6.4 Electronic structure of atoms (electron configurations)  (Page 6/15)

Electron configurations and the periodic table

As described earlier, the periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their electron configurations are added to the table ( [link] ), we also see a periodic recurrence of similar electron configurations in the outer shells of these elements. Because they are in the outer shells of an atom, valence electrons play the most important role in chemical reactions. The outer electrons have the highest energy of the electrons in an atom and are more easily lost or shared than the core electrons. Valence electrons are also the determining factor in some physical properties of the elements.

Elements in any one group (or column) have the same number of valence electrons; the alkali metals lithium and sodium each have only one valence electron, the alkaline earth metals beryllium and magnesium each have two, and the halogens fluorine and chlorine each have seven valence electrons. The similarity in chemical properties among elements of the same group occurs because they have the same number of valence electrons. It is the loss, gain, or sharing of valence electrons that defines how elements react.

It is important to remember that the periodic table was developed on the basis of the chemical behavior of the elements, well before any idea of their atomic structure was available. Now we can understand why the periodic table has the arrangement it has—the arrangement puts elements whose atoms have the same number of valence electrons in the same group. This arrangement is emphasized in [link] , which shows in periodic-table form the electron configuration of the last subshell to be filled by the Aufbau principle. The colored sections of [link] show the three categories of elements classified by the orbitals being filled: main group, transition, and inner transition elements. These classifications determine which orbitals are counted in the valence shell    , or highest energy level orbitals of an atom.

1. Main group elements (sometimes called representative elements ) are those in which the last electron added enters an s or a p orbital in the outermost shell, shown in blue and red in [link] . This category includes all the nonmetallic elements, as well as many metals and the intermediate semimetallic elements. The valence electrons for main group elements are those with the highest n level. For example, gallium (Ga, atomic number 31) has the electron configuration [Ar] 4 s ² 3 d ¹⁰ 4 p ¹ , which contains three valence electrons (underlined). The completely filled d orbitals count as core, not valence, electrons.
2. Transition elements or transition metals . These are metallic elements in which the last electron added enters a d orbital. The valence electrons (those added after the last noble gas configuration) in these elements include the ns and ( n – 1) d electrons. The official IUPAC definition of transition elements specifies those with partially filled d orbitals. Thus, the elements with completely filled orbitals (Zn, Cd, Hg, as well as Cu, Ag, and Au in [link] ) are not technically transition elements. However, the term is frequently used to refer to the entire d block (colored yellow in [link] ), and we will adopt this usage in this textbook.
3. Inner transition elements are metallic elements in which the last electron added occupies an f orbital. They are shown in green in [link] . The valence shells of the inner transition elements consist of the ( n – 2) f, the ( n – 1) d , and the ns subshells. There are two inner transition series:
   1. The lanthanide series: lanthanide (La) through lutetium (Lu)
   2. The actinide series: actinide (Ac) through lawrencium (Lr)