Student Resources

Figure 2.13 Electron Filling Diagram for the Element Iron (Fe).  The lefthand diagram shows the partial filling of the electron orbitals during the assembly of an iron atom. Filling begins with the lowest energy orbital, which is the 1s obital. This is followed by the 2s orbital and the three 2p orbitals. Note that all of the 2p orbitals are at the same energy level. Thus, they are filled according to Hund’s Rule: each orbital at the same energy level will fill with one electron alone, until there are no other options. Then the electrons will pair to complete the filling of the 2p orbitals. The figure to the right denotes the completed electron filling diagram for the element iron. Note that the final electrons are placed in the 3d orbitals, and that they are placed to obey Hund’s Rule, leaving 4 of the 6 electrons at this energy level in the unpaired state.
We can write the final electron configuration of iron (Fe) using a notation that denotes each electron shell and subshell and the superscript notation that refers the number of electrons present in the orbitals at that energy level. Remember that s-subshells can house 2 electrons, p-subshells can house 6 electrons, d-subshells can house 10 electrons, and f-subshells can house a maximum of 14 electrons.

From the diagram above, we can notice a few things. First, if we add up the superscript numbers, this should equal the total number of electrons in that atom. For example if we take the sum of the superscript numbers from the diagram above, we get:

2 + 2 + 6 +2 + 6 + 2 + 6 = 26

We can use this as a way to double check and make sure that we have constructed the electron configuration correctly.
The second piece of information that we can find very quickly is the number of electrons present in the outermost shell of the atom. The outermost shell of the atom is referred to as the valence shell, and it contains the most available and reactive electrons within the atom. The valence shell will always correspond to the highest shell number present in the atom that contains electrons. The other electron shells are known as inner electron shells. In the case of iron, the valence shell would be the 4th shell (or period on the periodic table). Shells 1-3 would then be the inner shells. Electrons housed in the valence shell are known as the valence electrons. Valence electrons are important because they are the ones that are most involved in chemical bonding, although sometimes electrons that are close to the valence electrons in energy, may also participate in bond formation. Most often these exceptions involve electrons in the d-subshell.
Note that due to the energy levels of the different subshells, that the only subshells that can exist in the valence shell are the s- and p-subshells. Both the d- and the f-subshells only fill after the s-orbital of the next higher shell level. Thus, they can never contribute to the number of valence electrons. Since the s-orbital can only house 2 electrons, and the p-orbitals can house 6 electrons, the maximum number of valence electrons possible for any atom is 8.

Concept Review

How many electrons are in the valence shell of Phosphorus? Phosphorus has an electron configuration of:

P = 1s22s22p63s23p3

Solution

In the analysis of the electron configuration of phosphorus, we can see that the 3rd shell is the valence shell. For phosphorus there are 2 electrons in the 3s orbital, and there are 3 electrons in the 3p orbitals. This is a total of 5 electrons in the outer shell or 5 valence electrons. Note that due to the energy levels of the different subshells, that the only subshells that can exist in the valence shell are the s- and p-subshells
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Practice using the electron energy diagram to predict the electron configurations of the following atoms:

Boron, B

Carbon, C

Nitrogen, N

Oxygen, O

Fluorine, F

Neon, Ne

Solution:

B: 1s22s22p1
C: 1s22s22p2
N: 1s22s22p3
O: 1s22s22p4
F: 1s22s22p5
Ne: 1s22s22p6

Electron Configurations and the Periodic Table

It is not always convenient to use the energy filling diagram to predict the electron configuration of the elements. Therefore, to be able to predict electron configurations more easily, the periodic table has been arranged to represent the electron orbital filling patterns. This is what gives the periodic table its stair-step or jagged appearance. Once you know how to read the periodic table for electron configurations, it becomes very easy to write out the electron configuration of any element. Figure 2.14 shows a version of the periodic table that depicts the layout of the electron configurations.