Name: ………………………………………………
Class:…………………………………………………
Number:……………………………………………
TURK MAARIF COLLEGE
Academic Year 2020-2021
Number:……………………………………………
10th Forms-1st Term
Chemistry Section 2
ATOM AND PERIODIC SYSTEM
ATOMIC MODELS
The smallest particle found in matter is called atom. Since ancient times, many
theories and atomic models have been developed describe atoms.
1)Dalton Atomic Model
Dalton discovered that atoms were fullfilled spherical
solid particles that cannot be divided, created or
destroyed. He proposed his theory by showing that there
are same type of atoms in an element and different
elements can combine together forming compounds.
At this time-zoneprotons were discovered by Rutherford and electrons
were discovered by J.J.Thomson.
2)Thomson Atomic Model (Plum Pudding Model)
Thomson developed a ‘Plum Pudding Model’ and
believed that atoms contain high mass of
positively charged particles (protons) and small
mass of negatively charged particles (electrons)
which were randomly found in the positive part
just like plums in a pudding. He also believed that atoms were neutral so number
of positively charged particles was always equal to the number of negatively
charged particles.
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3)Rutherford Atomic Model (Nuclear Model)
Rutherford investigated α-Beams Experiment
(Gold Leaf Experiment) and found that there’s
a positively charged part in the centre of the
spherical atom and called it nucleus. He believed
that protons were in the nucleus and electrons
were orbiting around the nucleus.
At this time-zone neutrons were discovered by Chadwick.
4)Bohr Atomic Model (Planetary Model)
Bohr believed that atoms were spherical,
protons and neutrons were in the nucleus of
atoms and electrons were found in the
circular energy levels around the nucleus
at 2-D positions-just like planets around
the sun.
5)Modern Theory (Quantum Model)
Schrödinger and Heisenberg discovered that
protons and neutrons are in the nucleus of atom
while electrons are found in the specific energy
levels in a cloud at 3-D positions around the
nucleus. Orbitals are regions in which electrons
have high probabilities to be found.
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SUB-ATOMIC PARTICLES
There are 3 types of sub-atomic particles inside an
atom: protons, neutrons and electrons
First sub-atomic particle discovered was electron.
Electrons were discovered by J.J.Thomson. Secondly,
protons were discovered by Rutherford and they were
told the name ‘protons’ by Stoney. Presence of
neutrons were predicted by Rutherford but discovered
by Chadwick.
Table 1. The charge, mass and locations of the particles in an atom.
Particle
Charge
Mass
Location
Proton
____
1
Nucleus
Neutron
0
____
Nucleus
Electron
-1
0 (1⁄1840 - almost zero!)
In the _______
around the _______
Be CaReFuL!
Remember the charge on each particle like this:
p+

Protons are positive

Neutrons are neutral

(so electrons must be negative ones)
n
e-
Atomic Number (Proton Number)
Each element has its own atomic number which is also called proton number.
It is just like ID Card number and special for each element.
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It tells us how many protons there are in one atom of that element. Thats why
it’s also called nuclear charge.
Atoms of the same element have same number of protons. Different elements
have different number of protons.
Elements are arranged according to their atomic numbers in the Periodic Table.
Atomic number = number of protons
If atoms are neutral, the positive charges are cancelled out by an equal number
of negative charges. So, number of protons (+) must always equal the number of
electrons (-)
In neutral atoms, Atomic number = number of protons = number of electrons
Mass Number
Mass number is sum of protons and neutrons in the nucleus of an atom.
Mass number = Number of protons + Number of neutrons
So, Number of neutrons = Mass number - Number of protons
Example:
Example:
Lithium’s atomic number is 3 and its mass number is 7.
How many protons, neutrons and electrons are in a lithium atom?
Li

Number of protons = Atomic number = 3

Number of electrons = Number of protons = 3

Number of neutrons = Mass number – Number of protons = 7 – 3 = 4
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Atomic Symbols
We can write the symbol
7
3𝐿𝑖
to show the mass number and atomic number of
lithium.
Atomic number is written at the bottom and to the left of the symbol.
Mass number is written at the top and to the left of the symbol.
Mass
number
Li
7
Atomic
number
3
Symbol
Number of neutrons = 7-3 = 4
Exercise: Find number of protons, neutrons and electrons for
Atomic Symbol
Number of
protons
Number of
neutrons
40
20𝐶𝑎
and
31
15𝑃 .
Number of
electrons
𝟒𝟎
𝟐𝟎𝐂𝐚
𝟑𝟏
𝟏𝟓𝐏

Positively and negatively charged particles are called ions. The number
written at upright of the element symbol is called ‘Net Charge’

Ions can be monatomic such as S2- and Mg2+ and polyatomic such as NH4+
and CO32-.
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
When an atom loses electron, it gets positive charge and becomes a
cation.
Ex.: 13Al3+ has 13 protons and 13-3=10 electrons.

When an atom gains electron, it gets negative charge and becomes an
anion.
Ex.:8O2- has 8 protons and 8+2=10 electrons.

Neutral atoms, have same number of protons and electrons.

Cations have more protons than electrons since they lose electrons

Anions have more electrons than protons since they gain electrons.
Exercise:
56
X2+ has 32 electrons. How many neutrons are there in its
24
X2- has 24 neutrons. How many electrons does it have?
nucleus?
Exercise:
Exercise:
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Isotopes
Atoms of an element with same number of protons and different number of
neutrons are called isotopes.
Since isotopes have same number of protons, they have same atomic numbers.
Isotopes have different number of neutrons. That’s why they have
different mass numbers.
Since isotopes have different masses, they have different physical
properties. For example, the ones with larger mass will have higher
melting and boiling points.
Isotopes have same number of electrons, so they have similar chemical
properties since they are in same group of the periodic table and react similarly
to each other.
Exercise:
Find number of protons, neutrons and electrons in these two isotopes of
chlorine:
35
17𝐶𝑙
and
37
17𝐶𝑙
Atomic Symbol
Number of
protons
Number of
neutrons
Number of
electrons
𝟑𝟓
𝟏𝟕𝐂𝐥
𝟑𝟕
𝟏𝟕𝐂𝐥
Isotones:
Elements that have same number of neutrons but different number of protons
are called isotones.
Example:
20
12𝑋
and
17
9𝑌 
They both have 8 neutrons while their atomic and so
proton numbers are different!
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Isobars: Elements that have mass number but different atomic number are
called isobars.
Example:
𝟐𝟐
𝟏𝟑𝒁
and
𝟐𝟐
𝟒𝑻
They both have 22 as mass number while their atomic
numbers are different!
Isoelectronics: Different elements that have same number of electrons are
called isoelectronics.
𝟏𝟔 2- 𝟐𝟎
𝟖𝑶 , 𝟏𝟎𝑵𝒆
and
𝟐𝟒
2+
𝟏𝟐𝑴𝒈 
They both have 10 electrons while their atomic
numbers are different so they are different elements! They just have same
electronic configurations since some of them are ions!
Exercise.Look at the groups of elements below and decide if they are isotopes,
isotones, isobars, or isoelectronics.
a) 35
17𝐶𝑙 and
b) 24
10𝑋 and
37
17𝐶𝑙
24
15𝑌
c) 168𝑂2− and
39 3+
d) 13
𝐴𝑙 ,
: _______________________
: _______________________
24
2+
12𝑀𝑔
14 3−
7𝑁
and
: _______________________
23
+
11𝑁𝑎
: _______________________
PERIODIC SYSTEM
Periodic Table

Periodic Table is a list of all of the elements. It’s a little like
the alphabet of chemistry.

The periodic table tells us several things… Elements were first
arranged according to their mass numbers by Mendelev and
Meyer.. But this was not the right one! The right version was developed by
Moseley in which elements were arranged according to their atomic numbers.
Appearance of a Oxygen atom in the Periodic Table is shown in the diagram
below.
Mass number (number
of protons + number of
neutrons)
Atomic number (number
of protons)
O
16
8
8

Atomic Number: 8

Mass number: 16

Number of protons: 8

Number of electrons: 8

Number of neutrons: 16-8=8
The Arrangement of the Electrons

Electrons move in the shells or energy levels around the nucleus.

Each energy level can hold a certain number of electrons.

Inner energy levels are always filled before outer ones.

1st shell can hold 2electrons.

2nd shell can hold 8 electrons,

3rd shell can hold 8 electrons and but it can expand upto 18 electrons.

4th shell can hold 8 electrons.
3rd shell can hold
8 electrons but it can
expan upto 18 electrons
2nd shell can hold 8 electrons
1st shell can hold 2 electrons
The First 20 Elements in the Periodic Table
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There are Two Ways to Represent Atomic Structure:
A. Electronic Configurations
With electronic configuration elements are represented numerically by the
number of electrons in their shells and number of shells.
So Let’s try it...
How to draw a Lithium atom
1. First, look at the Periodic Table
2. Then, determine the number of protons and hence
the number of electrons.
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(There are _3_ protons, and so _3_ electrons
in a neutral lithium atom)
3. Arrange the electrons in levels, always filling up an
inner (lower energy) level.
3
Li
These will be arranged _2_ in the first level and _1 in the second level. This is
written as _2 , 1_.
Write the electronic configuration for the following elements;
𝟒𝟎
a) 𝟐𝟎𝐂𝐚=2,8,8,2
b)
d) 𝟐𝟑
𝟏𝟏𝐍𝐚=2,8,1
𝟏𝟔
𝟖𝐎=2,6
e) 𝟑𝟓
𝟏𝟕𝐂𝐥=2,8,7
𝟐𝟖
c) 𝟏𝟒𝐒𝐢=2,8,4
f) 𝟏𝟏𝟓𝐁=2,3
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B) Dot & Cross Diagrams
With dot & cross diagrams elements and compounds are represented by dotsor
crosses to show electrons, and circles to show the shells.
For example;
14
N
7
O
16
8
35
17
Cl
8
2,6
2.8.7
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Draw diagrams to show the arrangement of the electrons in
𝟏𝟐
a) Carbon atom - 𝟔𝐂 =2,4
𝟑𝟎
b) Phosphorus atom- 𝟏𝟓𝐏
𝟐𝟎
c) Neon atom - 𝟏𝟎𝐍𝐞
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
In the Periodic Table, vertical columns are called __groups___ and horizontal
rows are called _period__

Elements were first divided into groups of 3 members-according to their
similarities as called ‘Triats Rule’ by Döbereiner. Then, they were divided
into groups of 8 as called ‘Octaves Rule’ by Newlands.

In modern periodic table, elements are arranged in 18 groups and 7 periods,

Groups contain elements with ___similar chemical__ properties.

Elements in the same group have the same number of __electrons___ in the
outer levels. The number of ____electrons___ in the outer level is same as
the __group__ number.
Group 1
6th Period

Barium is in Group 2 so it has _2_ electrons in its outer level.

The group 0 elements are known as the __noble__ __gases__ because
they are almost completely ___stable___. thought of as being ‘full’
levels.

Number of energy levels (shells) of an atom is the _period_ of it.

Calcium is in 4th period so it has _4_ energy levels (shells) around its
nucleus.
Group 8
4th Period
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Classification of Elements According to
Their Place In Periodic Table
Left part of the periodic table is mostly metals while right part is non-metals.
Group 1,2, and most of group 3 elements are metals while group 5,6 and 7
elements are mostly non-metal.
There are some elements that show both metallic and nonmetallic properties as
well-they are called semi-metals or metalloids. (B, Si, Ge, As, At, Te, Po,
Sb)
Noble gases (Group 8 or 0) don’t show any metallic or nonmetallic property since
they are very unreactive, have full outer shell and monatomic gases at room
temperature.
APPEREANCE
MALLEABILITY
MELTING AND
BOILING POINT
CONDUCTIVITY
STATE AT RTP
CHARGE
METALS
Shiny
Malleable(Easy to
change shape)
High
NONMETALS
Dull (Matte)
Brittle
Good conductor of
heat and electricity
Solid except mercury
which is liquid
Gets positive charges
Poor conductor of heat and
electricity
Solid, liquid or gas
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Low
Gets negative charges
DUCTILITY
HARDNESS
COMPOUNDS
OXIDES
Ductile(Easy to pull
into wire)
Hard
Forms ionic
compounds with
nonmetals
Basic oxides
Brittle
Usually soft except carbon in
structure of diamond
Forms ionic compounds with
metals and covalent compounds
with nonmetals
Acidic oxides
GROUP 1-Alkali Metals
They have 1 electron in their outer shell, want to lose 1 electron to have a
full outer shell so get 1+ charge. They are all metals at room temperature.
GROUP 2-Alkaline Earth Metals
They have 2 electrons in their outer shell, want to lose 2 electron to have a
full outer shell so get 2+ charge. They are all metals at room temperature.
TRANSITION METALS
Between group 2 and 3, there are transition metals that form colourful
compounds and can get different electrical charges such as iron, copper,
nickel, gold, silver, zinc, etc..
GROUP 7-HALOGENS
They have 7 electrons in their outer shell, want to gain 1 electron to have a
full outer shell so get 1- charge. They are all diatomic non-metals at room
temperature that can be solid, liquid or gas.
GROUP 8-NOBLE GASES
They have full outer shells with 2 or 8 electrons so don’t want to lose or
gain any electrons, they are stable, very unreactive monatomic gases at
room trmperature.
HeliumUsed in balloons
NeonUsed in advertising lights
ArgonUsed in light bulbs
XenonUsed in lasers
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TRENDS IN THE CHANGE OF PERIODIC PROPERTIES
There’s a regular change in some physical and chemical properties of elements
within a group or period.
Some of these properties are;
1. Metallic and nonmetallic properties
2. Atomic volume or size
3. Ionisation energies
4. Electron affinity
5. Electronegativity
1.Metallic and nonmetallic properties:
When you go down a group, metallic properties increase since metals get
stronger as number. of protons and their nuclear attraction power
increase. While, when you go across a period metallic properties decrease
since there are more non-metals and
finally noble gases! As metallic
properties increase, non-metallic
properties decrease! That’s why nonmetallic properties decrease down a
group and increase across a period.
Example. Arrange metallic properties of following elements in ascending
order:
12Mg, 13Al, 19K, 20Ca
Example. Arrange nonmetallic properties of following elements in
descending order: 3Li, 8O,
17Cl
2.Atomic volume or size
Atomic size is distance between nucleus and outermost shell of an atom.
Atomic size increases as you go down a group since no. of shells increase.
When you go right a period, atomic size decrease since no. of shells stay
same but no. of protons and attraction force of nucleus increase and it
gets outer electron more.
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When we look at isoelectronics, the greater the
nuclear charge, the smaller ionic radius since as
nuclear charge attracts atoms get smaller! 😊
Example. Arrange atomic sizes of following
elements in ascending order: 9F, 5B,
19K, 13Al
Example. Arrange nonmetallic properties of following ions in descending
order:
10Ne, 8O
2-
, 7N3, 13Al3+,
2+12Mg
3.Ionisation energies
1st ionisation energy is amount of energy to remove 1 electron from
gaseous state of an element.
To remove 2nd electron, we need 2nd ionisation energy and so on..
X(g) + I1
X+(g) + e-
X+(g) + I2
X2+(g) + e-
X2+(g) + I3
X3+(g) + e-
As you go down a group, ionisation energy
always decrease. Because, as atomic radius
gets larger, its gets easier to remove an
electron from outer shell since nuclear
attraction between nucleus and outer
electron decrease. So, we need less energy
to overcome these forces.
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As you go across a period, ionisation energy usually increase. Because, as
atomic radius gets smaller it gets harder to remove an electron from
outer shell since nuclear attraction increase. So, we need more energy to
overcome these forces.
But; there are 2 exceptions between groups 2-3 and 5-6. Because,
between these groups electrons also change their orbitals and we need
less energy to remove them suddenly.
So the correct order is this for ionisation energy
1A<3A<2A<4A<6A<5A<7A<8A
Example. Arrange ionisation energies of following elements in ascending
order: 11Na, 15P,
16S, 17Cl
Example. Arrange ionisation energies of following elements in descending
order:
17Cl, 16S, 13A, 12Mg
Be CaReFuL! 😊
You can understand group of a substance by looking at its different
ionisation energy values. If there’s a big jump between 2 ionisation
energies, that means at that point the electron changes its energy level
and now its not at the outer level more, it’s closer to the nucleus, it
requires more energy to remove the electron so the number of ionisation
energy at first step of jump is group number.
18
Example.Look at the ionization energies of element X shown below and
decide its group number. Explain your answer.
Ionisation energies (kJ/mol) 1st IE 2nd IE 3rd IE 4th IE 5th IE
Element X
1260
2260
6940
7300
8500
____________________________________________________________
____________________________________________________________
____________________________________________________________
____________________________________________________________
4.Electron affinity
It is the amount of energy given out when an electron is added to the gaseous
state of an element.
X(g) + e-
X-(g) + E
It is similar to ionisation energy and decreases down a group, increases acrossa
period. In a period, halogens have the highest electron affinity since noble
gases don’t have any electron affinity value.
Example. Arrange electron affinities of following elements in increasing order:
12Mg, 14Si, 15K, 18Ar
19
5.Electronegativity
A measure of the tendency of an atom to attract electrons in a molecule is
called electronegativity. It’s influenced by same factors with electron affinities.
In general, the greater electron affinity of an atom; the greater
electronegativity of an atom. F, O, N are elements with strongly high
electronegativities. On the other hand, inert gases have no or little
electronegativities.
Example. Arrange electronegativities of following elements in decreasing order:
5Ne, 8O, 7N, 13Al, 15P
SUMMARY
Properties that increase down a group and decrease across a period 
Atomic size, metallic properties
Properties that decrease down a group and increase across a period 
Ionisation energy, nonmetallic properties, electron affinity, electronegativity
(There are exceptions)
GOOD LUCK! 😊
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