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Electronic Configuration

We talked about quantum numbers and atomic orbitals. In this section, we will focus our attention mainly on writing the electronic configuration of atoms and the rules associated with it. First, let's talk about the ground state electronic configuration. The ground state configuration means the electronic configuration at the lowest energy state.
There are certain rules that should be applied to the filling of orbitals. In order to do that properly, we need to know the Aufbau principle. According to the Aufbau principle, the filling order of electrons obeys a general pattern in which the electrons try to occupy the orbitals in such a way as to minimize the total energy; that is, they occupy the lowest energy orbitals first and then step-by-step go to the next available higher energy levels successively. Of course, there are some exceptions to these generalizations. Some students find it hard to remember the order of filling. The diagram given below may help you. So take a close look. Filling order can be depicted as follows:


This is an easy way to remember the general order of electron-filling in the subshells. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 4p, 5s, . . ., and so on.

An s orbital is spherical in shape. All p orbitals have dumbbell shape with two lobes aligned along an axis (see page 47). All d orbitals have slightly complex shapes (see page 48) and are beyond the scope of our discussion. Each orbital can accommodate a maximum of two electrons. Hence, an s subshell, (only one orbital) can accommodate a maximum of two electrons. Similarly, p (three orbitals), d (five orbitals), and f (seven orbitals) subshells can have maximum of six, ten, and fourteen electrons, respectively.

Some possible combinations of quantum numbers for atomic orbitals


Example 3-1

Write the electronic configuration of lithium.



From the periodic table, we can get the atomic number of lithium. The lithium atom has 3 electrons. As we know, 1s subshell is the first one to be filled. The s orbital can hold a maximum of 2 electrons. We have one electron remaining. It will occupy the 2s subshell which is the next energy level. So the third electron will occupy the 2s subshell. Hence, the configuration of lithium is 1s22s1.


Example 3-2

Write the electronic configuration of oxygen in its ground state form.



The oxygen atom contains 8 electrons. The first 2 electrons will go to the 1s level. The next 2 electrons will occupy the 2s level. We have 4 electrons remaining. What is the next subshell according to the filling order? It is 2p. The p subshell can hold a maximum of 6 electrons in its orbitals. So the remaining 4 electrons will occupy the 2p level. Hence, the configuration of oxygen is 1s2 2s2 2p4.

Hund's Rule

W​e have learned the order of filling the subshells. Now let's take a closer look at the filling of electrons in an orbital level. Each orbital can be occupied by a maximum of 2 electrons, and these electrons will have opposite spins as dictated by the spin quantum number. Hund's rule describes the way the electrons fill up the orbitals. According to the Hund's rule, each electron starts filling up each orbital of a given subshell. After all the orbitals in a given subshell have been filled singly (half-filled), then the electrons start pairing. Let's look at some examples.

Example 3-3

Write the electronic configuration of sulfur and also show the filling of electrons with orbital notation.


Sulfur atom has 16 electrons. The electronic configuration is written as 1s2 2s2 2p6 3s2 3p4. To see the significance of the Hund's rule, look at the 3p subshell. In the 3p subshell, we have 3 orbitals.


​Note that the electrons first occupy singly in the orbitals. Altogether there are 4 electrons in the 3p subshell. Instead of filling the orbitals in pairs, the first 3 electrons start filling the three orbitals singly, and then the remaining electron occupies the orbital with the other electron as paired electrons. (See the orbital notation of 3p subshell shown above.) If this were not the case, you would have seen two electron-paired orbitals followed by an empty orbital.
Another idea we want to touch on concerns paramagnetic and diamagnetic substances. Substances that have unpaired electrons are called paramagnetic. Substances that have only paired electrons are called diamagnetic.

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