CBSE Class 11 Chemistry Periodic Classification Of Elements Notes

SUMMERY

CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES

Dobereiner’s classification: In triads, the atomic mass of the middle element is approximately the average of the other two elements. This is known as Law of Triads.

Newlands classification:

Newland arranged elements in the increasing order of their atomic masses. He noted that the properties of every eighth element, starting from a given element, is similar to that of the first element. Mendeleev’s classification: Dimitri Mendeleev classified the elements in the increasing order of their atomic weights. He founded that the properties of elements repeat after a regular interval. Based on this observation, he proposed a periodic law which states that “The properties of elements are the periodic functions of their atomic weights.” That is, when elements are arranged in the increasing order of their atomic weights, their properties repeat after regular interval.

Modern Periodic table

Henry Moseley’s work on the atomic spectra of elements proved that atomic number is a more Fundamental property than atomic mass. Based on this serration, he modified the Mendeleev’s periodic law as “the physical and chemical properties of elements are the periodic functions of their atomic numbers”. This is known as Modern Periodic law. Based on modern periodic law, numerous forms of periodic tables have been proposed. The most Commonly used is the long form of periodic table. Classification of Elements & Periodicity in Properties In this periodic table, the elements are arranged in the increasing order of their atomic number. It Contains 7 horizontal rows called periods and 18 vertical columns called groups. Elements having similar outer electronic configurations are arranged in same group or family. The groups are numbered from 1 to 18. Due to the similar outer electronic configuration and same valency, the elements present in the same group have similar properties.

Classification of Elements & Periodicity in Properties

Atomic Radius

It is defined as the distance from the center of the nucleus to the outermost shell having electrons. The atomic size decreases from left to right in a period. This is because in a period, the electrons are Added to the same valence shell. Thus the number of shells remains same, but the effective nuclear charge increases. So the atomic radius decreases. In a given period, alkali metals (group 1) have the maximum size and halogens (group 17) have the minimum size. Down a group, the atomic radius increases from top to bottom. This is because of the increase in no. of shells and shielding effect. (in atoms with higher atomic number, the inner electrons partially shield the attractive force of the nucleus. So the outer electrons do not experience the full attraction of the nucleus and this is known as shielding effect or screening effect).Atomic radius of noble gases is larger than that of halogens. This is because noble gases are monoatomic So van der Waal’s radius is used to express the atomic radius which is greater than covalent radius or metallic radius.

Ionic radius

It is defined as the half of the inter nuclear distance between cations and anions of an ionic crystal. The variation of ionic radius is same as that of atomic radius. Generally a cation is smaller than its parent atom (e.g. Na+ is smaller than Na atom). This is because a cation has fewer electrons, but its nuclear charge remains the same as that of the parent atom. An anion is larger than its parent atom (e.g. Cl- is larger than Cl atom). This is because the addition of one or more electrons would result in an increased electronic repulsion and a decrease in effective nuclear charge.

Isoelectronic species:

Atoms and ions having the same number of electrons are called isoelectronic species. E.g. O2-, F-, Ne, Na+, Mg2+ etc. (All these contain 10 ectrons).Among isoelectronic species, the cation with greater positive charge will have the smaller radius. This is because of the greater attraction of electrons to the nucleus. The anion with greater negative charge will have the larger radius. Here the repulsion between electrons is greater than the attraction of the nucleus. So the ion will expand in size.

Ionisation enthalpy (ΔiH)
It is defined as the energy required to remove an electron from the outer most shell of an isolated gaseous atom in its ground state. It may be represented as: X(g) + ΔiH → X+ (g) + e–.Its unit is kJ/mol or J/mol.

Factors affecting ionization enthalpy

The important factors which affect ionization enthalpy are:
a) Atomic size: Greater the atomic size (atomic radius), smaller will be the ionization enthalpy.
b) Nuclear charge: The value of ionization enthalpy increases with nuclear charge.
c) Shielding effect: As the shielding effect increases, the electrons can easily be removed and so the Ionisation enthalpy decreases.
d) Presence of half-filled or completely filled orbitals increases ionization enthalpy. 

Variation of ΔiH along a period and a group

Along a period, ionisation enthalpy increases from left to right. This is because of the decrease in atomic radius and increase in nuclear charge. Thus alkali metals have the least ΔiH and noble gases have the most. Down a group, ΔiH decreases due to increase in atomic radius and shielding effect. Thus among alkali metals, lithium has the least ΔiH and francium has the most. In the second period of modern periodic table, the first ionisation enthalpy of Boron is slightly less than that of Beryllium. This is because of the completely filled orbitals in Be (1s² 2s²). Similarly the first ionisation enthalpy of N is greater than that of Oxygen. This is because N has half filled electronic configuration (1s² 2s² 2p³), which is more stable and so more energy is required to remove an electron.

Electron gain enthalpy (ΔegH)

It is the heat change (enthalpy change) when an electron is added to the outer most shell of an isolated
gaseous atom. It can be represented as X(g) + e–→ X-(g).
positive electron gain enthalpy because of their completely filled (stable) electronic configuration. Electron gain enthalpy also depends on atomic size, nuclear charge, shielding effect etc. As the atomic
size increases ΔegH decreases. When nuclear charge increases, electron gain enthalpy increases and become more negative. Shielding effect decreases ΔegH. Presence of half-filled or completely filled orbitals makes ΔegH less negative.

Periodic variation of ΔegH

From left to right across a period, ΔegH become more negative. This is because of decrease in atomic Radius and increase in nuclear charge. So the ease of addition of electron increases and hence the ΔegH. Down a group, ΔegH becomes less negative. This is due to increase in atomic radius and shielding effect.
Electron gain enthalpy of fluorine is less negative than chlorine. This is because, when an electron is added to F,it enters into the smaller 2nd shell. Due to e smaller size, the electron suffers more repulsion from the other electrons. But for Cl, the incoming electron goes to the larger 3rd shell. So the electronic repulsion is low and hence Cl adds electron more easily than F. Due to the same reason ΔegH of Oxygen is less negative than S. Thus in modern periodic table, alkali metals have the least –ve ΔegH and halogens have the most –ve ΔegH. Among halogens, the negative ΔegH decreases as follows.
Cl> F > Br > I.The negative electron gain enthalpy is also called electron affinity.

Electronegativity

Electronegativity of an atom in a compound is the ability of the atom to attract shared pair of electron of Electrons. It is not a measurable quantity and so it has no unit. There are different scales for measuring the Electronegativity depends on atomic size and nuclear charge. As the atomic radius increases,Electronegativity decreases. Greater the nuclear charge, greater will be the electronegativity. Generally Electronegativity increases across a period and decreases along a group. So in modern periodic table, F has the maximum electronegativity and Fr has the minimum electronegativity.

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