IONS
Formation of Ions
 Ions are atoms which have gained or lost electrons to obtain a full outer shell.
 Ions are always charged, because electrons have a negative charge.
 Atoms that lose electrons form positive ions because they now have more
protons than electrons
 Positive ions are also known as cations (pronounced ‘cat irons’. See diagram,
above!)
 Metals tend to form positive ions (cations)
 Loss of electrons is known as OXIDATION
 Atoms that gain electrons form negative ions because they now have more
electrons than protons
 Negative ions are also known as anions
 Non-metals tend to form negative ions (anions)
 Gain of electrons is known as REDUCTION
e.g. chlorine
Electron configuration is 2,8,7
To gain a full outer shell, it can either gain one electron or lose seven electrons.
It is easier to gain one electron than to lose seven, so chlorine gains an electron to
form an anion with a charge of -1. The electron will come from another element
during chemical reaction. Chlorine is said to have been reduced.
The ion formed is called chloride and has the formula Cl-.
The ion still has 17 protons and 18 neutrons (the same as the chlorine atom), but it
now has 18 electrons, and it now has electron arrangement [2,8,8]-.
Work out the electron arrangement and the formula of the ion formed by
sodium:
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IONIC BONDING
Ionic bonding in sodium chloride
Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it
gave away that electron it would become more stable.
Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it
could gain an electron from somewhere it too would become more stable.
The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both
become more stable.
The sodium has lost an electron, so it no longer has equal numbers of electrons and
protons. Because it has one more proton than electron, it has a charge of 1+. If
electrons are lost from an atom, positive ions are formed. Positive ions are
sometimes called cations.
The chlorine has gained an electron, so it now has one more electron than proton.
It therefore has a charge of 1-. If electrons are gained by an atom, negative ions
are formed. A negative ion is sometimes called an anion.
The nature of the bond
The sodium ions and chloride ions are held together by the strong electrostatic
attractions between the positive and negative charges.
The formula of sodium chloride
You need one sodium atom to provide the extra electron for one chlorine atom, so
they combine together 1:1. The formula is therefore NaCl.
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Some other examples of ionic bonding
magnesium oxide
Again, noble gas structures are formed, and the magnesium oxide is held together
by very strong attractions between the ions. The ionic bonding is stronger than in
sodium chloride because this time you have 2+ ions attracting 2- ions. The greater
the charge, the greater the attraction.
The formula of magnesium oxide is MgO.
calcium chloride
This time you need two chlorines to use up the two outer electrons in the calcium.
The formula of calcium chloride is therefore CaCl2.
Ionic bonding is often represented by showing the electronic structure of the ions
like this
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IONIC COMPOUNDS
When Ionic compounds are formed large numbers of these ions pack together in a
regular pattern in a giant structure, called an IONIC LATTICE. They are held
together by electrostatic attraction between the opposite electrical charges. The
electrostatic attractions are strong and operate equally in all directions
throughout the crystal. This attraction is the ionic bond.
An example is the sodium chloride lattice, which is shown as an ‘exploded’ view
below.
Na+
ionCl ion
Each sodium ion is in contact with six chloride ions; each chloride ion is in contact
with six sodium ions.
In reality, the ions are touching, as shown in the following diagrams.
single layer
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Properties of ionic compounds
1. The strong forces which hold the lattice together must be over come before
the solid will melt. This requires a high temperature. Therefore, ionic substances
are solids at room temperature and have high melting points and high boiling
points.
2. Ionic compounds can conduct electricity if their ions are free to move. In the
solid, the ions are rigidly held in the lattice, so these substances do not conduct
as solids. The ions are free to move in the liquid state (when the lattice has been
broken down by heat) or in solution (when the lattice has been broken down by
dissolving). Ionic substances therefore conduct when molten or in solution.
The ions are attracted to the electrode of opposite charge and on reaching it
undergo a chemical change. This is called ELECTROLYSIS. Positive ions gain
electrons at the cathode (reduced) and negative ions lose electrons at the anode
(oxidised).
3. Most ionic compounds are soluble in water.
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Determining simple ion formulae using the Periodic Table
The chemical properties of an element depend on the number of electrons in the
highest energy level. Once we know the electronic structure (the electron
arrangement) of an element, we can work out its Group. We can then predict its
chemical properties.
Remember:
 metals lose electrons to form cations. They get oxidised.
 non-metals gain electrons to form anions. They get reduced.
An easy way to remember the charge on an ion is given in the table below:
Group
1
2
3
Charge on
Ion
+1
+2
+3
4
no
ions
5
6
7
-3
-2
-1
0
no
ions
Other common (but scarier looking) types of ions
Some common ions are more complicated than the ones we have met so far because
they contain more than one element. The formulas and charges of these ions should
be learned by heart because they cannot be worked out simply by looking at the
Periodic Table:
Ion formula and charge
Name of Ion
NH4+
OHCO32SO42NO3PO43-
ammonium
IGCSE TOPIC 10.4: ATOMS COMBINING
hydroxide
carbonate
sulphate
nitrate
phosphate
6
Atomic Structure Worksheet 2
Use your Periodic Table to complete the following table. Beware of the ions hiding among
the atoms....
Particle
Mass
Number
Atomic
number
P
31
15
K
F
Number
of protons
Number of
electrons
19
19
9
Cl-
35
17
Na+
23
11
H
Number of
neutrons
9
2,8,8
1
1
Ar
18
22
O2-
10
8
H+
1
1
Ca2+
N3-
20
14
Electronic
Structure
18
7
Al3+
10
Li+
3
C
6
6
S2-
16
18
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2
7
COVALENT BONDING




Non-metal atoms need to gain electrons to achieve a full outer shell.
When two non-metals bond, neither one will give up an electron, so the atoms
share a pair of electrons.
This means that each atom has a full outer shell, despite some of the
electrons in the outer shell being shared with another atom.
A covalent bond is defined as a shared pair of electrons.
Here are two hydrogen atoms forming a covalent
bond to become H2.
Each H atom only has one electron. As it only has
other H atoms to bond with, each H atom pairs with
another to share their electrons, so each can have
2 in the shell making it full and therefore stable.
It now exists as a hydrogen molecule (H2). It is
very difficult to split them up again (i.e. to break
the covalent bond) but very easy to separate one
H2 molecule from another. Hence, it has a low
boiling point so at room temperature it exists as a
gas.
A covalent bond can be represented by a dot and cross diagram. ONLY THE
OUTER, BONDING ELECTRONS NEED TO BE SHOWN.
Steps: 1. Decide how many electrons are in the outer shells of the atoms in the
formula (remember the Periodic Table shows this by the Group in which the
element sits).
2. Decide how many electrons are needed to fill the shell. This will equal the
number of covalent bonds or shared pairs needed.
3. Draw the atoms overlapping and making the electron from one as dots, the other
as crosses and place the correct number of dot/cross pairs in the overlap.
4. Count the number of electrons in each atoms shell including those in the overlap
and check the shells are now all full. If they are not then increase the sharing
until they are.
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Oxygen, O2
Chlorine, Cl2
Covalent bonding can also occur in compounds (i.e. 2 or more different nonmetals).
Hydrogen chloride, HCl
Carbon dioxide, CO2
Methane, CH4
Water, H2O
Ammonia, NH3
Nitrogen, N2
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The dot and cross diagram can be simplified, and the covalent bond shown as a
single line between two atoms. Every time you see a diagram with a line of this
sort, it means the atoms are covalently bonded together. For the molecules above
they can be drawn as:
H
H-Cl
O=C=O
H–C–H
H
H-O-H
H-N-H
H
–
N –– N
(3 lines show 3 pairs
of shared electrons,
called a triple bond)
Oxygen (O2) is shown as O=O (2 lines show 2 pairs of shared electrons, called a
double bond).
Task: to draw dot and cross diagrams of the following simple covalent molecules:
Hydrogen fluoride (HF); Fluorine (F2); Ethane (C2H6).
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SIMPLE MOLECULAR SUBSTANCES
Molecular substances are formed when two or more atoms of non-metallic elements
share pairs of electrons. Molecular substances may be elements (e.g. Cl 2 O2) or
compounds (e.g.H2O NH3). Molecular substances consist of individual molecules in
which the atoms are joined together by strong, covalent bonds. The separate
molecules are held together by weak attractive forces called van der Waals’ forces.
INTERMOLECULAR FORCES
These weak forces of attraction between neighbouring molecules are easily overcome
which is why simple molecules like O2, CO2, CH4 are gases at room temperature (they
have already boiled!)
Weak forces between molecules
Strong covalent bonds
Covalent bonds are directional, so individual molecules have a definite threedimensional shape. For example:
..
..
methane
(tetrahedral)
ammonia
(pyramidal)
..
..
water
(V-shaped)
Properties of simple molecules
1. Before the solid can melt, the weak van der Waals’ forces between molecules must
be overcome. Since this requires very little energy, some molecular substances melt, or
even boil, below room temperature. Molecular substances are usually gases, liquids or
low-melting solids, i.e. they have low melting and boiling points.
2. Simple molecular substances do not conduct electricity, because they have no
charged particles which are free to move. Their electrons are involved in bonding or are
lone pairs; molecules do not have an overall electric charge.
3. Most covalent substances are insoluble in water. They will, however, dissolve in
organic solvents such as petrol and propanone.
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MACROMOLECULAR SUBSTANCES
Atoms which share electrons can also form giant structures. Macromolecular
substances have continuous structures of atoms which are bonded by covalent
bonds. Each piece is one enormous molecule which extends indefinitely in two or
three dimensions. The covalent bonds are strong and exist throughout the crystal.
Examples: quartz, diamond, graphite.
Quartz is silicon dioxide SiO2. Diamond and graphite are both forms of carbon;
they contain only carbon atoms, but the atoms are bonded together differently.
Properties
1. Before the solid can melt, a large number of strong covalent bonds, which hold
the crystal together, must be overcome. This requires a lot of heat energy.
Therefore, macromolecular substances are solids at room temperature and have
very high melting and boiling points.
2. Macromolecular substances do not conduct electricity, because they contain no
ions.
An important exception is GRAPHITE, which can conduct electricity as a solid
because it contains electrons which are free to move.
3. Macromolecular substances are usually insoluble in water and insoluble in all
other solvents.
GRAPHITE
Graphite is an infinite two-dimensional
arrangement of atoms. It consists of
layers of carbon atoms in which each
carbon atom is covalently bonded to
THREE others. The layers are held
together by weak van der Waals’
forces.
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Properties of graphite
1. Graphite is a dark grey, greasy solid with a dull shine.
2. Graphite is soft and slippery and can be used as a lubricant. Since the forces
between the layers are weak, they can be easily overcome. This allows the layers of
atoms to slide over each other easily.
3. Graphite is a good conductor of electricity. This is because of electrons which
are free to move within a layer.
DIAMOND
Diamond is a giant, rigid, three-dimensional
structure of carbon atoms and each carbon atom is
bonded covalently to FOUR others.
Properties of diamond
1. Diamond is a hard, colourless solid which sparkles in the light. It is the hardest
naturally occurring substance known and is used to tip cutting tools such as drill
bits to make them long-lasting.
2. Since the strong bonds make its structure so rigid and hard to break down,
diamond has a very high melting point (3350oC)
3. Diamond does not conduct electricity, because there are no ions or free
electrons in it to carry charge.
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SILICON DIOXIDE
Silicon Dioxide occurs naturally as the mineral quartz. This is the main component
of sand. It is a very similar macromolecule to diamond and is also very hard, has a
high melting point and does not conduct electricity.
= Si
=O
METALS
Metals have giant structures in which their atoms are tightly packed together in a
regular pattern. The electrons in the highest occupied energy levels (outer shells)
are separated from their atoms and are free to move through the whole structure.
Metals are usually described as a lattice of ions in a ‘sea’ of electrons.
+
+
+
+
+
+
+
+
+
+
+
+
+
+
free electron
+
metal ion
The metal ions are held together in a regular structure by their attraction to the
electrons between them.
Properties of metals
1. Metals are good conductors of electricity, even when solid, because the free
electrons can move through the lattice, carrying charge.
2. Metals are good conductors of heat.
3. Most metals can be hammered into different shapes (malleable) or drawn into
wires (ductile), because the layers of atoms can slide over each other.
4. Most metals have high melting points, because it requires a large amount of
energy to break up the lattice. Important exceptions are the group 1 metals (e.g.
sodium melts at 98oC) and mercury, which is liquid at room temperature.
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Summary Questions
Topic 4: Atoms Combining
1/
2/
3/
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4/
5/
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6/
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7/
8/
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9/
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