enthalpy. The first ionization enthalpies of elements having atomic numbers up to are plotted in Fig. . .
The periodicity of the graph is quite striking. You will find maxima at the noble gases which have closed electron shells and very stable electron configurations. On the other hand, minima occur at the alkali metals and their low ionization enthalpies Fig. .
Variation of first ionization enthalpies ( ∆ i H) with atomic number for elements with Z = to can be correlated with their high reactivity. In addition, you will notice two trends the first ionization enthalpy generally increases as we go across a period and decreases as we descend in a group. These trends are illustrated in Figs. .
(a) and . (b) respectively for the elements of the second period and the first group of the periodic table. You will appreciate that the ionization enthalpy and atomic radius are closely related properties. To understand these trends, we have to consider two factors : (i) the attraction of electrons towards the nucleus, and (ii) the repulsion of electrons from each other.
The effective nuclear charge experienced by a Fig. . (a) First ionization enthalpies ( ∆ i H) of elements of the second period as a function of atomic number (Z) and Fig. .
(b) ∆ i H of alkali metals as a function of Z. . (a) . (b) valence electron in an atom will be less than the actual charge on the nucleus because of “shielding” or “screening” of the valence electron from the nucleus by the intervening core electrons.
For example, the s electron in lithium is shielded from the nucleus by the inner core of s electrons. As a result, the valence electron experiences a net positive charge which is less than the actual charge of + . In general, shielding is effective when the orbitals in the inner shells are completely filled. This situation occurs in the case of alkali metals which have single outermost ns -electron preceded by a noble gas electronic configuration.
When we move from lithium to fluorine across the second