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IGCSE1

� Between metals and non-metals.

� Electrons are transferred from the metal

atoms to the non-metal atoms.

� The atoms get full outer energy levels and

become more stable.

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� Sodium +chlorine � sodium chloride

� The sodium ion has obtained an electron configuration like neon.

� The chlorine atom has an electron configuration like argon.

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� The oppositely charged ions attract each

other and are pulled or bonded together by

strong electrostatic forces.

� This bonding is called ionic or electrovalent

bonding.

1. Draw diagrams to represent the bonding in

each of the following ionic compounds.

a) Magnesium fluoride ������

b) Potassium fluoride ��

c) Lithium chloride ��

d) Calcium oxide � �

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� The ions are packed together in a regular

arrangement called a lattice.

� Here oppositely charged ions attract one

another strongly.

Ionic structures are:1. Solids at room temperature

2. High melting and boiling points (due to strong intermolecular forces).

3. Usually hard substances.

4. Cannot conduct electricity when solid, because ions are not free to move.

5. Mainly dissolve in water (water is able to bond with both negative and positive ions which breaks up the lattice).

6. Usually conduct electricity when molten or in an aqueous solution (ions are free to move).

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� When ionic compounds combine their

respective charges need to balance.

e.g. how many�� ions will need to bond with

an���� ion to balance?

So the formula will be ����

Using the table write the

formula for:

a) Copper (I) oxide

b) Zinc phosphate

c) Iron (III) chloride

d) Lead bromide

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� Between non-metal atoms

� Sharing of electrons

� Each atom wants a noble gas electron

configuration i.e. 8 electrons in its outer shell.

� It gets these by sharing electrons with other

atoms.

� Between non-metal atoms

� Sharing of electrons

e.g. ��

This shared pair of electrons is

called a single covalent bond.

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� Other covalent compounds include:

1. Methane (���)

Methane is tetrahedral

2. Ammonia (���)

Ammonia is pyramidal

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2. Water (���)

Water has a bent or V-shape

2. Carbon dioxide (���)

Carbon dioxide is linear in shape.

� �

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1. Draw diagrams to represent the bonding in each of the following covalent compounds.

a) Tetrachloromethane ��b) Oxygen gas �c) Hydrogen sulphide ���d) Hydrogen chloride ��e) Ethene ���f) Methanol �� �g) Nitrogen ��

2. Why is the water molecule bent?

� Covalent compounds can either be simple molecular or giant molecular.

Simple Molecular compounds� Formed from only a few atoms.

� Strong covalent bonds between atoms in a molecule (intramolecular bonds).

� Weak bonds between the molecules (intermolecular bonds).

� These weak bonds can be Van der Waals bonds. These increase in strength as the molecule gets bigger.

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Giant Molecular compounds

� Formed from hundreds of thousands of atoms.

� Examples are diamond, graphite and silicon (IV)

oxide and plastics such as polythene.

Simple structures � They are usually gases, liquids or solids with low melting and

boiling points.� MP are low due to the weak intermolecular forces of attraction.

Giant structures� High melting points as structure is held together by strong

covalent bonds.

� Generally do not conduct electricity when molten or dissolved in water.*

� *Some molecules react with water to form ions (e.g. HCl) which can then conduct electricity.

� Generally do not dissolve in water.

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� When an element can exist in more than one

physical form in the same state is said to

exhibit allotropy or polymorphism.

� Each physical form is called an allotrope.

� E.g. sulphur, tin, carbon and iron.

� The allotropes of carbon are diamond and

graphite.

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� Graphite is a form of carbon in which the carbon atoms form layers. These layers can slide over each other, so graphite is much softer than diamond.

� It is used in pencils, and as a lubricant. Each carbon atom in a layer is joined to only three other carbon atoms.

� Graphite conducts electricity.

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� Diamond is a form of carbon in which each

carbon atom is joined to four other carbon

atoms, forming a giant covalent structure.

� As a result, diamond is very hard and has a

high melting point.

� It does not conduct electricity.

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Graphite can conduct electricity due to the vast electron delocalization within the carbon layers. These valence electrons are free to move, so are able to conduct electricity. However, the electricity is only conducted within the plane of the layers.

Diamond is a carbon tetrahedral shape and so carbons are linked with each other and are not free to move in plates like graphites can. That is also why graphite is a lubricant and diamond is a cutting abrasive.

� Silica, which is found in sand, has a similar structure to diamond.

� It is also hard and has a high melting point, but contains silicon and oxygen atoms, instead of carbon atoms.

� The fact that it is a semi-conductor makes it immensely useful in the electronics industry: most transistors are made of silica.

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� Buckminsterfullerene is yet another allotrope

of carbon.

� It is actually not a giant covalent structure,

but a giant molecule in which the carbon

atoms form pentagons and hexagons - in a

similar way to a leather football.

� It is used in lubricants.

� Uses of diamond and graphite.

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Soda Glass� Found in bottles and windows � It is made by heating a mixture of sand

(silicon (IV) oxide), soda (sodium carbonate) and lime (calcium oxide).

Borosilicate Glass

� Pyrex is borosilicate glass � boron oxide + silicon (IV) oxide� stronger than soda glass� used in cooking utensils and laboratory

glass ware.

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� Form a giant structure

� Atoms packed closely together in a regular arrangement

� valence electrons tend to move away from their atoms

� These form a “sea” of delocalised electrons which surround a lattice of positively charged metal ions.

� Positively charged ions held

together by the strong

attraction to the mobile

electrons.

� The electrostatic force between

the electrons and the metal ions

acts in all directions.

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1. High melting and boiling points – due to the strong forced of attraction between the metal ions and electrons.

2. Good conductors of electricity – delocalised electrons are free to move.

3. Good conductors of heat – atoms are closely packed and can pass “vibrations” on to the next atom.

4. Malleable (can be bent) and ductile (can be drawn into a wire) – the positive ions are arranged in layers which can slide past eachother when a force is applied.

Why does electrical conductivity of

metals decrease with an increase in

temperature?

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Complete the worksheet handed out to you.

Spare copies can be found on Moodle.