Periodic Classification of Elements | Class 10 CBSE | Web Notes

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 the 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 the 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.

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 the same number of valence electrons. The number of shells increases as you 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 the same number of valence electrons (i.e., 1).
Group 17 elements are fluorine (F), chlorine (Cl), etc. Their outermost shells contain 7 electrons.

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

Like group 1 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 the 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:

2^nd Period Elements	Li	Be	B	C	N	O	F	Ne
Electronic Configuration	2,1	2,2	2,3	2,4	2,5	2,6	2,7	2,8

3^rd period elements & their electronic configuration:

3^rd Period Elements	Na	Mg	Al	Si	P	S	Cl	Ar
Electronic Configuration	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 is 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 the same, they will form similar formulae with hydrogen, oxygen, etc. Thus, they show similar chemical properties.

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.

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 from 1 to 4, then decreases from 4 to 0.
When going down a group, valency remains the same.

It refers to the radius of an atom, i.e., the distance between the center of the nucleus and the outermost shell.
E.g., the atomic radius of a hydrogen atom is 37 pm (picometer, 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 pulls the electrons closer to the nucleus, reducing atomic size. E.g.:

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

Here, Li has the largest atom, and O has the smallest atom.
Atomic size increases down the group due to the addition of new shells. This increases the distance between the outermost electrons and the nucleus, so the atomic size increases despite the increase in nuclear charge. E.g.:

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

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, and 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, the tendency to gain electrons increases from left to right up to the 17^th group. It decreases in the 18^th group.
The tendency to gain electrons decreases down a group.
These trends help predict the nature of oxides formed by the elements because generally, metal oxides are basic, and non-metal oxides are acidic.

Select a Topic 👇

Download PDF