Student Resources

Figure 2.11 Orbital Shape Diagrams for the s, p, d, and f Subshells. A pair of electrons with opposite spin can be housed within each of the electron orbitals shown above.
We can summarize the availability of electron orbitals below:

(Back to the Top)

Electron Orbital Filling Rules

Now that we have defined the physical spaces that electrons can occupy, we need to determine the order of electron orbital filling. There are three major rules that we need to follow when filling electron orbitals. Recall that the Pauli Exclusion Principle excludes electrons from having the same quantum numbers, and thus, they cannot exist in the same place at the same time. Second according to the Aufbau principle, electrons orbiting one or more atoms will fill the lowest available energy levels before filling higher energy levels. Thus, the lowest energy levels must be filled with electrons before orbitals with higher energy can house electrons. The final rule that we need to follow is Hund’s rule which states that when electron orbitals have equal energy level, electrons must fill each of those orbitals as single electrons before they can begin to pair with electrons of opposite spin states. This is because electrons are negatively charged and naturally repel each other. Thus, they exist as far apart as possible. This is the most energetically stable way that electrons can fill orbitals.
Due to the strange shapes of the orbitals within the larger subshells, the energy levels of the orbitals don’t always correspond directly with their shell level. For example, the d-subshell orbitals have a higher energy than the s-subshell orbital from next shell level (i.e. the d-subshell orbitals from the 3rd shell have a higher energy and will fill with electrons later that the s-subshell from the 4th shell). Figure 2.12 depicts the energy levels from orbitals in the different subshells.