2. From the Syllabus
the structure of the periodic table is based on the atomic number
and the properties of the elements
the elements of the periodic table show trends across periods and
down main groups, including in atomic radii, valencies, 1st
ionisation energy and electronegativity as exemplified by groups
1, 2, 13–18 and period 3
3. Text Book references.
Pearson Chemistry 11 Chapter 3.2 p47 and 3.3 p53.
Essential Chemistry Chapter 9 p70
STAWA Sets 9/10/11
4. Organisation
All the chemical elements are arranged in order of increasing
atomic number.
Atomic number = number of protons in the nucleus
Elements are arranged into rows and columns depending on their
electronic configuration and their chemical properties.
The organisation of the Periodic Table, arose from four major
principles.
5. Modern Periodic Table
1. Atomic number, rather than atomic mass, was to determine
the order.
2. Electrons around the nucleus were arranged in increasing
energy order and repeating patterns of electron
configuration were observed.
3. Valence electrons (outer shell) determined the chemical
properties.
4. Recurrence of similar properties in groups were observed.
6. Periods are the horizontal rows of the periodic table
The period number of an element is the same as the number of
shells
occupied by the electrons in its atoms.
Periods
7. 1 2 3 4 5 6 7 8 9 1
0
1
1
1
2
1
3
1
4
1
5
1
6
1
7
1
8
Groups
Groups are the vertical columns. Valence electrons
can be worked out from the group number ( for
group 1, 2 and 13 – 18 only) as the last number
equals the valence electrons.
8. Group Names
There are special groups that are found in the Periodic Table.
Elements found in these groups exhibit similar properties and
are placed in these groups due to their electron
configuration.
Group 1 – Alkali metals
Group 2 – Alkaline Earth metals
Group 17 – Halogens
Group 18 – Nobel gases
Groups 3 – 12 – Transition metals
Lanthanides and Actinides (located at the bottom)
9. Elements
The electron configuration of an element tells you the period and
the group.
SODIUM: has 11 electrons and an electron configuration of
2,8,1. This configuration shows that sodium atoms have 3
shells and 1 valence electron. Therefore, sodium is placed in
Period 3, group 1.
NITROGEN: has 7 electrons and an electron configuration of
2,5. This configuration shows that nitrogen atoms have 2
shells and 5 valence electrons. Therefore, nitrogen is placed in
Period 2, group 15.
10. Metals/non-metals
Elements are classified as metals and non-
metals.
In the Periodic Table the metals are found
on the left hand side and the non-metals
are found on the right hand side. They are
separated by the zigzag line.
Some elements near this line
exhibit both properties and are
called metalloids.
11. Trends
There are many different trends that can be observed in the Periodic
Table
Some of the trends observed are:
Chemical reactivity
Atomic size
Ionisation energy
Electronegativity
Metallic and non-metallic nature
12. Core Nuclear Charge
Core charge or Core Nuclear charge can be defined as the net or
effective positive charge influencing the outer shell electrons.
It is the number protons in the nucleus minus the number of inner
shell electrons. It increases across the periodic table groups but
remains the same down a group.
Lithium, Sodium and Potassium have the same core charge of +1
Fluorine, Chlorine and Bromine have a core charge of +7
Core charge can be used to explain the trends in the periodic table.
13. Atomic size
An atom doesn’t have a defined boundary to set the limit of its
size
The radius of an atom cannot be measured directly as they are
so small.
However, there are methods available to gain an estimation of
the sizes of atoms relative to each other. Generally by measuring
the space between two nuclei it is possible to calculate the
radius.
14. Atomic size
Atomic size increases down the group of the
Periodic Table
As we move down a group, electrons are added and
therefore more shells are added.
15. Atomic size
Atomic size generally decreases from left to right across a period.
Across a period, each atom has its valence electrons in the same
energy shell. However, each element has one more proton than the
previous one.
The Core Charge increases across the period.
This creates an attraction between the positive nucleus and the
negative electrons. As we add more protons, the attraction becomes
greater and therefore the electrons are pulled closer to the nucleus.
16. Electronegativity
The electronegativity of an element is the measure of how strongly
an atom can attract an electron to itself.
This is particularly important when elements are bonded to each
other as it determined if they share electrons like in covalent bonds
or lose/gain electrons like in ionic bonds.
The degree of attraction depends on the number of protons in the
nucleus, core charge and the number of energy shells the electrons
are in, how far the electrons are away from the nucleus.
17. Electronegativity
Moving down the group the
electronegativity decreases.
This is because the electrons in each
subsequent shell is further away from
the positive nucleus and therefore are
not as strongly attracted, and are also
easier to remove.
Decreasing
18. Electronegativity
Electronegativity increase as you move across a period from left
to right.
This is due to an increase in the number of protons. Increase in
core charge. The greater positive charge in the nucleus attracts
the electrons to a greater extent, giving it a greater
electronegativity.
19. Electronegativity
Fluorine is more electronegative than oxygen because while their
valance electrons are in the same energy shell, fluorine has one more
proton in the nucleus than oxygen. The greater positive charge attracts
the electrons to a greater extent giving it a greater electronegativity.
Fluorine has a greater electronegativity than chlorine. This is because
chlorine has its valance electrons in the 3rd energy level /shell and
fluorine has its valence electrons in the 2nd energy level/shell. Even
though chlorine has more protons the distance between the protons
and the electrons is greater which means they are not as strongly
attracted giving chlorine a lower electronegativity.
20.
21. Ionisation energy
When atoms gain or lose electrons they become ions.
The energy required to remove an electron from a gaseous atom
is known as the ionisation energy.
Removal of one electron results in the formation of a positive ion
with a +1 charge. Therefore,
A(g) A+
(g) + e-
The energy required to remove the first (outermost) electron is
called the first ionisation energy.
22. Ionisation energy
To remove a second electron it requires more energy
than the first, as the atom is already a charged ion.
A+
(g) A2+
(g) + e-
The amount of energy required is known as the second
ionisation energy.
23. Ionisation energy
Moving down the group the ionisation
energy decreases.
This is because the electrons in each
subsequent shell is further away from the
positive nucleus and therefore are not as
strongly attracted, and easier to remove.
Decreasing
24. Ionisation energy
Ionisation energy increase as you move across a period from
left to right.
This is due to an increase in the number of protons. The
greater positive charge in the nucleus attracts the electrons to
a greater extent. This means more energy is needed to remove
the electrons.
25.
26.
27. Questions to do
Pearson Chemistry 11 Chapter 3.2 p52 and 3.3 p61.
Essential Chemistry Set 9 p50.
STAWA
WACE Study Guide 1.10 p17 – 1.15 p22.
28. Metallic Nature
Metals are characterised as materials that give up their
outer-shell electrons.
The metallic nature of an element is greater the easier it is to
remove its out-shell electron.
29. Metallic Nature
Moving down the group the metallic nature
increases
This is because the electrons in each
subsequent shell is further away from the
positive nucleus and therefore are not as
strongly attracted, therefore they are easier to
remove to form the characteristic positive ion
of a metal
Increasing
30. Metallic Nature
The metallic nature decreases as you move across a period
from left to right.
This is due to an increase in the number of protons. The
greater positive charge in the nucleus attracts the electrons to
a greater extent, making it harder to form the cation.
31. Non-Metallic Nature
Non-metals are characterised as materials that gain
electrons.
The non-metallic nature of an element is greater
when it easier it is to attract an electrons from
another element.
32. Non-Metallic Nature
Moving down the group the non-metallic
nature decreases
This is because the electrons in each
subsequent shell is further away from the
positive nucleus and therefore positive charge
of the nucleus is further away from the
valence shell, making it harder to attract the
electrons to form a negative ion
Decreasing
33. Non-Metallic Nature
The non- metallic nature increases as you move across a
period from left to right.
This is due to an increase in the number of protons. The
greater positive charge in the nucleus attracts the electrons
to a greater extent, making it easier to attract other
electrons to become a negative ion.
34. Chemical reactivity
The Chemical reactivity of an element is determined by how
easily it will lose or gain electrons.
This trend is directly related to the non-metallic and metallic
nature of elements as metals and non-metals react differently.
Group 17 will be the most reactive non-metals and Group 1 the
most reactive metals. Why is this so?.
