It's also worth mentioning that 3 of the giant covalent structures in the diagram above only have carbon atoms in them (it does tell you that but this is important). The word allotrope just means that there are the same atom present, but arranged in a different way. You should rejoice if you see a question asking you to compare the properties of allotropes of carbon, all you have to do is compare graphite vs diamond vs fullerene.
Structure and Bonding
Chemists use theories of structure and bonding to explain the physical and chemical properties of materials. Analysis of structures shows that atoms can be arranged in a variety of ways, some of which are molecular while others are giant structures. Theories of bonding explain how atoms are held together in these structures. Scientists use this knowledge of structure and bonding to engineer new materials with desirable properties. The properties of these materials may offer new applications in a range of different technologies.
The Three Types of Chemical Bond
When atoms join together to make molecules there are only 3 possible ways that they can do this. It is essential that you can work out what sort of bonding takes place in a molecule. This can be deduced using this simple rule:
Look at what types of atoms are joining together.
Non-metal and non-metal= covalent bond
Non-metal and metal= ionic bond
Metal and metal= metallic bond
Here's how you can tell if you have a metal or a non metal:
When a metal atom reacts with a non-metal atom electrons in the outer shell of the metal atom are transferred. Metal atoms lose electrons to become positively charged ions. Non-metal atoms gain electrons to become negatively charged ions. The ions produced by metals in Groups 1 and 2 and by non-metals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0). The diagram below shows how a sodium atom (2,8,1) gives it's electron to chlorine (2,8,7) to become a sodium ion (2,8) and a chlorine ion (2,8,8). The sodium ion is positive (it's lost an electron), and the chlorine ion is negative (it's gained an electron). The oppositely charged ions then attract in a giant ionic lattice. Ionic bonding always leads to a giant ionic lattice being formed.
Properties of Ionic Compounds
1. High melting point and boiling point (strong electrostatic forces between ions, and multiple forces of attraction per ion).
2. Do not conduct electricity unless molten or dissolved (need free moving electrons to carry charge.)
The diagram on the right shows what a giant ionic lattice looks like. You should be able to recognise this in an exam. Notice how there's 1 sodium ion to every 1 chloride ion? That is why it has the formula NaCl.
When atoms share pairs of electrons, they form covalent bonds. These bonds between atoms are strong.
Covalently bonded substances may consist of small molecules (called simple covalent structures).
Some covalently bonded substances have very large molecules, such as polymers.
Some covalently bonded substances have giant covalent structures, such as diamond and silicon dioxide.
The diagram below shows three common simple covalent molecules. Notice how the atoms share electrons to mimic a noble gas (i.e. fill their outer shell of electrons). These diagram are called "dot and cross" diagrams, and you will be expected to be able to draw these yourself. Notice how the dots are from one atom, and the crosses are from the other.
Properties of Simple Covalent Substances
Because they exist as molecules, with strong bonds between the atoms, but weak bonds between the molecules, they generally have low melting and boiling points. They can't conduct electricity because they have no delocalised (free) electrons to carry charge. We talk about the intermolecular forces (forces BETWEEN molecules) and their relative strength when discussing simple covalent substances. For example, the intermolecular forces between molecules of oxygen are weaker than the intermolecular forces between molecules of water.
Oxygen boils at at -183oC, whilst water boils at (+)100oC (I've put a plus in front there to emphasise that it is positive, but we don't really need to put a plus sign in front).
You may be asked why chlorine (Cl2) is a gas at room temperature. The answer is simple;
It has weak forces of attraction between molecules (intermolecular forces) and so requires little energy to overcome these.
Giant Covalent Bonding
Certain elements lend themselves to forming multiple covalent bonds, with multiple atoms involved. The most common substances to be aware of are diamond, graphite, graphene, buckminsterfullerene (yes, that IS the real name) and silicon dioxide. These structures have a variety of properties due to their structures. These are summarised in a table below the diagrams.
Metals consist of giant structures of atoms arranged in a regular pattern.
The electrons in the outer shell of metal atoms are delocalised and
so are free to move through the whole structure. The sharing of delocalised electrons gives rise to strong metallic bonds. The bonding in metals may be represented in the following form:
We have a fairly abstract and verbose (wordy) definition for metallic bonding. It's worth learning the diagram above and this definition, because together this can be worth up to 6 marks:
Here we go:
"Positive metal ions held in place by a sea of electrons, acting like glue".
Told you it was very non-scientific sounding!
This type of bonding only occurs between metal atoms, so be careful. Be careful of what I hear you say? Well of the thousands of exam papers I've marked it's noticeable that some people get really confused when it comes to identifying bonding. Use the periodic table to see if you have metals or non metals in your compound before answering. If this doesn't seem like good advice then I suggest you go back to the top of the page and start again.....
Properties of Metals
You need to know what the properties are AND the reasons for them. In case you don't know some of these words below here's a quick reminder:
Malleable= bendable/can be shaped
Ductile= stretchy/can be drawn out into long strands e.g. wires
Lustrous (or Lusterous?) =shiny