Danyal Education (Contact: 9855 9224)
“A commitment to teach and nurture”
Periodicity: The Periodic Table
Candidates should be able to:
on
I)
Periodic Trends:
a) describe the Periodic Table as an arrangement of the elements in the order of increasing proton (atomic)
number
b) describe how the position of an element in the Periodic Table is related to proton number and electronic
structure
al
Ed
uc
ati
c) describe the relationship between group number and the ionic charge of an element (*) (#)
d) explain the similarities between the elements in the same group of the Periodic Table in terms of their
electronic structure
e) describe the change from metallic to non-metallic character from left to right across a period of the Period
Table
f) describe the relationship between group number, number of valency electrons and metallic/non-metallic
character
g) predict the properties of elements in Group I and VII using the Periodic Table
II)
Group Properties:
a) describe lithium, sodium and potassium in Group I (the alkali metals) as a collection of relatively soft, low
density metals showing a trend in melting point and in their reaction with water
ny
b) describe chlorine, bromine and iodine in Group VII (the halogens) as a collection of diatomic non-metals
showing a trend in color, state and their displacement reactions with solutions of other halide ions
c) describe the elements in Group 0 (the noble gases) as a collection of monatomic elements that are chemically
un-reactive and hence important in providing an inert atmosphere, e.g. argon and neon in light bulbs; helium
in balloons; argon in the manufacture of steel (*) (#)
Da
d) describe the lack of reactivity of the noble gases in terms of their electronic structures
III)
Transition Elements: (*) (#)
a) describe the transition elements as metals having high melting points, high density, variable oxidation state
and forming coloured compounds
b) state that the elements and/or their compounds are often able to act as catalysts (see also 6.1(d)).
* not in combined Science syllabus
# not in N level Science syllabus
O Level Chemistry – The Periodic Table
1
Danyal Education (Contact: 9855 9224)
“A commitment to teach and nurture”
The Periodic Table is an arrangement of elements in
the order of increasing proton number.
The elements are classified generally into metals and
non-metals.
Features of the Periodic Table
The elements in the same group of the Periodic Table have
the same number of valence electrons, hence similar
electronic configuration.
Example:
Elements in Group 1 have 1 valence electron.
Lithium: 2,1
Sodium: 2,8,1 Potassium: 2,8,8,1
Elements in Group VII have 7 valence electrons:
Fluorine: 2,7 Chlorine: 2,8,7 Bromine: 2,8,18,7
al
Ed
uc
ati
The Periodic Table divides the elements into periods
and groups.
A group is a vertical column of elements. A period is
a horizontal row of elements.
Explain the similarities between the elements in the same
Group of the Periodic Table in terms of their electronic
structure.
on
What is the Periodic Table?
The Periodic Table has eight groups of elements,
numbered from I to 0. Group 0 is also sometimes
called Group VIII.
Elements in the same group have the same number
of valence electrons, and have similar chemical
properties.
The Periodic Table has seven periods of elements,
numbered 1 to 7. The proton number increases across
a period, from left to right.
Elements in the same period have the same number
of electronic shells.
The block of metals between Group II and III are
known as the transition metals. They can exist in
more than one oxidation state. Eg Fe can form ions
Fe2+ and Fe3+.
ny
Describe the relationship between group number and
ionic charge of an element.
Da
Elements in the same group have the same number
of valence electrons.
Hence, elements in the same group will have the
same valency and same ionic charge.
Examples:
Potassium and Sodium in Group I will form ions of
charge +1.
Fluorine and chlorine in Group VII will form ions of
charge -1.
Describe the properties of transition metals.
Transition metals have high melting points, high
density, variable oxidation state and form
colored compounds.
The elements and/or their compounds are often
able to act as catalysts.
As a result of these similar electronic configurations, the
elements in the same group will have similar chemical
properties.
Describe the change from metallic to non-metallic
character from left to right across a period of the Periodic
Table.
Metals are grouped on the left-hand side of each period.
Non-metals are grouped on the right-hand side.
From left to right across a period, there is a decrease in
metallic properties and an increase in non-metallic
properties.
Describe the relationship between group number, number
of valence electrons and metallic/non-metallic character.
Elements in the same group have the same number of
valence electrons.
Eg Group I elements have 1 valence electron. Group VII
elements have 7 valence electrons.
Group I, II, III elements form ions with charge same as
their group number. All the elements in these groups are
metals (except Boron at top of Group III)
Group IV and V elements have a maximum oxidation state
that is the same as the group number of the element.
Eg. Carbon in Group IV has max oxidation state +4
The changes from non-metallic to metallic character when
going down these groups are obvious.
Group VI and VII elements form ions of charge -2 and -1
respectively. All the elements in these groups are nonmetals (except Polonium at bottom of Group VI)
Group 0 elements do not form ions and are inert gases.
They exist as mono-atomic elements.
All the elements in this group are non-metals.
O Level Chemistry – The Periodic Table
2
Danyal Education (Contact: 9855 9224)
“A commitment to teach and nurture”
Group Properties
Group Trends
Group I elements (also known as “alkali metals”)
Group I trends:
Going down the group,
 Melting and boiling point decreases
 Density increases
 Reactivity increases
al
Ed
uc
ati
Chemical Properties:
 Reactive metals
They are stored in oil to prevent them from reacting
with air and water.
 React with cold water to form alkali and hydrogen
 Strong reducing agents
They lose their one valence electron readily, so they
will behave as strong reducing agents.
Group VII trends:
Going down the group,
 Melting and boiling point increases
 Density increases
 Reactivity decreases
 Physical state at r.t.p. changes from gas (F2 ,Cl2) to
liquid (Br2) to solid (I2, At2)
 Colors intensity increases (from bright to dull
colors)
on
Physical Properties:
 Shiny
 Soft, can be cut by knife
 Relatively low melting and boiling points
 Relatively low densities
 Form ionic compounds that are soluble in water
Group VII elements (also known as “halogens”)
Group 0 trends:
Going down the group,
 Melting and boiling point increases
Physical Properties:
 Low melting and boiling points
 Colored molecules
Observations when Group I metal reacts with
cold water:
ny
Chemical Properties:
 Reactive non-metals
They react with metals to form ionic compounds,
and with non-metals to form covalent compounds.
 Strong oxidizing agents
They accept one valence electron readily to form
halide ions.
 Displacement reaction
A more reactive halogen can displace the less
reactive halide ion from its solution.





Group 0 elements (also known as “noble gases”)

Da
Physical Properties:
 Low melting and boiling points
 Mono-atomic elements
 Colorless gases at room temperature
 Insoluble in water
Metal piece darts about quickly around the
surface of the water
Effervescence observed (due to hydrogen gas
produced in water)
Hissing sound heard (due to the vigorous
release of hydrogen gas)
Metal piece becomes smaller in size
Flame observed (for some metals)
o Lithium (no flame)
o Sodium (yellow flame)
o Potassium (lilac flame)
pH of solution increases from pH7 to pH14
(due to the formation of a strong alkali)
Chemical Properties:
 Un-reactive
They do not react to form compounds as they have
full outer shell of electrons.
Due to their inert nature, noble gases are used when an
inert atmosphere is required.
 Argon used to fill light bulbs
 Helium used for filling balloons and airships
 Neon used in making neon lights
O Level Chemistry – The Periodic Table
3
Danyal Education (Contact: 9855 9224)
“A commitment to teach and nurture”
Neon
Argon
Xenon
Radon
as a component of breathing gases in diving tanks
to fill party balloons, hot air balloons
as inert gas shield during welding
to maintain superconductors at very low temperature
in gas chromatography
-
neon lights
fog lights
luminous warnings
advertising signs
-
to fill fluorescent and incandescent light bulbs
to fill double-pane insulated windows
as inert gas shield during welding
to flush out melted metals to eliminate porosity in casting
to provide an oxygen-and-nitrogen free environment for
annealing and rolling metals and alloys
-
to fill double-pane insulated windows
krypton lasers
as a gas component within halogen sealed beam headlights
to fill high performance light bulbs
-
to fill incandescent light bulbs
development in X-rays (when mixed with oxygen, it can
enhance the contrast in Computer Tomography (CT) imaging
as a gas filler in plasma display panels
-
in radiotherapy
to treat arthritis
Da
Krypton
Uses
-
ny
Noble Gas
Helium
al
Ed
uc
ati
on
Reaction of alkali metals with water
O Level Chemistry – The Periodic Table
4
Danyal Education (Contact: 9855 9224)
“A commitment to teach and nurture”
Going down Group VII, the reactivity of the element
decreases. Explain why.
Going down Group I, the number of electronic shells
increases, and atomic radius increases. As the
valence electrons get further from the nucleus, the
attractive force between the positive protons in the
nucleus and the negative valence electrons decreases.
Hence, going down Group I, the atom loses its
valence electrons more readily and becomes more
reactive.
Going down Group VII, the number of electronic
shells increases, and atomic radius increases. As the
valence electrons get further from the nucleus, the
attractive force between the positive protons in the
nucleus and the negative valence electrons decreases.
Hence, going down Group VII, it becomes more
difficult for the atom to gain a valence electron and
reactivity decreases.
Going down Group I, the melting point and boiling
point of the element decreases. Explain why.
Going down Group VII, the melting point and
boiling point of the element increases. Explain why.
Going down Group I, the number of electronic shells
increases, and atomic radius increases. As the
valence electrons get further from the nucleus, the
attractive force between the positive protons in the
nucleus and the negative valence electrons decreases.
Group VII elements exist as diatomic molecules with
simple molecular structure.
Going down Group VII, the molecular mass of the
element increases.
An increasing molecular mass results in increasing
intermolecular force of attraction.
The higher the intermolecular force of attraction, the
higher the energy needed to break these bonds, hence
higher melting point and boiling point.
al
Ed
uc
ati
on
Going down Group I, the reactivity of the element
increases. Explain why.
Hence, going down Group I, the metallic
bonding formed between the positive protons in
the nucleus and the delocalized valence
electrons becomes weaker.
The weaker the metallic bond, the lesser the
energy needed to break these bonds, and the
lower the melting point and boiling point of
these elements.
Going from left to right in the same Period, the
atomic radii of the elements decreases.
Da
ny
Elements in the same Period have the same number
of electronic shells, but the proton number of the
elements increase from left to right.
As number of protons increase, the attractive force of
the positive protons on the negative electrons also
increases, pulling the electronic shells closer
together.
Hence atomic radius decreases from left to right in
the same Period.
Group VII elements
Element
Color and physical state at room
temperature
Fluorine
Pale yellow gas
Chlorine
Yellowish-green gas / Pale green gas
Bromine
Dark red liquid
Iodine
Dark grey crystalline solid
Astatine
Black solid
Color in
gaseous state
Pale yellow
Yellowish-green
Brown
Purple
Black
Color in aqueous
state
# Reacts with water
Pale green
Reddish-brown
* Brown
(insoluble)
#Fluroine reacts explosively with water to form hydrogen fluoride and oxygen.
* Iodine is only slightly soluble in water, but solubility increases if iodide ions are present in the water. For eg, if
aqueous KI is added to the water, the presence of iodide ions enable the iodine to become soluble.
O Level Chemistry – The Periodic Table
5