CBSE Class 11 Chemistry Classification Of Elements And Periodicity In Properties Of Elements Notes Set 02

3.Classification of elements and periodicity in properties

Some Important Points and Terms of the Chapter

1.  Dobereiner's Triads:In 1817 a German chemist Doberneiner identified certain groups of three  elements. These groups of three elements having similar properties was called triads. When three elements were arranged in order of their increasing atomic masses, the atomic mass of the middle element was roughly the mean of the atomic masses of the other two element 

2.  New Lands Law of octaves:When elements were arranged in order of their increa sing relative atomic  masses.  The  properties  of  every  eight  elements  were  similar  to  the  first  one,  like  the eighth  note  of  a  musical  scale.This  repetition  in  the  properties  of  elements  is  just  like  the repetition of eighth node in an octave of music.

3.  Mendeleev's Periodic Law:The physical and chemical properties of elements are the periodic function of their atomic masses.

4.  Mendeleev's Periodic Table:When mendeleev started his work, 63 elements were known at that time. He selected hydrogen and oxygen as they  are very reactive and formed compounds with most  elements.  Mendeleev's  periodic  table  contains  vertical  columns  called  groups  and horizontal rows called periods. There were 7 periods and 8 groups. Noble gases were not known at that time. So there was no group of noble gases.The elements in each group of the periodic tables are similar to one another in many properties. The similar properties of the elements are repeated periodically

(a).Merits of mendeleev's classification

1.   Mendeleev's  periodic  law  predicted  the  existence  of  some  elements  that  had  not  been discovered at that time
2.   Could  predict  the  properties  of  several  elements  on  the  basis  of  their  position  in  the periodic table.
3.   Could accommodate noble gases when they were discovered.

(b)Limitations of mendeleev's classification :-

1.   The correct position could not be assigned to the hydrogen in the periodic table.
2.   Wrong order of the atomic masses of some elements could not be explained.
3.   The position of isotopes could not be explained.
4.   Uncertainty in prediction of new elements was there.

5.  Modern periodic law: Properties of elements are the periodic function of their atomic number.

6.  Modern  Periodic  Table:  This  table  was  prepared  was  Bohr  and  is  based  upon  the  electronic configuration  of  elements.  The  table  consists  of  18  vertical  columns  called  groups  Elements having  similar  outer  electronic  configurations  in  their  atoms  are  arranged  in  vertical  columns, referred  to  as  groups  .  According  to  the  recommendation  of  International  Union  of  Pure  and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18 and  the  table consists of 7 horizontal  rows  called  periods.  The  first  period  contains  2  elements.  The  subsequent  periods consists of 8, 8, 18, 18 and 32 elements, respectively. The seventh period is incomplete and like the sixth period  would  have a theoretical  maximum  (on the  basis of  quantum numbers) of 32 elements.  In  this  form  of  the  Periodic  Table,  14 elements  of  both  sixth  and  seventh  periods (lanthanoids and actinoids, respectively) are placed in separate panels at the bottom.

8. We can classify the elements into four blocks viz., s-block, p-block, d-block and f-block depending on the type of atomic orbital that are being filled with electrons.

9. s-Block Elements :The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1and ns2 outermost electronic configuration belong to the s-Block Elements.

10. p-Block Elements The p-Block Elements comprise those belonging to Group 13 to 18 and these together with the s-Block Elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns2np1 to ns2np6 in each period.

11. d-Block Elements These are the elements of Group 3 to 12 in the centre of the Periodic Table.
These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n-1)d1- 10ns0-2 .

12. f-Block Elements The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) – Lr (Z = 103) are
characterised by the outer electronic configuration (n-2)f1-14 (n-1)d0–1ns2. The last electron added to each element is filled in f- orbital. These two series of elements are hence called the Inner- Transition Elements (f-Block Elements).

13. Variation in Atomic Radius in Period: The atomic size generally decreases across a period It is because within the period the outer electrons are in the same valence shell and the effective nuclear charge increases as the atomic number increases resulting in the increased attraction of electrons to the nucleus.

14. Variation in Atomic Radius in Group: Within a family or vertical column of the periodic table, the atomic radius increases regularly with atomic number as). as we descend the groups, the principal quantum number (n) increases and the valence electrons are farther from the nucleus.
This happens because the inner energy levels are filled with electrons, which serve to shield the outer electrons from the pull of the nucleus. Consequently the size of the atom increases as reflected in the atomic radii.

15. The atomic radii of noble gases are not considered here. Being monatomic, their (non-bonded radii) values are very large. In fact radii of noble gases should be compared not with the covalent radii but with the van der Waals radii of other elements.

16. A cation is smaller than its parent atom because it has fewer electrons while its nuclear charge remains the same. The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge. For example, the ionic radius of fluoride ion (F– ) is 136 pm whereas the atomic radius of fluorine is only 64 pm. On the other hand, the atomic radius of sodium is 186 pm compared to the ionic radius of 95 pm for Na+.

17. Isoelectronic species :Atoms and ions which contain the same number of electrons.. For example, O2–, F–, Na+ and Mg2+ have the same number ofelectrons (10). Their radii would be different because of their different nuclear charges. The cation with the greater positive charge will have a smaller radius because of the greater attraction of the electrons to the nucleus. Anion with the greater negative charge will have the larger radius. In this case, the net repulsion of the electrons will outweigh the nuclear charge and the ion will expand in size.

18. Ionization Enthalpy: It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state. In other words, the first ionization enthalpy for an element X is the enthalpy change (ΔiH) for the reaction depicted in equation. X (g) X+(g) + e– . The ionization enthalpy is expressed in units of kJ mol–1. We can define the second ionization enthalpy as the energy required to remove the second most loosely bound electron; it is the energy required to carry out the reaction shown in equation X+(g)X2+(g) + e– . Energy is always required to remove electrons from an atom and hence ionization enthalpies are always positive. The second ionization enthalpy will be higher than the first ionization enthalpy because it is more difficult to remove an electron from a positively charged ion than from a neutral atom.

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