Part 1 and 2 are not very clearly defined in the question.
The electronic configuration of four metals, sodium, magnesium, aluminum and silicon are given in the image. From the electronic configuration of the given metals, it is clear that as we move from left to right in a periodic table (that is, across a period), one electron is added between any two successive elements. This extra electron is progressively added in next available electron orbital. Thus, the electronic configuration of any element can be written by simply adding 1 electron to the next available orbital in the electronic configuration of the element preceding the given element in the periodic table. Historically, elements with similar electronic configurations were clubbed together in groups. Each group has the same number of electrons in the outermost shell. As we move down a group, each element has 1 extra shell, as compared to preceding element. Each element in the group will have same number of valence electrons, for example, Sodium is part of alkali metals and each element has only 1 electron in outermost orbital, which can be lost to form a cation. Elements in a group have similar chemical properties and hence were clubbed together.
The elements that have fully filled orbitals are most stable, since they do not need to gain or lose electrons. Such elements, with fully filled outermost orbitals are known as noble gases. These elements are very stable and practically inert, since they do not react. All the noble gases are monoatomic gases at standard conditions.
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