Name: electron filling order in atomic orbitals
Uploaded: 2020-09-03T13:38:28.000Z
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K12 Standards

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NC K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.1.3

The Aufbau Principle and Hund's Rule

Recall that atomic orbitals have different energies and can accommodate two electrons each. The aufbau principle dictates the distribution of electrons among the subshells of the atom, and Hund&#8217;s rule of maximum multiplicity explains the filling of orbitals in the subshells.

The aufbau principle states that in the ground state, electrons fill atomic orbitals from lowest to highest energy to achieve the lowest-energy configuration. 

Though energy generally increases with shell number, the greater penetration of s orbitals often leads to the four-s and five-s orbitals having lower energies than the three-d and four-d orbitals, respectively. The order can be remembered using diagrams like this one, where the path of the arrow reveals the sequence in which electrons are assigned to orbitals.

Consider writing the electron configuration for carbon &#8212; an element with atomic number six. Certainly, the 1s orbital, which has the lowest energy, should be filled before the 2s orbital. Each orbital can hold a maximum of two electrons.

The fifth electron enters a 2p subshell. But which of the three 2p orbitals does it populate? 

Well! The orbitals in any subshell are presumed to be degenerate, which means that they have the same energy. Thus, the fifth electron can enter any of the three degenerate 2p orbitals. 

What about the sixth electron? Does it enter the 2p orbital with an electron or one of the vacant 2p orbitals?

According to Hund&#8217;s rule of maximum multiplicity, electrons singly occupy all orbitals of a given energy level before they start pairing, and the unpaired electrons cannot have opposite spins: their spins must be parallel.

Hence, for carbon, the two 2p electrons must occupy two different orbitals and have parallel spins. This way, electrons can spread over a larger area. This decreases their shielding of each other, thereby minimizing the energy of the atom.

For nitrogen, each of the three 2p orbitals is singly occupied.

For oxygen, once the degenerate 2p orbitals are filled singly, the last electron must pair with another 2p electron. The atom has two unpaired electrons.

The electron configuration of neon reveals that the outermost shell is filled to its maximum capacity of eight electrons. Neon has two electrons, called core electrons in the inner shell, and eight electrons, called valence electrons, in the outermost shell.

the aufbau principle and hund's rule

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and continuing across the periods of the periodic table, we add one proton at a time to the nucleus and one electron to the proper subshell until we have described the electron configurations of all the elements. This procedure is called the aufbau principle, from the German word aufbau (&#8220;to build up&#8221;). Each added electron occupies the subshell of lowest energy available, subject to the limitations imposed by the allowed quantum numbers according to the Pauli exclusion principle. Electrons enter higher-energy subshells only after lower-energy subshells have been filled to capacity. Figure 1 illustrates the traditional way to remember the filling order for atomic orbitals. 

<img alt="Electron configuration diagram; Aufbau principle; orbital energy levels: 1s-6d with arrows." src="/files/core/CDNSource/lessons/986/986Image_1.png" style="width: 700px;" data-altupdatedat="1754904545000" data-alt="Electron configuration diagram; Aufbau principle; orbital filling order with arrows." />

Figure 1 This diagram depicts the energy order for atomic orbitals and is useful for deriving ground-state electron configurations.

Consider writing the electron configuration for carbon&#8212;an element with atomic number six. Four electrons fill the 1s and 2s orbitals. The remaining two electrons occupy the 2p subshell. We now have a choice of filling one of the 2p orbitals and pairing the electrons or of leaving the electrons unpaired in two different, but degenerate, p orbitals. The orbitals are filled as described by Hund&#8217;s rule: the lowest-energy configuration for an atom with electrons within a set of degenerate orbitals is that having the maximum number of unpaired electrons. Thus, the two electrons in the carbon 2p orbitals have identical n, l, and ms quantum numbers and differ in their ml quantum number (in accord with the Pauli exclusion principle). The orbital diagram for carbon, with an electron configuration of 1s22s21p2 is: 

<img alt="Electron configuration diagram, 1s&#178; 2s&#178; 2p&#185;, illustrating orbital filling and electron spin." src="/files/core/CDNSource/lessons/986/986Image_2.png" style="width: 300px;" data-altupdatedat="1754904545000" data-alt="Electron configuration diagram for lithium; orbitals 1s&#178;, 2s&#178;, 2p&#185;; quantum shells." />

Nitrogen (atomic number 7) fills the 1s and 2s subshells and has one electron in each of the three 2p orbitals, in accordance with Hund&#8217;s rule. These three electrons have unpaired spins. Oxygen (atomic number 8) has a pair of electrons in any one of the 2p orbitals (the electrons have opposite spins) and a single electron in each of the other two. Fluorine (atomic number 9) has only one 2p orbital containing an unpaired electron. All of the electrons in the noble gas neon (atomic number 10) are paired, and all of the orbitals in the n = 1 and the n = 2 shells are filled. 

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/6-4-electronic-structure-of-atoms-electron-configurations">Openstax, Chemistry 2e, Section 6.4: Electronic Structure of Atoms.</a>

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

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JoVE Core

Electronic Structure of Atoms

KEY TERMS AND CONCEPTS

electron filling order in atomic orbitals

Electron filling order in atomic orbitals follows the aufbau principle and Hund's rule. To write an electron configuration, we build atoms in order of atomic number. Starting with hydrogen and moving across the periods of the periodic table, one proton is added to the nucleus and one electron is placed into the correct subshell. This process describes the ground-state electron configurations of the elements.

The aufbau principle comes from the German word aufbau, which means &quot;to build up.&quot; It says that each added electron enters the lowest-energy subshell available. The Pauli exclusion principle still applies, so electrons must follow the allowed quantum numbers. Higher-energy subshells do not fill until lower-energy subshells reach capacity.

A common filling-order diagram helps show the energy sequence of the atomic orbitals. Carbon is a useful example because it has atomic number 6. Four electrons fill the 1s and 2s orbitals, and the remaining two electrons go into the 2p subshell. At that point, the electrons can either pair in one 2p orbital or stay unpaired in two different p orbitals.

Hund's rule gives the lower-energy choice for degenerate orbitals, which are orbitals with the same energy. The rule states that electrons in a set of degenerate orbitals occupy the arrangement with the maximum number of unpaired electrons. In carbon, the two 2p electrons have the same n, l, and m_s quantum numbers, and they differ in m_l, which matches the Pauli exclusion principle.

Nitrogen, with atomic number 7, fills the 1s and 2s subshells and places one electron in each of the three 2p orbitals. These three electrons are unpaired. Oxygen, atomic number 8, has one pair of electrons in one 2p orbital and one single electron in each of the other two. Fluorine, atomic number 9, has only one 2p orbital with an unpaired electron.

Neon, atomic number 10, has all paired electrons. Its n = 1 and n = 2 shells are completely filled. This pattern shows how electron configurations build step by step across the first ten elements.

Electron Filling Order in Atomic Orbitals

electron-filling-order-in-atomic-orbitals

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Electron filling order in atomic orbitals follows the aufbau principle and Hund's rule. To write an electron configuration, we build atoms in order of atomic number. Starting with hydrogen and moving across the periods of the periodic table, one proton is added to the nucleus and one electron is placed into the correct subshell. This process describes the ground-state electron configurations of the elements. The aufbau principle comes from the German word aufbau, which means "to build up." It...

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key terms and concepts

Learn how electron configurations build by atomic number, why lower-energy orbitals fill first, and how Hund's rule shapes carbon, nitrogen, oxygen, fluorine, and neon.

Electron Filling Order in Atomic Orbitals | JoVE K12

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Electron Filling Order in Atomic Orbitals

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