https://www.youtube.com/watch?v=I5H1SeepnaU
Genius of Mendeleef
https://ed.ted.com/lessons/the-genius-of-mendeleev-s-periodic-table-lou-serico
Revising Periodic Table WS
http://www.rsc.org/learn-chemistry/resource/res00001090/revising-the-periodic-table
http://www.rsc.org/periodic-table
Periodic Videos- Videos Of all the elements in the Periodic Table
https://www.youtube.com/watch?v=6rdmpx39PRk&list=PL7A1F4CF36C085DE1
Periodic Table Animation of All the Elements
Investigating the Periodic Table with Experiments - with Peter Wothers
DOBEREINER'S
TRIADS
Elements were arranged in increasing order of atomic
masses in groups of
three called triad; in which the atomic mass of the middle element was
the average of the atomic mass of the other two elements.
E.g. Li, Na, K ; Cl, Br, I formed a triad.
Limitations of
Dobereiner's Triads
All the elements
could not be arranged in the form of triads.
NEWLANDS’ LAW
OF OCTAVES
According to the Newlands' Law of octaves “When elements are
arranged in increasing order of their atomic masses, the properties of every
eighth element was similar to that of first element and so on.”
Limitations of
Newlands’ Law of Octaves
The law was applicable only for lighter elements.
MENDELEEV’S PERIODIC LAW
Mendeleev’s Periodic Law states that
‘The physical and chemical properties of elements are the periodic function of
their atomic masses.’
Merits of
Mendeleev’s Periodic Table
1.
Mendeleev
left gaps for undiscovered elements and also predicted their properties.
2.
Incorrect atomic mass of some
elements was corrected.
Limitations of
Mendeleev’s Periodic Table
1. Correct position of hydrogen
could not be assigned.
2.
Anomalous position of isotopes.
3.
Anomalous pairing of elements: In
certain places in the periodic table Mendeleev placed an element with slightly
larger atomic mass before an element with slightly lower atomic mass.
E.g. Co (58.93) was placed before Ni (58.71)
4.
Position of rare earth and actinides could not be justified.
MODERN PERIODIC
LAW
‘The physical and chemical properties of elements are the
periodic function of their atomic numbers.’
Salient
Features of Modern Periodic Table
1. The vertical columns in the Modern
periodic table are called groups. There are 18 groups.
2. The horizontal rows in the Modern
periodic table are called periods. There are 7 periods in the Modern Periodic
Table.
3. There are 2 separate rows at the bottom of the
periodic table each containing 14 elements. The first row is called the
lanthanide series and the second row the actinide series.
Table
1: No. Of Elements In Each Period
|
Period |
No. of
elements in each period |
Name of
the period |
|
1 |
2 |
Very short |
|
2 |
8 |
Short |
|
3 |
8 |
Short |
|
4 |
18 |
Long |
|
5 |
18 |
Long |
|
6 |
32 |
Very long |
|
7 |
|
|
BRIDGE ELEMENTS
Elements of second period are known as bridge elements. They
resemble in certain properties with the elements of the third period placed
diagonally.
E.g. Li has similarities with Mg.
PERIODIC PROPERTIES
The occurrence of characteristic properties of elements at
definite intervals in the modern periodic table is called periodicity in
elements. The properties which appear at regular intervals in the periodic
table are called periodic properties. Chemical properties of an element depend
on the electronic configuration of the elements. As we move across a period the
electronic configuration is found to change so the chemical properties also
change gradually. As we move down a group the number of valence electron
remains same so there is a similarity in the chemical properties of the
elements in a group.
Atomic
Radius
The distance between the centre of the nucleus of an atom
and its outermost shell is called the atomic radius.
Factors
affecting the Atomic Radii
1.
Number of shells:
Atomic Radii α No. of shells
As the number of shells increases the
distance between the outermost shell and the nucleus
increases. So the atomic size increases.
2.
Nuclear Charge:
Atomic Radii α
Nuclear Charge is the positive charge on the nucleus of an
atom.
As the nuclear
charge increases the electrons in the outermost shell are attracted with increasing force. Thus the outermost shell
comes closer to the nucleus.
Ionic Radii
a.
Radius of a cation:
Radius of a cation is less than the corresponding atom as the number of shells
is less in the cation than in the atom.
E.g.
|
|
K |
K+ |
|
Electronic configuration: |
2,8,8,1 |
2, 8, 8 |
|
Number
of shells: |
4 |
3 |
Therefore atomic size of K+ is less than K.
b. Radius of an anion:
Radius of an anion is more than the corresponding atom as an extra electron is
added to the valence shell which results in an increase in inter electronic
repulsion and a decrease in the effective nuclear charge.
|
Note: The atomic radii of an inert gas are larger than the
preceding elements. This is because the
valence shell of an inert gas is complete and so the inter electronic
repulsion is maximum
|
Ionisation
Potential
It is defined as the amount of energy required to remove a
loosely bound electron from the valence shell of an isolated gaseous atom in
its ground state.
Ionisation energy can be expressed in terms of kJmol-1 or
eV [electron volt].
M (g) + IE → M+
(g) + e-
Successive
Ionisation Potential
M + IE1 → M+
+ e-, IE1 is 1st
ionisation energy
M+ + IE2 →
M+2 + e-,
IE2 is 2nd
ionisation energy
M+2 + IE3
→ M+3 + e-,
IE3 is 3rd
ionisation energy
IE1 < IE2 < IE3
This is because it is more
difficult to remove an electron from a positively charged ion than from a
neutral atom (in case of IE2 electron has to be removed from a
cation).
Factors
Affecting Ionisation Potential (IP)
1.
Atomic size:
IP α
The force of attraction on the valence
shell increases as the atomic size decreases. So ionisation
energy increases.
2.
Nuclear charge:
I P α Nuclear charge
The force
of attraction on the valence electron increases so the electrons are held more
firmly and thus ionisation
energy increases.
|
Note: He has the highest ionisation energy (2372kJ/mol and Cs has the
lowest ionisation energy (375 kJ/mol). Metals usually
have low ionisation energy and non- metals have high ionisation energy.
|
Electron affinity (EA):
The amount of energy released
when an isolated atom in the gaseous state accepts an electron to form an
anion.
X (g)
+ e- → X- (g) + EA
EA can be either positive or
negative. The process of adding an electron to the atom can be either
endothermic or exothermic. Electron affinity is positive for noble gases. EA is
expressed in kJ/mol or eV.
Factors
Affecting Electron Affinity (EA)
1.
Atomic radius:
EA α
A small atom accepts an electron more readily
than a large atom as the nucleus exerts a greater force of attraction on the
electrons.
2.
Nuclear charge: EA α Nuclear charge
Note:
EA of chlorine is more than fluorine. EA of sulphur is more than that of
oxygen. This is because of the small size of the atom, inter electronic
repulsion increases and hence EA decreases.
EA
is highest for halogens and least for alkali metals.
Electronegativity
It is defined as the tendency of an atom in a molecule to
attract the shared pair of electrons towards itself.
Factors
Affecting Electronegativity
1. Atomic radius: Electronegativity α
2. Nuclear Charge: Electronegativity α
Nuclear Charge
|
Note: Noble gases have zero
electronegativity. Fluorine
is the most electronegative element and caesium the least. |
Metallic
and Non Metallic Properties
The tendency of an atom to loose electrons to form cations
is called metallic character or electropositive character.
The tendency of an atom to gain electrons by to form anions
is called non-metallic character or electronegative character.
Factors Affecting Metallic and Non-metallic
Character
1. Atomic radius:
Metallic
character α
Atomic radius
Non
–metallic character α
2. Ionisation
potential:
Metallic
character α
Non
–metallic character α
Chemical Reactivity
The
reactivity of elements depends upon their tendency to loose or gain electrons
to complete their outermost orbit.
Greater
the tendency to loose electrons greater is the reactivity of the metals.
Greater the tendency to gain electrons greater is the reactivity of non-
metals.
Variation of Periodic
Properties Across the Period and Down the Group
Across the
Period
Ionisation potential, electron affinity, electronegativity, non-metallic
character and the number of valence electrons increases across the period.
Atomic radius and metallic character decreases across the period. Chemical reactivity
of an element first decreases and then increases across the period.
Down the Group
Ionisation potential, electron affinity, electronegativity and
non-metallic character decreases down the group. Atomic radius and metallic
character increases down the group. The number of valence electrons remains
same. Chemical reactivity of metals increases because the tendency to
lose electrons increases and reactivity of non-metals decreases because the
tendency to gain electrons decreases.
|
Note:
Across a period valency of an element first
increase from 1 to 4 and then decreases to 0. Down the group valency
remains same. |
PREDICTION OF
GROUP, PERIOD AND BLOCK OF A GIVEN ELEMENT
Period = The total number of shells in an atom of an
element.
Block = The orbital in which the last electron enters.
Group:
For s block elements= No. of valence electron
For p block elements = No. of valence electron + 10
For d block elements= No. of electrons in (n-1) d orbital +2
|
Exercise 1 1.
Predict
the block, period and group to which the following element belongs. A (Z = 10), X (Z = 26), Y (Z
=19), B (Z = 14), C (Z = 47) |
|
Note: · The
elements in the first group (Li, Na, K, Rb,Cs & Fr) are called alkali
metals. The elements in the second group (Ca, Sr, Ba …) are called alkaline
earth metals. The elements in the 17th group (F, Cl, Br, I, At) are called
halogens.
|
PERIODIC TABLE: SAMPLE TEST - 1
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