5. Periodic Classification of Elements | Class 10 CBSE | Web Notes | Part 3 | Modern Periodic Table

 5. PERIODIC CLASSIFICATION OF ELEMENTS

MAKING ORDER OUT OF CHAOS – THE MODERN PERIODIC TABLE

Henry Moseley (1913) showed that the atomic number (Z) is a more fundamental property than atomic mass.

Atomic number= Number of protons in an atom’s nucleus.

Thus, Mendeleev’s Periodic Law was modified as follows:

‘Properties of elements are a periodic function of their atomic number.’

Arrangement of elements based on increasing atomic number led to Modern Periodic Table. In this, more precise prediction of properties of elements is possible.

Modern Periodic Table rectified three limitations of Mendeleev’s Periodic Table.

• Positions of Co & Ni resolved based on atomic number.
• Isotopes have same atomic number so they are placed in the same group.
• Atomic number is a whole number. So, there is no confusion about the presence of an element between two elements. E.g. There is no element with atomic number 1.5 placed between hydrogen and helium.

Position of Elements in the Modern Periodic Table

The Modern Periodic Table has 18 vertical columns (groups) and 7 horizontal rows (periods).

Groups signify an identical outer shell electronic configuration. All elements in a group contain same number of valence electrons. The number of shells increases as go down the group.  E.g.

·  Group 1 elements are H, Li, Na, K, Rb, Cs & Fr.

Electronic configuration of H = 1.

Electronic configuration of Li = 2, 1.

Electronic configuration of Na = 2, 8, 1.

Here, all elements have same number of valence electron (i.e., 1).

·  Group 17 elements are fluorine (F), chlorine (Cl) etc. Their outermost shells contain 7 electrons.

There is an anomaly in case of the position of hydrogen. It can be placed in group 1 or 17 in the first period.

• Like group I elements (alkali metals), hydrogen has only one valence electron. Thus, it can lose an electron to achieve a stable configuration like alkali metals. Hence it can be placed in group 1.
• Like group 17 elements, it needs only one electron to complete its valence shell. Thus, it can gain an electron to achieve a noble gas configuration.

Elements in a period do not have the same number of valence electrons, but contain same number of shells. Also, the number of valence shell electrons increases by one unit, as the atomic number increases by one unit on moving from left to right. E.g.

2^nd period elements & their electronic configuration:

Li	Be	B	C	N	O	F	Ne
2,1	2,2	2,3	2,4	2,5	2,6	2,7	2,8

3^rd period elements & their electronic configuration:

Na	Mg	Al	Si	P	S	Cl	Ar
2,8,1	2,8,2	2,8,3	2,8,4	2,8,5	2,8,6	2,8,7	2,8,8

Atoms of different elements with the same number of shells are placed in the same period.

Number of elements in periods are based on how electrons are filled into various shells.

Maximum number of electrons that can be accommodated in a shell depends on the formula 2n² (n= number of the shell). E.g.

• K Shell: 2 × (1)² = 2 electrons. Hence 1^st period has 2 elements. They have only one shell (K).
• L Shell: 2 × (2)² = 8 electrons. Hence 2^nd period has 8 elements. They have 2 shells (K & L).
• M shell: 2 × (3)² = 18 electrons. 3^rd period has 3 shells (K, L & M). Last shell can accommodate only up to 8 electrons. Hence 3^rd period has only 8 elements.
• 4^th, 5^th, 6^th & 7^th periods have 18, 18, 32 & 32 elements respectively.

Mendeleev used formulae of compounds as a basic property to decide the position of an element. This was a good choice because elements are arranged in groups based on the number of valence electrons and valency.  Since valency in a group is same, they will form similar formulae with hydrogen, oxygen etc. Thus they show similar chemical properties.

Trends in the Modern Periodic Table

Valency

It is the number of electrons that must be lost or gained by an atom to attain a stable configuration.

It is determined by the number of valence electrons present in the outermost shell of its atom.

Valency of a metal = Number of valence electrons.

E.g. Electronic configuration of Mg (Z= 12) is 2, 8, 2.

∴ Valency of Mg = 2.

Valency of a non-metal = 8 – No. of valence electrons.

E.g. Electronic configuration of S (Z= 16) is 2, 8, 6.

∴ Valency of S = 8 – 6 = 2.

Valencies of first 20 elements:

Elements	Atomic No.	E. Config.	Valency
H	1	1	1
He	2	2	0
Li	3	2, 1	1
Be	4	2, 2	2
B	5	2, 3	3
C	6	2, 4	8 – 4 = 4
N	7	2, 5	8 – 5 = 3
O	8	2,6	8 – 6 = 2
F	9	2, 7	8 – 7 = 1
Ne	10	2, 8	8 – 8 = 0
Na	11	2, 8, 1	1
Mg	12	2, 8, 2	2
Al	13	2, 8, 3	3
Si	14	2, 8, 4	8 – 4= 4
P	15	2, 8, 5	8 – 5= 3
S	16	2, 8, 6	8 – 6= 2
Cl	17	2, 8, 7	8 – 7= 1
Ar	18	2, 8, 8	8 – 8= 0
K	19	2, 8, 8, 1	1
Ca	20	2, 8, 8, 2	2

In a period, from left to right, valency increases 1 to 4 then decreases from 4 to 0.

When going down a group, valency remains the same.

Atomic size (Atomic radius)

It refers to the radius of an atom. i.e., distance between the centre of the nucleus and outermost shell.

E.g. atomic radius of hydrogen atom is 37 pm (picometre, 1 pm = 10^–12 m).

In a period, atomic radius decreases from left to right. This is due to an increase in nuclear charge which pull the electrons closer to the nucleus reducing atomic size. E.g.

Atomic radii of 2^nd period elements in decreasing order:

Period II elements	Li	Be	B	C	N	O
Atomic radius (pm)	152	111	88	77	74	66

Here, Li has largest atom and O has smallest atom.

Atomic size increases down the group due to the addition of new shells. This increases distance between outermost electrons and nucleus so that the atomic size increases in spite of the increase in nuclear charge. E.g.

Atomic radii of 1^st group elements in an increasing order. Here, Li has smallest atom and Cs has largest atom.

Elements	Atomic radius (pm)
Li	152
Na	186
K	231
Rb	244
Cs	262

Metallic & Non-metallic Properties

In Periodic Table, metals are found on the left side and the non-metals are found on the right side towards the top. A zig-zag line separates metals from non-metals. E.g.

Elements of third period:

Elements with Atomic No.	Configuration	Metal / Non-metal
Na (11)	2, 8, 1	Metal
Mg (12)	2, 8, 2	Metal
Al (13)	2, 8, 3	Metal
Si (14)	2, 8, 4	Metalloid
P (15)	2, 8, 5	Non-Metal
S (16)	2, 8, 6	Non-Metal
Cl (17)	2, 8, 7	Non-Metal
Ar (18)	2, 8, 8	Non-Metal

In the middle, semi-metal or metalloid are found. They show intermediate properties of metals and non-metals. These borderline elements include boron, silicon, germanium, arsenic, antimony, tellurium & polonium.

Metals form bonds by losing electrons. So they are electropositive.

Metallic character decreases across a period and increases down a group because

• Across a period, the effective nuclear charge acting on the valence electrons increases. So, the tendency to lose electrons decreases.
• Down a group, the nuclear charge acting on valence electrons decreases as the outermost electrons are farther away from the nucleus. So, the electrons are lost easily.

Non-metals form bonds by gaining electrons. So, they are electronegative.

In a period, tendency of gaining electrons increases from left to right up to 17^th group. It decreases in 18^th group.

Tendency of gaining electrons decreases down a group.

These trends help to predict the nature of oxides formed by the elements because generally the metal oxides are basic and non-metal oxides are acidic.

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👉 Part 1: Early Attempts of classification
👉 Part 2: Mendeleev's Periodic Table
👉 Part 3: Modern Periodic Table