Metal valence ionization energy

Z the atomic number and I is an appropriate valence ionization energy (in most.cases the first ionization energy). E is the cohesive energy and E + CZ) is the solution energy of impurity atom (Z+1) in (Z) metal. 

The low ionization energies of elements at the lower left of the periodic table account for their metallic character. A block of metal consists of a collection of cations of the element surrounded by a sea of valence electrons that the atoms have lost (Fig. 1.53). Only elements with low ionization energies—the members of the s block, the d block, the f block, and the lower left of the p block—can form metallic solids, because only they can lose electrons easily. 

The valence electron configuration of the atoms of the Group 2 elements is ns1. The second ionization energy is low enough to be recovered from the lattice enthalpy (Fig. 14.18). Flence, the Group 2 elements occur with an oxidation number of +2, as the cation M2+, in all their compounds. Apart from a tendency toward nonmetallic character in beryllium, the elements have all the chemical characteristics of metals, such as forming basic oxides and hydroxides. 

The alkali metals in Group 1(a) have the lowest ionization energies, which is again expected since they always form cations with a +1 valence. There is little variation in I across the d-block and f-block elements, with a slight increase in / as the atomic number increases. 

Metal atoms have a small number of valence electrons. The nuclear attractive forces between the metal nuclei and their valence electrons are reduced by the inner electrons (which are closer to nucleus). Thus, the nucleus of a metal atom exerts only a small attractive force on its valence electrons and these electrons are able to move more freely. For this reason, metal atoms have very low ionization energies and electronegativities. 

A) Alkali metals have one electron in their outer shell, which is loosely bound. This gives them the largest atomic radii of the elements in their respective periods. Their low ionization energies result in their metallic properties and high reactivities. An alkali metal can easily lose its valence electron to form the univalent cation. Alkali metals have low electronegativities. They react readily with nonmetals, particularly halogens. 

In a) and P) the non bonding-hypothesis for 5 f electrons is retained, differences in cohesive energy being only due to promotion of outer electrons from one to another orbital state and ionization energies (or electron affinities) due to the different valence states attained. Therefore, any further discrepancy found with experimental values, is indicative of the metallic bonding introduced by delocalization of the 5f electrons (point y). 

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