8.1: Periodic Classification of the Elements

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, a periodic recurrence of similar electron configurations in the outer shells of these elements is observed. 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.

The horizontal rows are known as periods. Across a period, each consecutive element has an additional proton in the nucleus and an additional electron to the valence shell. The vertical columns are groups. Elements in any one group (or column) have the same number of valence electrons (Figure 1); 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. It is the loss, gain, or sharing of valence electrons that defines how elements react. The similarity in chemical properties among elements of the same group occurs because they have the same number of valence electrons.

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 the arrangement of the periodic table is understood; elements whose atoms have the same number of valence electrons are in the same group. The colored sections of Figure 1 show the three categories of elements classified by the orbitals being filled.

Figure 1: This version of the periodic table shows the configuration of each element. Note that down each group, the configuration is often similar.

Main group elements are sometimes called representative elements. These are the elements in which the last electron added enters an s or a p orbital in the outermost shell, shown in blue and red in Figure 1. This category includes all of the nonmetallic elements, as well as the metalloids and many metals. 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]4s²3d¹⁰4p¹, which contains three valence electrons (underlined). The completely filled d orbitals count as core, not valence, electrons.

The two far-left columns comprise the s-block and the six far-right columns constitute the p-block. The noble gases, which are a part of the p-block, all have eight valence electrons except for helium, which has two. These elements are highly stable and unreactive.

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 Figure 1) are not technically transition elements. However, the term is frequently used to refer to the entire d block (colored yellow in Figure 1).

The d-block consists of 10 columns. The principal quantum number of the d orbitals that fill across each row is equal to the row number minus one. In the fourth row, 3d orbitals fill, in the fifth row, 4d orbitals fill, and so on.

Inner Transition Elements: They are shown in green in Figure 1. The valence shells of the inner transition elements consist of the (n – 2)f, the (n – 1)d, and the ns subshells. Inner transition elements constitute f-block as the last electron enters an f orbital. The principal quantum number of the f orbitals that fill across each row is equal to the row number minus two. In the sixth row, the 4f orbitals fill, and in the seventh row, the 5f orbitals fill. There are two inner transition series:

1. The lanthanide series: lanthanide (La) through lutetium (Lu)
2. The actinide series: actinide (Ac) through lawrencium (Lr)

Lanthanum and actinium, because of their similarities to the other members of the series, are included and used to name the series, even though they are transition metals with no f electrons.

This text is adapted from OpenStax Chemistry 2e, Section 6.4: Electronic Structure of Atoms.

The electrons occupying the outermost shell of an atom are valence electrons, while electrons occupying the inner principal energy levels are core electrons. 

A sodium atom with a [Ne]3s¹ electron configuration has three principal energy levels. The complete inner principal energy levels with [He]2s²2p⁶ indicate that there are ten core electrons followed by the third outermost level containing the remaining one electron. Therefore sodium has one valence electron. Similarly, chlorine has ten core electrons and seven valence electrons.

Valence electrons are the farthest from the nucleus and are held most loosely. Hence, they are the easiest to lose or share and play an important role in chemical bonding.

Elements that have the same number of valence electrons exhibit similar chemical properties, as can be seen from the arrangement in the modern periodic table. 

For main-group elements, the lettered group number equals the number of valence electrons and the row number equals the highest principal quantum number of that element.

Each element in a group has the same number of electrons available for bonding. Down the group, the principal quantum number increases by one, whereas the number of valence electrons remains the same. 

The two far-left columns of the periodic table constitute the s-block. For these elements, the last electron enters an s-orbital. Group one elements except for hydrogen are called the alkali metals and are extremely reactive as they have only one valence electron. Group two elements are the alkaline earth metals with two electrons in the valence shell. 

The six far-right columns form the p-block. The valence shell of these elements has completely occupied s-orbitals and the last electron enters the p-orbital. From group three A to group eight A, the number of electrons in p orbitals increases by one. The noble gases have eight valence electrons except for helium, which belongs to the s-block, and only has two electrons.

The d-block consists of the ten columns placed between the s and p-block. These elements are referred to as the transition metals. The last electron enters the d-orbital of the principal shell number one less than the row number. In the fourth row, three d-orbitals fill, in the fifth row, four d-orbitals fill, and so on.

Inner transition elements constitute the f-block and have the last electron entering an f-orbital. The principal quantum number of the f-orbitals that fill across each row is two less than the row number. In the sixth row, the four f-orbitals fill, and in the seventh row, the five f-orbitals fill. Inner transition elements are arranged in the lanthanide series and the actinide series. 

The number of columns in each block indicates how many electrons can be filled in the sublevel of the block. Two  columns in the s-block correlate to an s-orbital with two electrons, six columns in the p-block represent three p-orbitals with six electrons, while ten columns in the d-block correspond to five d-orbitals with two electrons each.

Lastly, the f-block comprises fourteen columns indicating the maximum capacity of seven f-orbitals.