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1. Periodic classification of elements
Class X

2. Classification is the grouping of similar elements together and separating them from dissimilar ones.
In order to make the study of the elements known to us easy , the elements have been divided into
groups in such a way that those that have similar properties are placed in the same group.

3. Early attempts of the classification of elements:
Lavoisier
Lavoisier divided elements into two main types – metals and non metals.
But soon with the discovery of more elements it became clear that this classification was not enough.

4. Dobereiner’s Triads
• Dobereiner arranged certain elements with similar physical and chemical properties in groups of
three called triads.
• The basis of the arrangement was the atomic mass of the elements.
• In a particular triad, he showed that when the elements were arranged in the order of increasing
atomic masses, the atomic mass of the middle element is approximately the arithmetic mean of the
other two.

5. Triad of Li, Na and K
Some common properties of the elements of this triad:
i) All are metals
ii) All react with water to form alkalis and hydrogen .
iii) All are monovalent (i.e have a valency of 1).
• The atomic mass of sodium (Na) is arithmetic mean of first element lithium (Li) and third element
potassium (K).

6. Triad of Ca, Sr and Ba
The elements have a few common characteristics:
i) All are metals.
ii) All are divalent (i.e have a valency of 2)
iii) All impart a characteristic colour to the flame.
Mean of the atomic masses of the calcium (Ca) and barium (Ba) is almost equal to atomic mass of
strontium (Sr).
Arithmetic mean of atomic masses of calcium (Ca) and barium (Ba) =(40+137)/2=88.5
Actual atomic mass of the strontium (Sr) = 87.6

7. Triad of Cl, Br and I
Some common properties of the elements of this triad:
i) All are non metals and diatomic.
ii) All are highly reactive.
iii) All are monovalent.
iv) All react with water to form halogen acids.
Arithmetic mean of Chlorine (Cl) and Iodine (I) = (35.5+127)/2=81.25
Actual atomic mass of the second element = 80

8. Limitations of Dobereiner’s Triads
Dobereiner could identify only three triads from the elements known at that time.
As a result, this system of classification did not gain importance.

9. Newland’s Law of Octaves
• John Newlands arranged the 56 known elements in order of increasing atomic masses.
• He found that every eighth element had properties similar to the first. Thus there was a resemblance
in the properties of the first and the eighth, second and ninth, third and tenth and so on.
• He compared this to the octaves found in music and called it the Law of octaves.
According to the Newland’s Law of Octaves:
When elements are arranged in order of increasing atomic masses, there is a repetition of the
properties of every eighth element as compared to the first.

11. Limitations of Newland’s Law of Octaves
i) The law was found to be successful only up to the element calcium (atomic mass 40). After that every
eighth element did not have the same properties. For e.g Cu was placed below K in the same group.
ii) Newlands placed two elements in the same slot to fit elements in the table. For e.g Co and Ni were
placed in the same slot. Similarly La and Ce were placed in the same slot.
iii) When new elements were later discovered, their properties did not fit into the law of octaves.
iv) When the noble gases were discovered later, their inclusion in the table disturbed the entire
arrangement.
Thus Newland’s Law of Octaves worked well with lighter elements only.

12. Mendeleev’s Periodic Table
• Russian chemist who made significant contribution in the formation of periodic table.
• Mendeleev’s Periodic Law:
The properties of elements are a periodic function of their atomic masses.
• The repetition of similar properties after a definite gap of atomic masses is called periodicity in
properties.

13. Description of Mendeleev’s periodic table
• The 63 known elements were arranged in order of increasing atomic masses in vertical columns and
horizontal rows.
• There were 6 horizontal rows called Periods and a total of 8 vertical columns called Groups.
• Groups I to VII are subdivided into A and B subgroups. Group VIII did not have any subgroups.
• Elements with similar properties were placed in a Group. The formulae of the hydrides and oxides
formed by an element were treated as one of the basic properties for classification.
• All the elements in a particular group show regular gradation in their physical properties and chemical
reactivities.

15. Achievements of Mendeleev’s Periodic Table
1. Elements with similar properties were grouped together which made a systematic study of their
chemical and physical properties easier.
2. In some places elements with higher atomic masses had to be placed before those with lower atomic
masses. E.g. Co was placed before Ni and Te before I so that elements with similar properties could be
grouped together.
3. Atomic masses of some elements like Be (beryllium), Au (gold), In (indium) were corrected based on
their positions in the table. For example, the atomic mass of beryllium was corrected from 13.5 to 9.
4. Mendeleev predicted the existence and properties of some elements that had not yet been
discovered, on the basis of their position in his periodic table. He even named them by prefixing Eka
(Sanskrit numeral meaning one) to the name preceding the element in the same group. Scandium,
Gallium and Germanium were discovered later and their properties matched very closely with the
predicted properties of Eka - boron, Eka – aluminium and Eka – silicon respectively.
5. When the noble gases were discovered later, they could be placed easily in a new group without
disturbing the existing order.

16. Limitations of Mendeleev’s Periodic Table
1. Hydrogen resembles both the alkali metals and the halogens in properties.
Like the alkalis it combines with halogens, oxygen and sulphur and forms compounds having similar
formulae. E.g HCl and NaCl, H2O and Na2O, H2S and Na2S.
However like halogens it exists as diatomic molecules ( H2, Cl2)
So, Mendeleev could not justify its position.
2. Isotopes were discovered later and posed a challenge to the Mendeelev’s periodic law.
3. Atomic masses do not increase in a regular manner for e.g Cobalt (Co) has higher atomic weight but
was placed before Nickel (Ni) in the periodic table, similarly Te (Tellurium)was placed before I. This was
not explained by Mendeleev.

17. Modern Periodic Table
Modern Periodic Law given by Moseley:
Properties of elements are the periodic functions of their atomic numbers.
The periodicity or the repetition of similar chemical properties of elements is because of the repetition
of the same valence shell electronic configuration.

19. • Elements are arranged in order of increasing atomic numbers.
• The modern periodic table also has Groups and Periods.
• Groups:
There are 18 vertical columns or Groups, numbered from 1 to 18.
The elements in a group are separated by definite gaps of atomic numbers (8, 8, 18, 18, 32).
The elements in a group have the same number of valence electrons and valency.
The elements in a group have identical chemical properties.
The physical properties and the chemical reactivities show a regular gradation within the group.

20. • Periods:
There are 7 horizontal rows or periods.
Each period marks a new electronic shell getting filled.
In a period, the number of valence electrons increases by 1 unit as we move from left to right.
The number of elements in a period is based on how the electrons are filled into various shells.
First period has 2 elements and is called Very Short Period.
Second and Third periods have 8 elements each and are called Short Periods.
Fourth and Fifth Periods have 18 elements and are called Long Periods.
Sixth and seventh periods have 32 elements and are also called Long Periods.
14 elements ( Z = 58 to Z = 71) belonging to the 6th and 14 elements (Z= 90 to Z= 103) belonging
to the 7th Periods are placed at the bottom of the periodic table. These are known as Lanthanoids
and Actinoids respectively.

21. Advantages of the modern periodic table:
• The position of the elements in the periodic table makes it easy to predict and compare
their properties, and explains the reason for their specific position in the periodic table.
• It gives explanation for the periodicity of elements and tells the reason why all elements
in a group have similar properties, which differ from those of other groups.
• One position for all isotopes of an element is justified since isotopes have same atomic
number.
• The positions of some elements which were misfits in the Mendeleev’s periodic table are
now justified because it is based on atomic numbers and not atomic masses.

22. Trends in the Modern Periodic Table
1. Valency:
Valency (defined as the combining capacity of the atom of an element with the atoms of other
elements in order to acquire 8 e- or in some cases 2e- ) of an element is determined by the number of
valence electrons present in an atom.
Within a group, since the elements have same number of valence electrons, they have the same
valency.
The elements present in a period, have different valencies. The valency of elements in a period
first increases from 1 to 4 (as the number of valence electrons also increases from 1 to 4) and then
decreases from 4 to finally 0 for inert gases.
For e.g:
Element Na Mg Al Si P S Cl Ar
No. of valence e- 1 2 3 4 5 6 7 8
valency 1 2 3 4 (8-5=) 3 (8-6=) 2 (8-7 =) 1 (8-8=) 0

23. 2. Atomic size:
The atomic size is related to the radius of the atom ( defined as the distance between the centre of
the nucleus and the outermost shell of electrons of an isolated atom).
In a group, the atomic size / radius increases as we go down. This is because as we go from
element to another down the group the number of shells increases.

24. In a period as we move from left to right, the atomic size decreases. This is because, within a
period, the electrons are added to the same shell. However, at the same time, protons are being added to
the nucleus and the effect of increasing proton number is greater than that of the increasing electron
number. Thus the nucleus attracts the electrons more strongly, pulling the valence shell closer to it, thus
decreasing the atomic radius/ size.

25. 3. Metallic and non metallic properties:
The metallic character of elements is because of their tendency to lose electrons and form positive
ions ( i.e metals are electropositive in nature).
As we go down a group, the metallic nature of elements increases. This is because the atomic
size increases down a group and so the valence electrons are father away from the nucleus which
decreases the effective nuclear charge on the valence electrons. Thus these can be lost easily.
Across a period from left to right, the effective nuclear charge acting on the valence electrons
increases because of the decreasing atomic size. Thus tendency to lose electrons decreases,
metallic nature decreases.

26. Names of some Groups
Group 1 (IA) - Alkali metals
Group 2 (IIA) - Alkaline earth metals
Groups 13 (IIIA) – Boron family
Group 14 (IVA) – Carbon family
Group 15 (VA) – Nitrogen family
Group 16(VIA) – Chalcogens ( ore forming)
Group 17 (VIIA) – Halogens ( salt forming)
Group 18 (VIIIA) – Inert gases/ Noble gases/ Rare gases

27. The End