Covalent Bonding

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  • Created on: 13-05-15 12:10

Covalent Bonding

1. Molecules are the smallest parts of compounds that can take part in chemical reactions.

2. They're formed when two or more atoms bond together - it doesn't matter if the atoms are the same or different. Chlorine gas (Cl2), carbon mooxide (CO), water (H2O) and ethanol (C2H5OH) are all molecules.

3. Molecules are held together by strong covalent bongs.

In covalent bonding, two atoms share electrons, so they've both got full outer shells of electrons. Both the positive nuclei are attracted electrostatically to the shared electrons.

Here's some examples. These diagrams don't show all the electrons - just the ones in the outer shells.


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Double and Triple covalent bonds

Atoms don't just form single bonds - double or even triple covalent bonds can form too.

It's not over yet - the typical properties of simple covalent moleculaes are covered later on.

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Giant Covalent Structures

1. Giant covalent structures have a huge netowrk of covalently bonded atoms. (They're sometimes called macromolecular structures).

2. Carob atoms can form this type of structure because they can each form four strong, covalent bonds. There are two types of giant covalent carbon structure you need to know about, graphite and diamond.

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Graphite's structure means it has certain properties.

1. The weak bonds between the layers in graphite are easily broken, so the sheets can slide over each others - graphite feels slippery and is used as a dry lubricant and in pencils.

2. The 'delocalise' electrons in graphite aren't attatched to any particular carbon atoms and are free to move along the sheets, so an electric current can flow.

3. The layers are quite far apart compared to the length of the covalent bonds, so graphite has a low density and is used to make strong, lightweight sports equipment.

4. Because of the strong covalent bonds in the hexagon sheets, graphite has a very high melting point (it sublimes at over 3900K)

5. Graphite is insoluble in any solvent. The covalent bonds in the sheets are too difficult to break.


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Diamond is also made up of carbon atoms. Each carbon atom is covalently bonded to four other carbon atoms. The atoms arrange themselves in a tetrahedral shape - its crystal lattice structure.

Because of the its strong covalent bonds:

1. Diamond has a very high melting point - it actually sublimes over 3800K.

2. Diamond is extremely hard - it's used in diamond-tipped drills and saws.

3. Vibrations travel easily through the stiff lattice, so it's a good thermal conductor.

4. It can't conduct electricity - all the outer electrons are held in localised bonds.

5. Like graphite, diamond won't dissolve in any solvent.


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Dative covalent bonding

The ammonium ion (NH4+) is formed by dative covalent (or coordinate) bonding - it's an example the examiners love. It forms when the nitrogen atom in an ammonia molecule donates a pair of electrons to a proton (H+).


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