The main drawbacks of Mendeleef's periodic table are

(i) Some elements having similar properties were placed in different groups whereas some elements having dissimilar properties were placed in the same group.

e.g., alkali metals such as $\mathrm{Li}, \mathrm{Na}, \mathrm{K}$, etc., (IA group) are grouped together with coinage metals such as $\mathrm{Cu}, \mathrm{Ag}, \mathrm{Au}$ (IB group) though their properties are quite different. Chemically similar elements such as $\mathrm{Cu}(\mathrm{IB}$ group) and Hg (IIB group) have been placed in different groups.

(ii) Some elements with higher atomic weights are placed before the elements with lower atomic weights in order to maintain the similar chemical nature of elements.

i.e., $\quad { }_{18}^{39.9} \mathrm{Ar}$ and ${ }_{19}^{39.1} \mathrm{~K} ;{ }_{27}^{58.9} \mathrm{Co}$ and ${ }_{28}^{58.7} \mathrm{Ni}$, etc.

(iii) Isotopes did not find any place in the periodic table. However, according to Mendeleef's classification, these should be placed at different places in the periodic table. (All the above three defects were however removed when modern periodic law based on atomic number was given).

(iv) Position of hydrogen in the periodic table is not fixed but is controversial.

(v) Position of elements of group VIII could not be made clear which have been arranged in three triads without any justification.

(vi) It could not explain the even and odd series in IV, V and VI long periods.

(vii) Lanthanides and actinides which were discovered later on have not been given proper positions in the main frame of periodic table.

The long form of the periodic table is better than Mendeleef's periodic table because it classifies the elements on the basis of electronic configurations of their atoms. The characteristics of this table are

(i) The table consists of 9 vertical columns, called the groups and 7 horizontal rows, called the periods.

(ii) The groups are marked 0 to VIII out of which group I to VII are subdivided into subgroups $A$ and $B$.

(iii) The group IA elements ( $\mathrm{Li}, \mathrm{Na}, \mathrm{K}, \mathrm{Rb}, \mathrm{Cs}$ and Fr ) are known as alkali metals and the group IIA elements ( $\mathrm{Be}, \mathrm{Mg}, \mathrm{Ca}, \mathrm{Sr}, \mathrm{Ba}$ and Ra ) are known as alkaline earth metals. Elements in group VIII A ( $\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}$ and At ) are called halogens and elements in group VIII ( $\mathrm{He}, \mathrm{Ne}, \mathrm{Ar}, \mathrm{Kr}$, Xe and Rn ) are called noble gases or rare gases.

(iv) The group VIII has there similar elements placed together in one place. These are called transition triads, e.g., $\mathrm{Fe}, \mathrm{Co}$ and $\mathrm{Ni}, \mathrm{Ru}, \mathrm{Rh}$ and $\mathrm{Pd} ; \mathrm{Os}, \mathrm{Ir}$ and Pt etc.

(v) In the 6th and 7th period, 14 elements present called as lanthanides and actinides respectively.

(vi) Based on their electronic configuration, elements have been grouped into $s-, p-, d-$ and $f$-blocks. This has helped us to understand their properties more easily.

(vii) There is gradual change in properties seen from one end to the other.

The ionisation enthalpies decreases regularly as we move down a group from one element to the other. This is evident from the values of the first ionisation enthalpies of the elements of group 1 (alkali metals) and group 17 elements as given in table and figure.

Given trend can be easily explained on the basis of increasing atomic size and screening effect as follows

(i) On moving down the group, the atomic size increases gradually due to the addition of one new principal energy shell at each succeeding element. Hence, the distance of the valence electrons from the nucleus increases.

Consequently, the force of attraction by the nucleus for the valence electrons decreases and hence the ionisation enthalpy should decrease.

(ii) With the addition of new shells, the shielding or the screening effect increases. As a result, the force of attraction of the nucleus for the valence electrons further decreases and hence the ionisation enthalpy should decrease.

(iii) Nuclear charge increases with increase in atomic number. As a result, the force of attraction by the nucleus for the valence electrons should increase and accordingly the ionisation enthalpy should increase.

The combined effect of the increase in the atomic size and the screening effect more than compensates the effect of the increased nuclear charges. Consequently, the valence electrons become less and less firmly held by the nucleus and hence the ionisation enthalpies gradually decrease as move down the group.