There is deviation of ionisation enthalpy of some elements from the general trend as shown in figure. The first ionisation enthalpy of B is lower than that of Be and in case of nitrogen, the first ionisation enthalpy is higher than that of $O$. (Also, refer to Q. 27)

(a) Across the period, the nuclear charge increases and the atomic radius decreases. As a result, the tendency of the atom of an element to attract the shared pair of electrons towards itself increases and hence the electronegativity of the element increases. e.g., electronegativity of the elements of the 2nd period increases regularly from left to right as follows $\mathrm{Li}(1.0), \mathrm{Be}(1.5), \mathrm{B}(2.0), \mathrm{C}(2.5), \mathrm{N}(3.0), \mathrm{O}(3.5)$ and .

(b) The ionisation enthalpy decreases regularly as we move from top to bottom, as explained below

(i) On moving down a group from top to bottom, the atomic size increases gradually due to the addition of a new principal energy shell at each succeding element. As a result, the distance between the nucleus and the valence shell increases. In other words, the force of attraction of the nucleus for the valence electrons decreases and hence the ionisation enthalpy should decrease.

(ii) With the addition of new shells, the number of inner shell which shield the valence electrons from the nucleus increases. In other words, the shielding effect 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) Further, in a group from top to bottom nuclear charge increases with increase in atomic number. As a result, the force of attraction of the nucleus for the valence electrons increases and hence the ionisation enthalpy should increase.

The combined effect of the increase in atomic size and screening effect more than compensate the effect of the increased nuclear charge. Consequently, the valence electrons become less and less firmly held by the nucleus and hence the ionisation enthalpy gradually decreases as we move down the group.

As we move from left to right in a period, the number of valence electrons increases by one at each succeeding element but the number of shells remains same. Due to this effective nuclear charge increases. More is the effective nuclear charge, more is the attraction between nuclei and electron.

Hence, the tendency of the element to lose electrons decreases, this results in decrease in metallic character. Furthermore, the tendency of an element to gain electrons increases with increase in effective nuclear charge, so non-metallic character increases on moving from left to right in a period.

When an atom loses an electron to form cation, its radius decreases. In a cation, per electron nuclear forces increases due to decrease in number of electrons. As a result of this, effective nuclear charge increases and the radius of cation decreases. e.g., ionic radius of $\mathrm{Na}^{+}$ is smaller than the radius of its parent atom Na .

\begin{equation} \begin{array}{ll} &\mathrm{Na} \longrightarrow \mathrm{Na}^{+}+1 \mathrm{e}^{-} \\ \text {Electrons } & 11 \qquad \quad10 \\ \text { Nuclear charge } & 11 \qquad \quad 11 \\ \text { lonic size } & 186 \mathrm{pm} \quad 95 \mathrm{pm} \end{array}\end{equation}

On moving down the group, electronegativity decreases because atomic size increases. Fr has the largest size, therefore it is least electronegative.