2.3: Electron Configurations and the Periodic Table

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Learning Objectives

• To correlate the arrangement of atoms in the periodic table results in blocks corresponding to filling of the ns, np, nd, and nf orbitals

Blocks in the Periodic Table

As you have learned, the electron configurations of the elements explain the otherwise peculiar shape of the periodic table. Although the table was originally organized on the basis of physical and chemical similarities between the elements within groups, these similarities are ultimately attributable to orbital energy levels and the Pauli principle, which cause the individual subshells to be filled in a particular order. As a result, the periodic table can be divided into “blocks” corresponding to the type of subshell that is being filled, as illustrated in Figure 2.3.1.

For example, the two columns on the left, known as the s block, consist of elements in which the ns orbitals are being filled. The six columns on the right, elements in which the np orbitals are being filled, constitute the p block. In between are the 10 columns of the d block, elements in which the (n − 1)d orbitals are filled. At the bottom lie the 14 columns of the f block, elements in which the (n − 2)f orbitals are filled.

Because two electrons can be accommodated per orbital, the number of columns in each block is the same as the maximum electron capacity of the subshell: 2 for ns, 6 for np, 10 for (n − 1)d, and 14 for (n − 2)f. Within each column, each element has the same valence electron configuration—for example, ns1 (group 1) or ns2np1 (group 13). As you will see, this is reflected in important similarities in the chemical reactivity and the bonding for the elements in each column.

Note

Because each orbital can have a maximum of 2 electrons, there are 2 columns in the s block, 6 columns in the p block, 10 columns in the d block, and 14 columns in the f block.

Hydrogen and helium are placed somewhat arbitrarily. Although hydrogen is not an alkali metal, its 1s1 electron configuration suggests a similarity to lithium ([He]2s1) and the other elements in the first column. Although helium, with a filled ns subshell, should be similar chemically to other elements with an ns2 electron configuration, the closed principal shell dominates its chemistry, justifying its placement above neon on the right.

Example 2.3.1

Use the periodic table to predict the valence electron configuration of all the elements of group 2 (beryllium, magnesium, calcium, strontium, barium, and radium).

Given: series of elements

Strategy:

1. Identify the block in the periodic table to which the group 2 elements belong. Locate the nearest noble gas preceding each element and identify the principal quantum number of the valence shell of each element.
2. Write the valence electron configuration of each element by first indicating the filled inner shells using the symbol for the nearest preceding noble gas and then listing the principal quantum number of its valence shell, its valence orbitals, and the number of valence electrons in each orbital as superscripts.

Solution:

A The group 2 elements are in the s block of the periodic table, and as group 2 elements, they all have two valence electrons. Beginning with beryllium, we see that its nearest preceding noble gas is helium and that the principal quantum number of its valence shell is n = 2.

B Thus beryllium has an [He] 2s2 electron configuration. The next element down, magnesium, is expected to have exactly the same arrangement of electrons in the n = 3 principal shell: [Ne] 3s2. By extrapolation, we expect all the group 2 elements to have an ns2 electron configuration.

Exercise 2.3.1

Use the periodic table to predict the characteristic valence electron configuration of the halogens in group 17.

All have an ns2 np5 electron configuration, one electron short of a noble gas electron configuration. (Note that the heavier halogens also have filled (n − 1)d10 subshells, as well as an (n − 2)f14 subshell for Rn; these do not, however, affect their chemistry in any significant way.

Electron Configuration of Ions

How does the electron configuration of an atom change when it gains or loses an electron to become an ion? The answer to this important question depends on the location of the atom in the Periodic Table. For the representative elements (s- and p-blocks) the electrons are always removed in the opposite order that they were filled. For example, the last electron added in the electron configuration of magnesium is the second 3s-electron. Mg+ would have a configuration of 1s2 2s2 2p6 3s1. For Mg2+, the electron configuration would be 1s2 2s2 2p6. There is an exception to this process for transition metals (members of the d-block) and the lanthanoids and actinoids (members of the f-block). For these elements, the electrons are always removed from the outermost s-subshell first, then the d- or f-electrons.

 Sc [Ar] 4s2 3d1 Sc+ [Ar] 4s1 3d1 Sc2+ [Ar] 3d1 Sc3+ [Ne] 3s2 3p6

When adding an electron, place it into the next available orbital, following Hund's rule, Pauli's Exclusion Principle, and the Aufbau Principle. For example, when fluorine gains an electron, the electron configuration of F- will be 1s2 2s2 2p6.

Summary

The arrangement of atoms in the periodic table results in blocks corresponding to filling of the ns, np, nd, and nf orbitals to produce the distinctive chemical properties of the elements in the s block, p block, d block, and f block, respectively. Electrons can be added or removed from an atom to generate cations and anions, and the electron configurations of these ions can be predicted.

2.3: Electron Configurations and the Periodic Table is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.