Science
Examination Board: WJEC
Module: Chemistry 2 (Booklet 1)
Name:________________________
This module has the following 8 topics:
1. Atomic Structure.
2. Chemical Bonding, Structure and Properties.
3. The Production and Use of Metals.
4. Chemical Calculations.
5. Ammonia and Fertilisers.
6. Alkanes, Alkenes and Polymers.
7. Smart Materials.
8. Water.
Booklet 1 contains notes on topics 1 & 2.
You should read each topic and if you not
understand some ideas, then ask for help
from your chemistry teacher.
1
1. Atomic Structure
The Atom
Electron
Nucleus
Neutron
Proton
An atomic has a nucleus at its centre and shells of electron orbit the
nucleus. The nucleus is made up of protons and neutrons.
Particle
Charge
Relative Mass
Protons
+1
1
Neutrons
0
1
Electrons
-1
0.0005 (negligible)
Atomic Number and Mass Number
Mass Number
12
C
Atomic Number
6
2
Atomic Number is the number of protons in an atom.
There are the same number of electrons as protons.
Mass Number is the number of protons and neutrons in an atom.
e.g. Sodium =
23
Na
11
Protons = 11
Neutrons = 23 – 11 = 12
Electrons = 11
Atoms have no overall electrical charge. This is because the number
of electrons in the shells (orbits) is the same as the number of
protons in the nucleus.
Atoms of the same element have the same number of protons.
Atoms of different elements have different numbers of protons.
2. Chemical Bonding, Structure and Properties
How a material is used depends on its properties (how it behaves).
A materials properties depend on the structure and bonding within
that material.
Scientists are able to design materials with special properties for
particular uses e.g.
 Alloys such as stainless steel
 Polymers such as Kevlar
 Textiles such as polyesters
 Ceramics such as oven to table ware
3
Different materials join in different ways
Ionic Bonding
Metal and non-metal atoms
Covalent Bonding
Non-metal atoms only
Metallic Bonding
Metal atoms only
Giant Metallic Structures
Most elements in the Periodic Table are metals (Remember the step
under Boron, silicon etc.)
Properties of metals
Good conductors of electricity
(high electrical conductivity)
Good conductors of heat
(high thermal conductivity)
Malleable
(can be beaten into sheets)
Ductile
(can be drawn into a wire)
Lustrous (freshly exposed
surfaces are shiny)
High Melting Points
High Boiling Points
Properties of Non-Metals
Poor conductors of electricity
(low electrical conductivity)
Poor conductors of heat
(low thermal conductivity)
Not Malleable
(solids tend to be brittle)
Not Ductile
Not usually lustrous
Low Melting Points
Low Boiling Points
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Metallic bonding
Atoms of metals are tightly packed together in a giant lattice
similar to the lattice in ionic compounds.
metal
atoms
ions
Sea of free
electrons
The outer electrons separate from their atoms and become
delocalized, creating a ‘sea of electrons’. The atoms become
positive ions and are attracted to these electrons. The electrons
act as a sort of glue that holds the structure together.
This attraction is called metallic bonding and is the reason why the
positive metal ions do not repel each other.
This theory of the structure of metals can be used to explain
 Electrical Conductivity – the free electrons can carry an
electrical charge, so a current flows.
 Thermal Conductivity - the free electrons can take in heat
energy, which makes them move faster. They can then transfer
the energy throughout the lattice
 Malleability and Ductility – Metals are usually tough, not brittle.
When a metal is hit, the layers of the lattice just slide over
each other. The electrons act as a lubricant, the metallic bonds
do not break because the electrons are free to move.
force
5
This means that metals are can be bent and pressed into shape
or drawn out into wires.
 High Melting and Boiling Points –These are high because bonding
in metals is strong so lots of energy is needed to break them.
The more electrons the metal has in the outer shell the strong
the bond and the higher the melting and boiling points.
So group 1 metals have lower melting and boiling points than
group 2 and 3.
Group 2 metals are also harder and have a higher mechanical
strength than group 1 metals because they have more free
electrons. Group 3 metals have even higher melting and boiling
pints and are harder than group 2.
Alloys
These are mixtures of metals.
Metals can be made stronger by adding another element when the
metal is molten. The atoms of the new element spread through the
crystal structure.
force
force
Because the atoms of the added element are larger, they make it
difficult for layers of metal atoms to slide. This makes the metal less
malleable and less ductile.
By combining metals with different elements we can make alloys with
properties to suit different purposes. Stainless steel for cutlery and
sinks. Nickel alloys and titanium alloys are commonly used in spectacle
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frames. Some modern frames are also made from smart alloys that
allow the shape to be restored after bending or deformation.
Metallic Glasses
Metals have a regular lattice where the atoms are very ordered.
Metallic glasses are solids which are formed with a far more random
arrangement of atoms. This makes the lattice much less ordered.
These alloys have enhanced mechanical and anti-corrosive properties.
Some alloys have been produced in metallic glass form that have been
used for electronic components, tennis rackets, the aerospace
industry and replacement joints. These often contain zirconium which
was produced in metallic glass form in 2004. Metallic glasses have
improved elasticity and can be three times stronger than steel and
ten times more springy.
Giant Ionic Structures
Ions are electrically charged atoms or groups of atoms (molecules).
 Positvely charged ions are called cations.
They have lost electrons.
 Negatively charged ions are called anions.
They have gained electrons.
e.g.
Positive ions
Sodium Na+, Ammonium NH4+. Magnesium Mg2+, Aluminium Al3+
Negative Ions
Chloride Cl-, Hydroxide OH-, Oxide O2-, Phosphate PO43Atoms lose or gain electrons to reach the electronic arrangement of a
noble gas.
e.g. Sodium 2,8,1 loses 1 electron to become Na+ with the electron
arrangement 2,8 like Neon
Fluorine 2,7 gains 1 electron to become F- with the electron
arrangement 2,8 like Neon
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The sodium ion
Sodium ion:
Sodium atom:
11 protons
11 protons
=+11
=+11
11 electrons =-11
10 electrons =-10
Total charge = 0
Total charge = +1
+
one electron
is lost
Na
Na
Electron arrangement: [2.8]+
(full Outer Shell)
Electron arrangement: 2.8.1
(partially full outer shell)
The fluoride ion
Fluoride ion:
Fluorine atom:
=
+9
9 protons
=
+9
9 electrons =
-9
10 electrons =
-10
Total charge =
0
Total charge =
-1
9 protons
-
F
one electron
is gained
F
Electron arrangement: [2.8](full outer shell)
Electron arrangement: 2.7
(partially full outer shell)
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Ionic Bonding
Sodium Fluoride is an ionic compound formed by the reaction between
the metal sodium and the non-metal Fluorine.
Sodium
Na
+
Fluorine
F
Sodium Fluoride
NaF
During the reaction, one electron is transferred from each sodium
atom to each Fluorine atom.
Sodium has 1 electron
in its outer shell. If it
loses this electron, it
will have no partially
filled shells.
Chlorine has 7 electrons
in its outer shell. If it
gains 1 electron, it will
completely fill its outer
shell.
-
+
F
Cl
Na
2.8.1
Na+
F-
2.8.7
[2.8]+
[2.8.8]-
Positive charges attract negative charges strongly.
The smallest crystal of sodium fluoride contains
millions of sodium and fluoride ions held together in
a regular arrangement or ionic lattice by
electrostatic forces. The ions in sodium fluoride are
arranged in a cubic lattice.
9
Other Ionic Compounds
Sodium chloride and magnesium oxide are simple ionic compounds.
In each case, the metal and non-metal need to lose and gain the same
number of electrons.
Na
Mg
1 electron
2 electrons
Cl
Na+
Cl-
O
Mg2+
O2-
Sodium (2.8.1) needs to lose 1 electron but oxygen (2.6) needs to gain
2 electrons. Therefore, two sodium atoms are required for each
oxygen atom. The formula sodium oxide is Na2O.
Na+
Na
1 electron per
atom
O2-
O
Na+
Na
Magnesium (2.8.2) needs to lose 2 electrons but chlorine (2.8.7) needs
to gain 1 electron. Therefore, two chlorine atoms are required for
each magnesium atom. The formula for magnesium chloride is MgCl2.
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Cl-
Cl
Mg
1 electron for each
atom
Mg2+
Cl-
Cl
Properties of Ionic Compounds
Property
High melting Point
Do not conduct electricity in solid state
Conduct electricity when molten or
dissolved in water
Cause
Forces between the ions are strong
electrostatic forces which need a lot of
energy to break them
Ions are held in fixed positions and are
not free to move
The lattice breaks down and the ions
are free to move and conduct an
electric current
Covalent Molecular Substances
Covalent molecular compounds exist as neutral particles called
molecules. Molecules are formed from atoms sharing electrons. Two
atoms sharing a pair of electrons is called a covalent bond.
The covalent bond
When non-metal atoms react together, they need to gain electrons to
fill their outer shell and become stable.
H
incomplete
outer shells
H
They can only do this if they share electrons with each other.
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both atoms have a full
outer shell
H
H2
H
or
H–H
The atoms share electrons so there is a strong force that joins the
atoms together. This is called a covalent bond. The line between the
two atoms represents the covalent bond (a shared pair of electrons).
The formation of other molecules
Some common covalent molecular compounds represented by dot and
cross diagrams are;
Hydrogen chloride
Cl
H
HCl
or
H Cl
Oxygen
O
O O
O
O2
or
O=O
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Water
O
H
H
H 2O
or
H-O-H
Ammonia
H
H
N
H
NH3
or
H
N
H
H
13
Methane
H
H
H
C
H
H
CH4
or
H
C
H
H
Properties of covalent molecular compounds
Property
Low melting points
Cause
Weak attractive forces between
molecules easily broken
Gases or liquids at room temperature
Weak attractive forces between
molecules easily broken
Do not conduct electricity
Uncharged/No free electrons
Mostly insoluble in water
Remember the covalent bond is very strong and it needs a lot of energy to break
it.
Giant Covalent Structures
Some covalent substances exist as giant structures that have high
melting points because all the atoms are held by strong covalent
bonds. Silicon dioxide, graphite and diamond are examples of giant
covalent structures. Diamond and graphite are both forms of pure
carbon.
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Structures, Properties and Uses of Diamond and Graphite
Diamond – structure
All strong covalent bonds in the
tetrahedrons as all 4 electrons in the
outer shell of each carbon atom are
bonded together.
Appearance
Transparent and crystalline. Used as
gemstone in jewellery.
Extremely hard because all the bonds
are strong covalent bonds. Used to cut
glass & small industrial diamonds are
used in drill bits for oil exploration etc.
Electrical Insulator. No free electrons.
Hardness
Conductivity
Melting Point
Very high over 3500oC due to strong
covalent bonds needing lots of energy
to break them.
Graphite – structure
Strong covalent bonds inside layers of
hexagonal rings. Sea of electrons
between layers as there are only 3 of 4
outer electrons in each carbon atom are
bonded together.
Appearance
Hardness
Conductivity
Melting Point
Grey/black shiny solid
Very soft. Used as a lubricant & as the
“lead” in pencils. Weak forces between
layers are easily broken. Layers slide
over each other making graphite feel
slippery.
It is a non-metal that conducts
electricity.
Very high over 3600oC due to strong
covalent bonds.
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Carbon Nanotubes
 These were discovered in the 1990’s.
 They are like rolled up graphite layers and
they too have electrical conductivity.
 They have extremely fine diameters,about
10,000 times thinner than a human hair.
 Recent research as produced thin walled
tubes inside which metal crystals can be
grown.
 One of the proposed uses is for connections
in miniature electronic circuits. As
electronic circuits get smaller and smaller
conventional connections will not be
practical so nanotube technology may take
over.
The structure of
carbon nanotubes
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