The pattern of elements in the periodic table reflects the progressive filling of electronic orbitals. The two columns on the left—the alkali metals and alkaline earths—show the addition of 1 and 2 electrons into s‐type subshells. (See Figure 1.)
Figure 1. Filling of the s subshells.
The loss of these s‐subshell valence electrons explains the common +1 and +2 charges on ions of these elements, except for helium, which is chemically inert.
The six elements from boron through neon show the insertion of electrons into the lowest energy p‐type subshell. (See Figure 2.)
Figure 2. Filling of the 2 p subshell.
The same type of subshell is used to describe the electron configurations of elements in the underlying rows. (See Figure 3.)
Figure 3. Filling of the 3 p subshell.
The three long rows of metallic elements in the middle of the periodic table, constituting the rectangle from scandium (21) to mercury (80), are the transition metals. Each of these three rows reflects the filling of a d‐type subshell that holds up to 10 electrons. Figure 4 shows the valence subshell of the first series of transition metals. Notice the general increase in the number of electrons occupying the 3 d subshell.
Figure 4. Filling of the 3 d subshell.
The anomalous electronic configuration of chromium and copper is interpreted as the displacement of 1 electron from an s orbital into a d orbital; these two elements have only one electron in the 4 s subshell because the second electron was promoted into a 3 d subshell. This example warns you that there are exceptions to the general pattern of electronic configurations of the elements. The complicated electronic structure of the transition metals is a consequence of the similar energy of various subshells, like the 4 s and 3 d subshells, which leads to multiple valence states for single elements. Vanadium, for example, shows valences of +2, +3, +4, or +5.
The two rows at the bottom of the periodic table are designated as the lanthanides and actinides, respectively. The lanthanides belong between elements 57 and 72, while the actinides belong between elements 89 and 104. (See Figure 5.)
Figure 5. The correct placement of the lanthanides and actinides in the periodic table.
These two long rows of elements are traditionally moved to the base of the chart so the more important, lighter elements may be closer together for clarity. These two rows of metals each reflect the progressive addition of 14 electrons into an f‐type subshell. The lanthanides occur in only trace amounts in nature and are often called rare earths. All the actinides have large, unstable nuclei that undergo spontaneous radioactive decay. Elements with atomic numbers 93 and higher were synthetically produced.
The outermost electrons of an atom generally determine the chemical behavior of that element. They determine the atom's size, charge, and ability to exchange electrons with other atoms. If you understand how the periodic table displays the pattern of electron configurations, you are on your way to mastering chemistry. You should know that rows of the periodic table show the filling of various subshells and that congeners in columns have similar electron subshells that are filled to the same degree.