 bonds tetrahedrally.

And the there is an Oxygen between each Silicon atom.

The Si:O mole ration is 1:2 so it is sometimes called Silicon Dioxide - although this rather implies a simple molecular structure like Carbon

Dioxide- which is clearly not the case.

We used to write the formula as n(SiO₂) to imply the Giant Structure.

But examiners seem happy with SiO₂ now.

Sloppy really!

As you'd expect: High melting point,insulator.

Yawn!

Boron Nitride

Now this has never been on the syllabus.

But that didn't stop the examiners asking about it before, so why wouldn't they do it again some time?

The idea is that you should be able to apply your knowledge to see how this lattice might work and how it would be similar or different to lattices that you're supposed to know about.

Clearly Boron has three valance electrons because its in Group 3.

Nitrogen is in Group 5, so it has a lone pair and three electrons available to bond.

So you might assume that there would be some sort of lattice with three single bonds around every atom.

But Boron is electron-deficient - even when it forms three bonds its outer-shell only has six electrons.

So it would be better for it to accept the lone pair from Nitrogen and allow a dative covalent bond to form.

You can see why an examiner might want to ask you to work this out.

But what structure would it take up?

Because there are four covalent bonds around each atom it could take up the diamond structure - and therefore have similar properties to diamond.

An examiner might like to throw in the phrase isoelectric and ask you what you think it means - it means it has the same electronic structure as the Carbons in diamond.

Or it could take up a graphite structure - but the difference would be that there are no spare electrons to dissociate because the Boron would still need to accept the lone pair from Nitrogen to fill its outer shell.

So a double bond would form.

And the substance would be an electrical insulator despite being similar to Graphite - though it has other possible structures too.

Videos

Exam-style Questions

1. The diagram below represents a section of a crystal of Silicon Dioxide.

(a) Name an element which has a structure similar to this.

……………………………………………………………………………………......................................................................................

(1)

(b) Name the type of bonding between Silicon and Oxygen in this crystal

……………………………………………………………………………………......................................................................................

(1)

(c) Name the type of structure illustrated by this diagram

……………………………………………………………………………………......................................................................................

(1)

(d) Describe the motion of the atoms in this crystalline solid

………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………................................................................................................

(2)

(e) In terms of structure and bonding, describe what happens to the atoms in this crystal when it melts

………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………................................................................................................

(2)

(f) Explain why this crystal is a non conductor of electricity in the solid state and why graphite is a good conductor.

………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………..............................................

(4)

Answers

1(a) Carbon (1)

(b) covalent (1)

(c) macromolecule / Giant Covalent(1)

(d) vibrate (1)

about a fixed position (1)

(e) bonds break (or loosen) (1)

atoms (or molecules) free to move (1)

(f) electrons not free in Silicon Dioxide (1)

delocalised electrons in graphite (1)

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