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Chemistry
The Periodic Table cont’d
Presented by
A.A.S Lateef
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The periodic table
• Learning objectives, at the end of this lecture,
 Students should have understood general properties of the
various groups in the periodic table
 Students should be able to explain the justification behind
the position of hydrogen in the periodic table.
 Students should be explain the trends in reactivity among
the elements across the periods and down the group and
explain the reason for the variations.
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The Periodic Table
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Review of key points about periodic
table
• The Periodic Table is a way of classifying the elements.
• It shows them in order of their proton number. Lithium has
3 protons, beryllium has 4, boron has 5, and so on.(The
proton number is the lower number beside each symbol.)
• When arranged by proton number, the elements show
periodicity: elements with similar properties appear at
regular intervals. The similar elements are arranged in
columns.
• Look at the columns numbered 1 to 0. The elements in these
form families called groups.
• The rows are called periods. They are numbered 0 to 7.
• The heavy zig-zag line above separates metals from nonmetals, with the non-metals to the right (except for
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hydrogen).
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More about groups
 The group number is the same as the number of outer-shell electrons in the
atoms, except for Group 0. In Group I the atoms have one outershell
electron, in Group II they have two, and so on.
 The outer-shell electrons are also called valence electrons. And they
are very important: they dictate how an element behaves.
 So all the elements in a group have similar reactions, because they
have the same number of valence electrons.
 The atoms of the Group 0 elements have a very stable arrangement of
electrons in their outer shells. This makes them unreactive.
More about the periods
The period number tell us the number of electron shells in the atoms. So in the
elements of Period 2, the atoms have two electron shells. In Period 3 they have
three, and so on.
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Hydrogen
• Hydrogen sits alone. That is because it has one outer
electron, and forms a positive ion (H+) like the Group I
metals – but unlike them it is a gas, and usually reacts
like a non-metal.
Patterns and trends in the Periodic Table
• As you saw, the elements in a group behave in a similar
way. But they also show trends. For example as you go
down Group I, the elements become more reactive.
Down Group VII, they become less reactive.
• Across a period there is another trend: a change from
metal to non-metal. For example in Period 2, only
sodium, magnesium, and aluminium are metals. The
rest are non-metals. So if you know where an element
is, in the Periodic Table, you can use the patterns and
trends to predict how it will behave.
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Group 1: alkali metals
 The alkali metals are in Group I in the Periodic Table: lithium, sodium,
potassium, rubidium, caesium and francium. Only the first three of these
are safe to keep in the school lab. The rest are violently reactive.
 They react violently with water, giving hydrogen and a hydroxide.
 When heated and plunge them into gas jars of chlorine, they burst into
flame. They burn brightly, forming chlorides.
 The metals also burst into flame when you heat them and plunge
them into gas jars of oxygen. They burn fiercely to form oxides. These
dissolve in water to give alkaline solutions.
NB: They react in a similar way but show some trends in it. This is because
they have the same number of valence electrons.
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Why are they so reactive?
• The alkali metals are the most reactive of all the metals.
• Why? Because they need to lose only one electron, to gain a stable outer
shell. So they have a strong drive to react with other elements and
compounds, in order to give up this electron.
Why does reactivity increase down Group I?
• In reactions, the Group I atoms lose their outer electron, to gain a stable
outer shell.
• The more shells there are, the further the outer electron is from the positive
nucleus – so the easier to lose.
• And the easier it is to lose an electron, the more reactive the metal will be!
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The halogens-group vii
• A non-metal group
• Group VII is a group of non-metal elements. It includes fluorine, chlorine,
bromine, and iodine. These are usually called the halogens.
• They all:
 form coloured gases. Fluorine is a pale yellow gas and chlorine is a green
gas. Bromine forms a red vapour, and iodine a purple vapour
 are poisonous.
 form diatomic molecules (containing two atoms). For example, Cl2.
Why do they react in a similar way?
• The halogens react in a similar way because their atoms all have 7 valence
(outer-shell) electrons. Compare the fluorine and chlorine atoms:
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Why are they so reactive?
The halogen atoms need just one more electron to reach a stable
outer shell of 8 electrons. So they have a strong drive to react
with other elements or compounds, to gain this electron. That is
why they are so reactive.
When halogen atoms react with metal atoms they accept
electrons, forming halide ions. So the products are ionic.
But with non-metal atoms such as hydrogen and carbon, they
share electrons, forming molecules with covalent bonds.
Why does reactivity decrease down Group VII?
Halogen atoms react to gain or share an electron. The positive nucleus of the
atom attracts the extra electron.
The more shells there are, the further the outer shell is from the nucleus. So
attracting an electron becomes more difficult. So reactivity falls.
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The noble gases-Group 0 (or VIII)
• This group of non-metals contains the elements helium, neon, argon,
krypton and xenon.
These elements are all:
 non-metals
 colourless gases, which occur naturally in air
 monatomic – they exist as single atoms
 unreactive. This is their most striking property. They do not normally react
with anything. That is why they are called noble.
Why are they unreactive?
• Atoms react in order to gain a stable outer shell of electrons. But the atoms
of the noble gases already have a stable outer shell – with 8 electrons,
except for helium which has 2 (since it has only one shell):
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Across the Period
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Trends across the period
•
•
•
•
•
•
The number of valency (outer-shell) electrons increases by 1 each
time.
It is the same as the group number, for Groups I to VII.
The elements go from metal to non-metal. Silicon is in between. It
is like a metal in some ways and a non-metal in others. It is called a
metalloid.
Melting and boiling points rise to the middle of the period, then fall
to very low values on the right. (Only chlorine and argon are gases
at room temperature.)
The oxides of the metals are basic – they react with acids to form
salts.
Those of the non-metals are acidic – they react with alkalis to form
salts. But aluminium oxide is in between: it reacts with both acids
and alkalis to form salts. So it is called an amphoteric oxide.
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Valency
• Look at the last two rows in the table. One shows a typical
compound of each element. The other shows the valency of the
element in that compound.
• The valency of an element is the number of electrons its atoms
lose, gain or share, to form a compound.
• Sodium always loses 1 electron to form a compound. So it has a
valency of 1. Chlorine shares or gains 1, so it also has a valency of
1. Valency rises to 4 in the middle of the period, then falls again. It
is zero for the noble gases.
• Note that valency is not the same as the number of valency electrons.
But:
• the valency does match the number of valency electrons, up to
Group IV
• the valency matches the charge on the ion, where an element forms
ions.
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Reactivity across periods
• reactivity decreases across the metals. Aluminium is a lot less
reactive than sodium, for example. Why? Because the more
electrons a metal atom needs to lose, the more difficult it is.
(The electrons must have enough energy to overcome the pull
of the nucleus.)
• reactivity increases across the non-metals (apart from Group
0). So chlorine is more reactive than sulfur. Why? Because the
fewer electrons a non-metal atom needs to gain, the easier it is
to attract them.
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In summary,
Reactivity increases down the group in metals
(Potassium is more reactive than sodium) but
decreases across the periods in metals (sodium is
more reactive than magnesium) etc
It decreases down the group in non-metals (Chlorine
is more reactive than Iodine) but increases across the
periods in non-metals. (Chlorine is more reactive than
silicon) etc.
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Exercises
JAMB 1994, Q12
JAMB 1989, Q12
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Exercises
JAMB 1994, Q12
Option A
JAMB 1989, Q12
Option
C
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JAMB 1999, Q15
JAMB 1997, Q9
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JAMB 1999, Q15
Option
C
JAMB 1997, Q9
Option
C
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