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Fundamental Chemistry: Theory and Practice
Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
Carnegie College
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DH2K 34
Fundamental Chemistry: Theory and Practice
Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
Carnegie College
Contents
Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Equations
1
Covalent Bonding
1
Ionic Bonding
3
Nomenclature
5
–IDE and –ATE
5
-ATE and –ITE
5
Metallic Bonding
5
Covalent Bonding: A Second Look
6
Discrete Covalent Molecules
7
Covalent Network Molecules
7
Chemical Formula
8
Balance
12
The Bonding and Formula Progress Checklist
18
Answers to SAQs
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Fundamental Chemistry: Theory and Practice
Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
Carnegie College
Topic 3 – Bonding, Chemical Symbolism and
Balanced Chemical Equations
By the end of this section you should be able to:

Outline the three main types of bonding

Understand and use nomenclature accurately

Use chemical formula accurately

Balance chemical equations
There are three main types of bonding:



Covalent
Ionic
Metallic
Covalent Bonding
Atoms form bonds in order to increase stability; remember that an element is most
stable when it has a full outer electron shell, that is when it has the same outer electron
configuration as the noble gases.
i)
Covalent Bonding in the elements occurs in the non-metals. For example
hydrogen has one outer electron; to gain a full outer shell like the noble gas helium
it needs one more electron. Hydrogen is able to achieve this stability by sharing its
one outer electron with another atom of hydrogen. We call this shared pair of
electrons a single covalent bond. Covalent bonding also occurs in other elements
forming diatomic molecules, such as Oxygen (O2) and Nitrogen (N2).
In covalent bonding there is a sharing of electrons to give stability to the outer shell of
each atom involved. Oxygen has the electron arrangement 2)6. Four of the outer
electrons are paired up, leaving two others. Two oxygen atoms can share these two
outer electrons to give a stable covalent bond. Two bonds in fact. Nitrogen has the
electron configuration 2)5, so three electrons are left unpaired in the outer shell. A
nitrogen atom can therefore bond with another nitrogen atom to from N2, which has
three covalent bonds.
The group 7 elements (the halogens) also form diatomic molecules; how many pairs of
electrons are shared in the halogens?
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Fundamental Chemistry: Theory and Practice
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ii)
Carnegie College
Covalent Bonding in Molecular compounds occurs when more than one type
of non-metal element join together to increase stability. You will recognise some
common covalently bonded compounds.
Water (H2O), carbon dioxide (CO2) and methane (CH4) the gas that heats many
of our homes. In all of these examples electrons are shared giving a stable
electron arrangement.
For water we say that the molecular formula is H2O but the structural formula is
drawn as
H
H
O
For Carbon dioxide CO2 is the molecular formula while the structural formula is shown
O=C=O and for methane CH4 and
H
H
C
H
H
All discrete covalent molecules have low melting and boiling points and tend to be
liquids and gases at room temperature. Between all molecules in the liquid or solid
state weak forces called van der waals’ forces exist these forces become larger as the
size of the molecule increases, it is these forces that enable us to liquefy and solidify
gases. We can also say that since the electrons are held tightly in a bond and
because covalent compounds are electrically neutral they do not conduct electricity;
this is because there are no free electrons to carry the charge.
When molecules are formed between atoms of different elements polar covalent
bonding occurs. In polar covalent bonding the electrons are not shared equally
between the different elements. That is to say some elements have a greater affinity
for electrons than others. The atom with the greater electron attracting power acquires
a very slight negative charge shown δ- while those with less attracting power acquire a
slight positive charge δ+. Let’s consider water
This bonding allows
δ+
δ+
weak bonds to be formed within the structure
of water and gives
H
H
water higher melting and boiling points than
would be expected
Oδin a molecule of this size, these bonds are
called hydrogen
bonds and give water a lattice structure. This ordering of the water molecules means
that when water freezes ice is less dense than water and floats on the top, allowing, as
we know aquatic species to survive during the winter.
Hδ+
Hδ+
Hδ+
Oδ-
Hδ+
OδHδ+
Hydrogen Bond
Hδ+
Oδ-
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iii) Covalent Networks are unusual for molecules joined by covalent bonding in that
they have high melting and boiling points. A good example is diamond. A diamond is
many millions of carbon atoms linked together by single covalent bonds in a
tetrahedral arrangement. Diamonds are strong, rigid and solid at room temperature.
Graphite is another form of carbon that exists as a covalent network; in this case the
structure is made of layered rings linked together by van der waals’ forces. In graphite
some of the electrons are free to move about the structure, we say these electrons are
‘delocalised’ and are capable of conducting electricity. Carbon in the form of
graphite is unusual as it is the only non-metal that can conduct electricity.
Ionic Bonding
Ionic bonding occurs when one atom has a much larger ‘pull’ on electrons than
another. Alkali metals and Halogens form ionic bonds.
In ionic bonding one atom, the one with the greater electronegativity removes the
electron from another atom with a smaller electronegativity. Both end up with noble
gas electron arrangement in the outer shell. A chlorine atom with an electron
configuration 2)8)7 will try to pick up an electron to give an electron configuration 2)8)8
and form the chloride ion Cl-. A Sodium atom with an electron configuration 2)8)1 will
try to loss an electron to give an electron configuration 2)8) and form the sodium ion
Na+. In general metals lose electrons to achieve a full outer shell and form positive
ions, while non-metals gain electrons to achieve a full outer shell and form negative
ions.
Now try SAQ 1 on the following page
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1
What will be the size and type (positive or negative) formed by each of the elements
below?
Lithium
Calcium
Aluminium
Nitrogen
Oxygen
Fluorine
Check your answer at the end of this booklet.
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The overall charge on an ionic compound must be zero you will learn more about this
in a short while. All ionic compounds are solid at room temperature and do not
conduct electricity when solid, however when ionic compounds are dissolved (and
most are soluble in water) or when they are melted ionic compounds are able to
conduct electricity due to the movement of ions.
Nomenclature
Many chemicals have traditional names, but all chemicals have a scientific name.
Acetic acid is the traditional name of the acid in vinegar; ethanoic acid is the scientific
name. These names can tell you a lot about a compound. They can tell you both the
number and type of atoms that a molecule contains. Many compounds contain more
than one type of atom.
–IDE
If a chemical has a name that ends in –IDE, this indicates that it is made up of two
different types of atoms. So sodium chloride contains two types of atoms: Sodium and
Chlorine. The metal sodium retains its name and the chlorine is renamed chloride. (A
notable exception to this rule is Hydroxides(OH-), sodium hydroxide contains three
elements – sodium, oxygen and hydrogen.
The number of each atom can be denoted by the prefixes:




Mono
di
tri
tetra
one
two
three
four
So carbon dioxide contains two oxygen; its formula is CO2.
-ATE and –ITE
Names that end in –ate and –ite contain three different types of atoms, one of which is
oxygen. Compounds with names ending in ate have a greater proportion of oxygen
than those with names ending in ite.
Metallic Bonding
Metals make up the left and ‘middle’ section of the periodic table. Clearly identify the
metals on your periodic table.
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In metals the outer electrons are delocalised (delocalised means free to move). This
produces a regular matrix of positively charged ions in a sea of electrons. This
delocalisation is why metals are good at conducting electricity and heat. Unlike ionic
compounds metals conduct due to movement of electrons rather than movement of
ions. In addition since all the positive ions are of equal size, metals are malleable and
ductile. This means that metals can be hammered into different shapes and drawn
into wires.
+
+
+
+
+
+
+
+
+
+
+
+
+
+
Properties of metals:
Electrical conductivity
Thermal conductivity
Ductility
Malleability
they can conduct electricity
they can conduct heat
they can be drawn into wire
they can be beaten and rolled into sheets
Most metals except Mercury and the alkali metals have high melting and boiling points.
Covalent Bonding: A Second Look
Covalent bonding can be described at the ________ ____________ of electrons.
There are two forms of covalent bonding:

Covalent Discrete Molecules

Covalent Network Molecule
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Discrete Covalent Molecules
This type of molecule can be divided into two types; pure covalent bonding and polar
covalent bonding. Pure covalent bonding occurs between atoms of the same element,
for example Oxygen (O2) or Sulphur (S8). In these cases the electrons are shared
equally between the atoms. Where more than one type of atom is involved then
electrons are unequally shared; polar covalent bonding – this can lead to areas of
positive and negative charge within the molecule. These charged areas are much less
than the charge on a proton or an electron and are shown by the Greek symbol delta
+ or -. Water is a good example.
OδHδ+
Hδ+
Discrete covalently bonded compounds usually have low melting and boiling points.
Covalent Network Molecules
These are covalently bonded structures such as Carbon (diamond) or Silicon Dioxde
(Quartz and Sand).
These have four bonds and form very strong stable network structures such as
diamond, graphite and sand.
It takes a lot of energy to melt these compounds, due to the strong nature of the
bonds.
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Bonding Summary
Compound
Ionic: Solid does not conduct
electricity. Conducts electricity when
molten or dissolved
High Melting Points
Covalent: Does not conduct electricity
Covalent Network
High Melting Point
Solids
Covalent Discrete
Low Melting Point
Solids Low boiling
point Gasses
Question
Name the only covalently bonded element that can conduct electricity.
Chemical Formula
Many atoms can be found as ions. An ion is an atom that has gained or lost an
electron. Atoms gain or lose electrons to get full outer shells and increase their
stability.
The alkali metals produce a charge of one positive and would be written as X+. So
Sodium would be written as Na+. While Magnesium, which looses two electrons when
it forms ions, can be written as Mg2+. The halogens gain an electron and have a one
negative charge. So Chlorine would be written as Cl-. While oxygen in group 6 would
gain two electrons and be written as O2-.
For a compound to be stable, the electrical changes must balance out. That is to say
that the overall charge on the compound will be zero. So when a negative ion joins
with a positive ion the amount of the positive charge must equal that of the negative
charge. This is fine when the charges are 1+ and 1-. Then we just need one of each
atom.
However, if we have an atom with a charge of 2-, and another with a charge of 1+, we
would need two of the 1+ atoms to balance out the 2- of the other.
Let’s look at Oxygen. This has a charge of 2- and so is written as O2- so if this reacts
with Sodium which is 1+, how many Sodium do we need?
That’s right two.
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So the chemical formula would be:
O2-(1) Na+(2). This could be written as ONa2. But due to convention the positive ion,
ie the hydrogen or metal is named first and it would be written as Na2O.
If you know the charge on each of the ions you can easily work out the chemical
formula. The way to do this is to cross multiply. So if the negative ions have a charge
of 3- and the positive ion a charge of 2+, you simply multiply the positive ion with the 3
and the negative ion with 2, ie
X3- Y2+ the X is multiplied by the 2 giving X2 and the Y is multiplied by 3 to give Y3.
This would give six positive and negative changes and the formula would be X2Y3. We
can do this as a series of steps until we become more confident.
For simple negative ions, that is those that are found on the periodic table the ending
in the name is –ide, so a chlorine atom becomes a chloride ion and an oxygen atom
becomes an oxide ion.
Lets us try that with Potassium Oxide.
Step 1: Write down the names of the ions involved.
Potassium
Oxide
Step 2: Write down the symbols
K+
O2-
Step 3: Write down the charges
K
1
O
2
Step 4: Cross multiply the charges
K
1
2
O
2
1
Step 5: Write down the formula of the compound
K2O1
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Step 6: Simplify if possible
K2O
Try the examples below
2
a)
Sodium chloride
b)
Magnesium oxide
c)
Calcium chloride
d)
Aluminium oxide
Check your results with those given at the end of the booklet.
Sometimes an ion is made up of a group of non-metals that carry an overall charge;
these are called compound ions and can be found on page 4 of your data book.
Examples include the hydroxide ion OH-, sulphate ion SO42- and the ammonium ion
NH4+. In this case it is the size of the charge that is swapped in order to get the
formula. Follow the steps outlined in the simple example below to get the answer.
For example Sodium carbonate
Sodium
Na+
1
2
carbonate
CO322
1
Na2 (CO3)1
The brackets are important in many compounds involving compound ions as it tells us
that the number outside the bracket refers to the whole ion and not just part of it. If the
number is 1 then it is appropriate at this stage to remove both the number and the
brackets so our answer becomes:
Na2CO3
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Try the examples below
3
a)
Potassium Dichromate
b)
Ammonium Carbonate
c)
Sodium Nitrate
d)
Magnesium phosphate
Check your results with those given at the end of the booklet.
In addition to the above it is important to note that the transition elements are capable
of forming ions of more than one size. For example copper forms Cu+ ions and Cu2+
ions and iron forms Fe2+ and Fe3+ ions. In examples like these the size of the charge
is always indicated using roman numerals, so a compound containing the Fe2+ ion
would for example be shown as Iron (II) chloride while a compound containing Fe3+
would be Iron (III) chloride.
Try the examples below.
4
a)
Lead(II) nitrate
b)
Iron(III)sulphate
c)
Copper(I) chloride
d)
Silver(I) nitrate (In fact unless told otherwise it is safe to assume that silver will
always form 1+ ions).
Check your results with those given at the end of the booklet.
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Balance
When chemicals come together and react, the starting material and the product often
appear completely different.
Example 1: Think of common table salt. This is made of the reactive alkali metal
sodium and the poisonous gas chlorine, neither of which you would want to eat; yet
when these two elements react, a food seasoning is produced.
The combining of sodium and chlorine can be written as a word equation where the
reactants are shown on the left and the products on the right of the arrow.
Reactants
Products
Sodium + chlorine
Sodium chloride
This is called a word equation
We can replace the words with chemical formula, which can be worked out using the
rules of valency and by knowing which elements form diatomic molecules.
The equation above becomes
Na + Cl2
NaCl
When we count up the numbers of each element on each side of the arrow they should
be equal.
Element
Sodium
Chlorine
Number of LHS
1
2
Number on RHS
1
1
As you can see the number of each type of element is not balanced on both sides of
the equation. We have to add more molecules of some compounds to make the
equation balance. In chemistry it is very important that any equations we write should
balance.
So how can we balance this equation?
We have 2 chlorines on the LHS but only one on the RHS; we need to add a 2 to the
LHS.
Na + Cl2
2NaCl
You will note that this 2 also affects the number of sodium on the RHS. Whenever a
number is added in front of a compound it affects all the elements in that compound,
so now we have 2 sodium on the RHS and only one on the LHS, to make the equation
balance a 2 will need to be placed in front of the sodium on the LHS.
2Na + Cl2
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2NaCl
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Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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The equation is now balanced.
Element
Sodium
Chlorine
Number of LHS
2
2
Number on RHS
2
2
Example 2: A Human Example
Think of the food you eat. You take in sugars and you breathe out carbon dioxide and
water.
This can be written as a word equation:
Glucose + Oxygen
Carbon Dioxide + Water
We could write this in its chemical form:
C6H12O6 + O2
CO2
+ H2O
A question of balance
When we count up the numbers of each element on each side of the arrow they should
be equal.
Element
Carbon
Hydrogen
Oxygen
Number of LHS
6
12
8
Number on RHS
1
2
3
As you can see the number of each type of element is not balanced on both sides of
the equation. We have to add more molecules of some compounds to make the
equation balance.
So how can we balance this equation?
C6H12O6 + O2
CO2
+ H2O
We need 12 Hydrogen and six Carbons on the RHS, so we need six H2O and six CO2
so giving:
C6H12O6 + O2
6CO2
+ 6H2O
We now have 18 oxygen on the RHS and only eight on the LHS. So to make it
balance we need to add six O2 to the LHS.
C6H12O6 + 6O2
© Carnegie College
6CO2
+ 6H2O
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Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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This equation now balances out:
Element
Carbon
Hydrogen
Oxygen
Number of LHS
6
12
18
Number on RHS
6
12
18
Try the examples in SAQ 5 on the following page.
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Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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5
1)
When methane (CH4) burns in oxygen (O2), carbon dioxide (CO2) and water
(H2O) are produced. Write a word equation, a formula equation and a balanced
chemical equation for this reaction.
2)
Hydrogen and oxygen combine in an explosive reaction to form water, write word,
formula and balanced chemical equations for this reaction.
3)
Balance each of the following formula equations
a
Mg + HCl
MgCl2 + H2
b
NaOH + H2SO4
Na2SO4 + H2O
c
AgNO3 + Cu
Ag + Cu(NO3)2
d
C2H4 + O2
CO2 + H2O
SAQ 5 continues over the page
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SAQ 5 continued
e
C6H12O6
C2H5OH + CO2
f
SnO2 + HCl
SnCl4 + H2O
Check your results with those given at the end of the booklet.
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Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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Additional information can be added to balanced chemical equations by the use of
state symbols.
s- solid
l- liquid
g- gas
aq- aqueous( dissolved in water)
So our first example would become
2Na(s) + Cl2(g)
© Carnegie College
2NaCl(s)
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Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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The Bonding and Formula Progress Checklist
Tick the boxes only if you can:
Topic
Outline the three main types of bonding
Understand and use nomenclature accurately
Use chemical formula accurately
Balance chemical equations
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Understand(?)
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Fundamental Chemistry: Theory and Practice
Topic 3 – Bonding, Chemical Symbolism and Balanced Chemical Reactions
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Answers to SAQs
Answer to SAQ 1
1+
2+
3+
321-
Answers to SAQ 2
Sodium chloride
Magnesium Oxide
Calcium chloride
Aluminium oxide
NaCl
MgO
CaCl2
Al2O3
Answer to SAQ 3
Potassium dichromate
Ammonium carbonate
Sodium Nitrate
Magnesium phosphate
K2Cr2O7
(NH4)2CO3
NaNO3
Mg3(PO4)2
Answer to SAQ 4
Lead (II) nitrate
Iron(IIl) sulphate
Copper(I)chloride
Silver(I)nitrate
Pb(NO3)2
Fe2(SO4)3
CuCl
AgNO3
Answer to SAQ 5
Methane + Oxygen
carbon dioxide + water
CH4 + O2
CO2 + H2O
CH4 + 2O2
CO2 + 2H2O
Hydrogen + Oxygen
© Carnegie College
Water
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H2 + O2
2H2 + O2
H2O
2H2O
Mg + 2HCl
MgCl2 + H2
2NaOH + H2SO4
Na2SO4 + 2H2O
2AgNO3 + Cu
2Ag + Cu(NO3)2
C2H4 + 3O2
2CO2 + 2H2O
C6H12O6
2C2H5OH + 2CO2
SnO2 + 4HCl
SnCl4 + 2H2O
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