1e – Bonding and Structure
Bonding and Structure
As you have seen in Activity 7, compounds can be separated into two types, known as
covalent and ionic based on the type of chemical bond that causes their formation. The
only true way to classify a compound as covalent or ionic compounds is by its ability to
conduct electricity.
Covalent compounds cannot conduct electricity in any state.
Ionic compounds can conduct electricity but only when liquid or in solution.
Covalent Bonds
A covalent substance is one where atoms are joined together by one or more covalent
bonds, which is formed when two atoms share a pair of electrons. This is done in order
for each atom to have an electron arrangement similar to that of a noble gas, with a full
outer electron shell.
It is not only compounds that can be described covalent. Some elements exist as
molecules where their atoms are joined together by covalent bonds.
Take hydrogen as an example, it has one proton in its nucleus and one electron orbiting
the nucleus in the first electron shell. As we know the first shell can hold a maximum of
two electrons before it is considered full.
Each hydrogen atom can share its one electron with another hydrogen atom, this means
that each hydrogen effectively has two electrons, filling its outer shell. This shared pair
of electrons is what is known as a covalent bond.
A covalent bond is a shared pair of electrons between two atoms
This can be shown by drawing a diagram showing the outer electrons and how they
overlap:
Covalent bond
H
+
H

Page 35
H
H
1e – Bonding and Structure
There are also examples of substances where two or more covalent bonds form between
the same two atoms, these are known as double and triple covalent bonds.
For example oxygen and nitrogen both exist as diatomic molecules, nitrogen has five
outer electrons and oxygen has six outer electrons.
When an oxygen atom forms a diatomic molecule it must share two of its electrons with
another oxygen atom, this forms a double covalent bond.
Double covalent bond
+
O

O
O
O
When a nitrogen atom forms a diatomic molecule it must share three of its electrons
with another oxygen atom, this forms a triple covalent bond.
Complete this diagram by adding the electrons to show a triple covalent bonded
nitrogen molecule:
N
+
N

N
Draw another diagram for the formation of hydrogen chloride:
Page 36
N
1e – Bonding and Structure
Valency for Covalent Substances
When covalent compounds are formed, atoms share electrons with one another in order
to achieve a full outer shell. Every electron shared becomes part of a pair (a covalent
bond), which allows atoms to gain the share of electrons from other atoms until a full
outer shell is achieved.
Since covalent compounds normally form between non-metal elements we will examine
groups 4 to 0.
Complete this table:
Group
4
5
6
7
0
Example element
C
N
O
F
Ne
Outer electrons
5
Electrons needed
to complete outer
shell
3
Valency Number
3
Diatomic Elements
When two atoms join together they form a molecule. A molecule can be of an element,
where all atoms are the same, or a compound, where the atoms are different.
There are seven elements that, under normal conditions, exist as
“diatomic molecules”. This means that two of their atoms join together
to form a molecule
When these seven are present as an element their formula is written to reflect the fact
that they exist as diatomic molecules.
You must learn these seven elements, they are:

Hydrogen, H2

Chlorine, Cl2

Oxygen, O2

Bromine, Br2

Nitrogen, N2

Iodine, I2

Fluorine, F2
Page 37
1e – Bonding and Structure
Attraction and Repulsion
When any two charged particles are close together they will interact with one another,
either attracting or repelling each other.
If both particles are positive or both are negative then they will repel one another.
e.g. two protons will repel one another.
If the particles have opposite charges then they will attract one another.
e.g. a proton and an electron will be attracted to one another.
Like charges repel and opposite charges attract
So when two atoms come together, how can the covalent bond hold the two atoms
together? Why don’t the two positive nuclei repel one another and fly apart?
This is because a balance is reached between attraction and repulsion. The two positive
nuclei will repel one another but they both have a shared attraction for the pair of
bonding electrons. As a result the two atoms are held in a stable formation.
Complete the diagram, showing the repulsion of the two nuclei and their shared
attraction for the bonding electrons:
+
+
Page 38
1e – Bonding and Structure
Covalent Structures
As we have seen, some compounds exist as covalent molecular compounds. These
molecules have very weak forces of attraction between the molecules but strong
covalent bonds inside the molecules.
When melting or boiling a covalent
molecular substance it is only the
weak forces of attraction that must
be broken, not the strong covalent
bonds.
As a result covalent molecular
compounds have low melting and
boiling points and can exist as solids,
liquids
and
gases
at
room
temperature.
Some covalent substances exist as giant three-dimensional covalent networks. This is
where the atoms are joined together by covalent bonds but rather than being in
molecules the atoms bond with one another in a continual way.
Carbon can exist in many different forms, one of which is
diamond. Diamond is a three-dimensional network of carbon
atoms where each carbon has four covalent bonds joining it to 4
different carbons.
When melting or boiling a covalent network substance it is
only the strong covalent bonds that must be broken.
As a result covalent network compounds have very high melting
and boiling points and can only exist as solids at room
temperature.
Page 39
1e – Bonding and Structure
Molecule v Network
Physical Properties
Carbon and silicon are both found in group four of the periodic table and both form
covalent compounds with oxygen. Let’s compare carbon dioxide with silicon dioxide:
Carbon dioxide
Silicon dioxide
Exists as covalent molecules
Exists as a covalent network
Gas at room temperature
Solid at room temperature
Melting Point: -78 °C
Melting point: 1700 °C
Boiling Point: -57 °C
Boiling point: 2230 °C
Chemical Formulae
For covalent molecular substances the chemical formula tells us the number of atoms of
each element present in the molecule
e.g.
carbon dioxide is CO2 this tells us that there is one carbon atom and two oxygen atoms
in the molecule.
Water is H2O and this tells us that there are two hydrogen atoms and one oxygen atom in
the molecule.
For covalent network substances the formula tells us the ratio of the atoms in the
network, not the definite number present.
e.g. silicon dioxide is SiO2, this tells us that there is one silicon atom for every two
oxygen atoms in the network.
Page 40
1e – Bonding and Structure
Summary – Covalent Substances
covalent
network (x3)
conduct
pair
solid
molecular (x3)
electrons (x2)
elements
attraction
ratio
Covalent substances can be either ____________________ or compounds.
They are
formed when atoms form ____________________ bonds. A covalent bond is described as
a shared ____________________ of ____________________ between two atoms.
A
covalent bond holds the atoms together due to the two nuclei having a shared
____________________ for the pair of ____________________.
Covalent compounds can be easily distinguished from ionic compounds, as they cannot
____________________ electricity in any state.
A covalent substance can form one of two structures, it can be either a covalent
____________________
or
____________________.
These
two
structures
are
distinguished by their physical properties; covalent ____________________ have low
melting and boiling points and can exists as solids, liquids or gases at room temperature,
while covalent ____________________ have very high melting and boiling points and
only exist as ____________________ at room temperature.
The formula of a covalent ____________________ substance gives the exact number of
atoms present inside the molecule, whilst a covalent ____________________ structure
gives the ____________________ of atoms in the network.
Page 41
1e – Bonding and Structure
Ionic Bonding
An ionic substance is one that is made up, not of atoms, but of ions. As covered
previously an ion is a charged particle, formed when an atom either gains or loses
electrons.
Similarly to covalent bonding, atoms need to
achieve a full outer electron shell, similar that
of their nearest noble gas. Unlike covalent
bonding, however, no electrons are shared to
form a bond. Instead one atom will donate
electrons to the other forming a positive ion
(having lost electrons) and a negative ion
(having gained electrons).
An ionic bond is simply an electrostatic attraction between a positive and a negative ion.
Take sodium chloride as an example. Sodium has an electron arrangement of 2,8,1 and
chlorine has an arrangement of 2,7. Sodium needs to lose one electron to achieve an
arrangement of 2,8 and chlorine needs to gain one electron to also achieve an
arrangement of 2,8.
1e
e.g.
+
+
Empty shell
Cl
Na
+
+
attraction
Full outer shell
Na+
Cl-
Another key difference from covalent bonding is that ionic bonding can only occur in
compounds, not in elements.
Page 42
1e – Bonding and Structure
Ionic Structure
The structure adopted by ionic compounds is known as a crystal lattice structure. Each
ion is surrounded by the ion of opposite charge forming a very rigid and well-ordered
structure.
Different compounds will form lattices of different shapes and it is this arrangement of
ions that causes ionic compounds to form crystals of various shapes.
Sodium chloride forms a cubic shaped crystal lattice where every sodium ion and every
chloride ion is surrounded on six sides by an oppositely charged ion
When melting or boiling an ionic substance it is
the strong ionic bonds that must be broken.
As a result ionic compounds have high melting
and boiling points and only exist as solids at
room temperature.
This is an image of a sodium chloride crystal. You can see that it
forms a cubic shape, just like the lattice shape above.
Although ionic compounds have high melting points they are lower
that that of a covalent network. Below is a comparison of sodium
chloride and silicon dioxide.
Structure
Melting point / °C
Boiling point / °C
Sodium chloride
crystal lattice
801
1413
Silicon dioxide
covalent network
1700
2230
Many ionic compounds will dissolve in water, they are said to be soluble. The reason for
their solubility is that, when added to water, the crystal lattice structure breaks down
and the ions are free to move in the solution.
Page 43
1e – Bonding and Structure
Valency and Ionic Compounds
For ionic compound the valency number is the number of electrons that must be lost or
gained in order to form a full outer shell. Group 4 and 0 atoms do not form ions.
Group
Example element
Outer electrons
1
2
3
Na
Mg
1
2
4
5
6
7
Al
P
S
Cl
3
5
6
7
Electrons gained or
lost
2
lost
Charge
2+
–
Valency
2
1
No
Ions
Page 44
1
gained
0
No
Ions
1e – Bonding and Structure
Summary – Ionic Compounds
Ionic
positive
oppositely
soluble
attraction
crystal
high
negative
atoms
ion
electrons
substances
are
made
____________________
up
forms
of
when
ions
not
____________________.
an
atom
either
gains
or
An
loses
____________________. During a chemical reaction this involves one atom donating its
electrons, forming a ____________________ ion, to another atom, forming a
____________________ ion.
An ionic bond is the ____________________ that the positive and negative ions have for
one another.
The structure of an ionic compound is known as a ____________________ lattice
structure
and
consists
of
a
well-ordered,
three-dimensional
lattice
of
____________________ charged ions.
When an ionic substance is melted or boiled, strong ionic bonds must be broken and, as
a result, these substances have ____________________ melting and boiling points. A
liquid (or molten) ionic compound has lost its rigid and well-ordered structure and the
ions are able to move over one another and flow.
Another way to break down the lattice structure is to dissolve the substance. If the
substance is ____________________ in water then the lattice structure is lost and the
ions are again able to flow in the solution.
Page 45
1e – Bonding and Structure
Electricity and Structure
An electrical current is defined as a flow of charged particles
For any substance to be an electrical conductor it must have charged particles, which
are free to move. Metals and carbon in the form of graphite have electrons, which are
free to move, and this is why metals and graphite are conductors of electricity.
Ionic compounds, as we have seen, are made of charged particles known as ions. Ionic
compounds can only conduct electricity when they are molten or in solution. Why is
this?
Ionic substances have a rigid crystal lattice structure,
however when they are melted or dissolved this rigid
structure is lost.
This image shows sodium chloride dissolving. When added
to water the ions are surrounded by water molecules and
are able to leave the lattice structure.
Ionic compounds can conduct electricity only as a liquid or when in solution because
they have ions that are free to move.
Solid ionic compounds contain ions, but these ions are rigidly held in place and are
unable to move
Covalent compounds, both molecular and network are made of atoms, which we learned
earlier must have a neutral charge, and not ions. Without charged particles it is not
possible for covalent compounds to conduct electricity.
Covalent compounds cannot conduct electricity, as they contain no ions.
 Check Test 1.12
Page 46
1f – Formulae and Reaction Quantities
Ionic Formulae
When working with an ionic compound it is important to be able to understand and write
an ionic formula. An ionic formula gives both the charge on the ions in the compound
and the ratio of the ions in the crystal lattice.
Writing an ionic formula links together two skills you already have:
1. To write a chemical formula using valency rules, group ions and Roman numerals
where appropriate
2. To work out the charge of an ion using the periodic table and electron
arrangements and write the symbol for the ion.
e.g. sodium chloride
1. Write the formula
1
1
Na Cl
Na1 Cl 1


Na Cl
2. Write the symbols for the ions
Na – 2,8,1 – loses 1 e – Na+
Cl – 2,8,7 – gains 1 e – ClTo write the ionic formula you simply write the ions in the ratio of the chemical
formula:
Na Cl is in the ration 1:1 which gives an ionic formula of
Na+ ClIf any ion is in the formula with a number other than one it must be put into brackets.
e.g.
Magnesium bromide is MgBr2 the ions are Mg2+ and Br- so the ionic formula is
Mg2+(Br-)2
Give the ionic formula of the following compounds:
a) magnesium nitrate
b) sodium oxide
c) aluminium oxide
d) iron (II) sulphide
 Check Test 1.13
Page 47
1f – Formulae and Reaction Quantities
Chemical Equation
Chemical reactions are often described in the form of a chemical equation. A chemical
equation is just like a word equation, except written with chemical formulae.
When writing a chemical formula it is important to first write a word equation, then
simply replace each chemical name with its chemical formula.
Note that if any of the seven diatomic elements are present as elements, their formula
reflects the fact that they are diatomic. e.g. hydrogen is written as H2.
For example:
magnesium +
Mg
soidum hydroxide
NaOH
+
bromine

magnesium bromide
Br2

MgBr2
+
hydrochloric acid

+
HCl

sodium chloride
NaCl
+
water
+
H 2O
For each of the following word equations, write the formula equation:
a) zinc + copper chloride  zinc(II) chloride + copper
b) lithium sulphate + barium chloride  lithium chloride + barium sulphate
c) carbon + oxygen  carbon dioxide
d) iron (II) oxide + carbon monoxide  iron + carbon dioxide
e) hydrogen + chlorine  hydrogen chloride
Page 48
1f – Formulae and Reaction Quantities
Showing the State
When writing chemical formulae it can be useful to also show the state of matter in
which the chemical would be. This is done using state symbols after the chemical
formula.
State of Matter
State symbol
solid
(s)
liquid/molten
(l)
gas
(g)
in solution (aqueous)
(aq)
For example if solid sodium reacts with water to form hydrogen gas and a solution of
sodium hydroxide, we could describe this reaction as:
Na(s)
+
H2O(l)
NaOH(aq)

+
H2(g)
1. Complete the following examples to show their state symbol.
a) solid sulphur
b) liquid nitrogen
c) ice
d) magnesium chloride solution
2. Write a formula equation including state symbols.
a)
Copper(II) chloride solution reacts with sodium carbonate solution to form
sodium chloride solution and solid copper(II) carbonate.
b)
Copper(II) carbonate powder reacts with hydrochloric acid solution (HCl) to
form copper(II) chloride solution, water and bubbles of carbon dioxide.
 Check Test 1.14
Page 49
1f – Formulae and Reaction Quantities
Finding a Balance
As you have seen, a chemical reaction can be described through the use of both word
equations and formula equations. These two methods rely on the same idea, reactants
go into the reaction and products come out.
When you look at some formula equations you might notice that not everything is
balanced. For example let’s look at the reaction of hydrogen and oxygen to make
water:
hydrogen
+
oxygen
+
H2

O2
water

H2O
As you can see diatomic molecules of hydrogen and oxygen react together to form a
water molecule. We can show this as a diagram:
= an atom of hydrogen
= an atom of oxygen
H2
+
O2

H2O
There is an imbalance between the reactants and the products, as there are two
hydrogen atoms and two oxygen atoms on the reactants side, but only one oxygen atom
and two hydrogen atoms on the products side. This means that the equation needs to be
balanced.
In order to balance an equation you must follow some rules:
1. First count the number of atoms of each element on both sides of the equations,
this can be noted down in a small table.
H2
+
H
2
O2

H2O
H
2
O
2
Page 50
O
1
1f – Formulae and Reaction Quantities
2. By looking at your count, decide which elements are lacking and whether they are
lacking as reactants or products.
There is a shortage of oxygen on the products side.
3. Next you can add reactants or products to the equation in order to correct the
imbalance. Whenever you add something you must recount including the atoms
you have added.
Note that only complete ‘formula units’ can be added, you cannot just add an
atom of an element unless it is a reactant or product of the reaction.
H2
+
H
2
O2

H2O
+H2O
H
2
4
O
2
O
1
2
4. Now you can repeat step 2 and 3 until the count is equal on both sides.
There is a shortage of hydrogen on the reactants side.
H2
+H2
+
H
2
4
O2

H2O
+H2O
H
2
4
O
2
2
O
1
2
5. The count is now balanced, with four hydrogen atoms and two oxygen atoms on
either side.
You can now rewrite the equation totalling up the number of each reactant and
product, as with formulae the number 1 need not be written.
2H2
+
O2

This is a balanced equation.
Page 51
2H2O
1f – Formulae and Reaction Quantities
Balance the following equations:
1.
H2
+
Cl2

HCl
2.
CH4
+
O2

CO2
3.
CO
+
O2

CO2
4.
Al(OH)3
+
HCl

5.
Fe
O2

Fe2O3
+
+
H2O
AlCl3 +
 Check Test 1.15
Page 52
H2O
1f – Formulae and Reaction Quantities
Gram Formula Mass
Earlier in this topic you saw that every atom of an element has mass, due to protons and
neutrons in the nucleus, and that an average of all possible isotopes is called the
Relative Atomic Mass or R.A.M.
Using these masses and the chemical formula it is possible to calculate the Gram
Formula Mass or GFM.
Due to the incredibly small size of atoms it is not logical to try and measure the mass of
one atom or even a few million atoms. Instead we measure the mass of one mole of
atoms. The term mole simply represents a number, an extremely large number and by
measuring this quantity of atoms and molecules we can measure sensible and easy to
handle masses.
The GFM is defined as the mass of one mole of a substance.
To calculate the GFM you must start with the correct chemical formula.
1. First, write down the chemical formula of the relevant substance
e.g. H2O
2. Using the data booklet, note the RAM of each element in the substance
e.g. H = 1
O = 16
3. Add up the total mass of the substance taking into account the chemical formula
e.g. H2O :
2xH=2x1=2
1 x O = 1 x 16 = 16
Total = 18
The GFM of water is 18 which means that one mole of water has a mass of 18g.
Note that if the formula has brackets then the number out of the brackets multiplies
everything inside of the brackets. e.g. Mg(NO3)2 is 1 x Mg, 2 x N and 6 x O.
Calculate the GFM of the following substances:
a) NaCl
b) O2
c) AgCl
d) Al2(SO4)3
 Check Test 1.16
Page 53
1f – Formulae and Reaction Quantities
Mass and Moles
Knowing the formula of a substance, you can now calculate its GFM.
This forms a relationship with the mass of a substance (m) and the number of moles (n)
of a substance as shown on page 1 of the data booklet
n
m
GFM
number of moles = mass of a substance / GFM
This can be rearranged to give
m  n x GFM
e.g.
If you were asked to weigh out 2 moles of sodium chloride, what mass would you
collect?
m
e.g.
mass of a substance = number of moles / GFM
= n x GFM
= 2 x 58.5
= 117 g
GFM
= NaCl
= 23 + 35.5
= 58.5
If you measured out 50 g of calcium carbonate, how many moles of calcium
carbonate have you collected?
n
= m / GFM
= 50 / 100
= 0.5 mol
GFM
= CaCO3
= 40 + 12 + (3 x 16)
= 100
Calculate the mass of:
a) 0.1 mol of O2
b) 5 mol of NaBr
c) 3.5 mol of CuCl2
d) 0.5 mol of K2SO4
Page 54
1f – Formulae and Reaction Quantities
Calculate the number of moles in:
a) 10 g of H2
b) 12.8 g of SO2
c) 20 g of He
d) 12.7 g of I2
 Check Test 1.17
Page 55
1f – Formulae and Reaction Quantities
Concentration
Concentration (C) is a measure of the quantity of chemical present in a given volume.
Chemical concentration is normally given as the number of moles in a litre of a solution,
mol l-1 (moles per litre).
This forms a relationship with the volume of a solution (V) and the number of moles (n)
of a substance as shown on page one of the data booklet
nCV
number of moles = concentration x volume
This can be rearranged to give
C
n
V
concentration = number of moles / volume
n
C
volume = number of moles / number of moles
and
V
The volume must be in litres before being used in the calculation.
To change a volume in cm3 into litres we must divide by 1000.
e.g 100 cm3 = 100 ÷ 1000 = 0.1 litres
Convert the following volumes into litres:
a) 75 cm3 =
b) 150 cm3 =
c) 10 cm3 =
d) 700 cm3 =
Page 56
1f – Formulae and Reaction Quantities
e.g.
If you were asked to make a 500 cm3 solution containing 2 moles of sodium
chloride, what concentration would this solution have?
c
e.g.
500 cm3
=n/v
= 2 / 0.5
= 4 mol l-1
= 500 / 1000
= 0.5 litres
If you had 750cm3 of a 1.5 mol l-1 solution, how many moles of the chemical
would you have?
n
750 cm3
=cxv
= 0.75 x 1.5
= 1.125 mol
= 750 / 1000
= 0.75 litres
Calculate the concentration of a 800 cm3 solution containing:
a) 1.5 moles of MgCl2
b) 5 mol of NaBr
c) 3.5 mol of CuCl2
d) 0.5 mol of K2SO4
Calculate the number of moles in a 200 cm3 solution with a concentration of:
e) 2 mol l-1
f) 1.4 mol l-1
g) 0.01 mol l-1
h) 0.05 mol l-1
Page 57
1f – Formulae and Reaction Quantities
Concentration and Mass
If we look at both of our calculations side by side we can see that they have one variable
in common, the number of moles (n).
n
m
GFM
nCV
With this common link, we can now know more about a solution by using both
calculations together.
e.g.
e.g.
If you were asked to make a 500 cm3 solution of sodium chloride with a
concentration of 1.4 mol l-1, what mass of sodium chloride would you need?
n
=cxv
= 1.4 x 0.5
= 0.7 mol
500 cm3
= 500 / 1000
= 0.5 litres
m
= n x GFM
= 0.7 x 58.5
= 40.95 g
GFM
= NaCl
= 23 + 35.5
= 58.5
If you were to make 750cm3 of a potassium sulphate (K2SO4) solution using 17.4 g
of solid potassium sulphate, what concentration would the solution have?
n
= m / GFM
= 17.4 / 174
= 0.1 mol
GFM
= K2SO4
= 78 + 32 +64
= 174
c
=n/v
= 0.1 x 0.75
= 0.075 mol l-1
750 cm3
= 750 / 1000
= 0.75 litres
Page 58
1f – Formulae and Reaction Quantities
Calculate the mass of magnesium chloride (MgCl2) required to make a 1.5 litre solution
with a concentration of:
a) 2 mol l-1
b) 1.4 mol l-1
c) 0.01 mol l-1
d) 0.05 mol l-1
Calculate concentration of a 350 cm3 solution containing:
a) 20.6 g of sodium bromide (NaBr)
b) 4 g of ammonium nitrate (NH4NO3)
c) 6.38 g of copper(II) sulphate
d) 21.2 g of sodium carbonate
 Check Test 1.18
Page 59