How to find electronic configuration of atoms in an excited state?

Electronic configuration refers to the arrangement of electrons in the atomic orbitals. It helps determine the valence electrons present in an atom which in turn plays a vital role in controlling the chemical behavior of elemental atoms.

But an interesting fact is that the electronic configuration of an atom in the ground state is different from that in the excited state. Why is that and how to find the electronic configuration of an atom in the excited state? Let’s find out through this article.

What is ground state electronic configuration?

The ground state electronic configuration of an atom represents the most stable electronic arrangement at the lowest possible energy. It is the most stable state.

What is excited state electronic configuration?

“An excited state configuration is a higher energy arrangement (it requires energy input to create an excited state). Valence electrons are the electrons utilized for bonding. These are the ‘outer’ electrons (i.e. the electrons which travel furthest away from the nucleus) and the highest energy electrons.”

One or more electrons in a higher energy state can consequently be lost or transferred to a different orbital in order to reduce the total energy of the atomic electrons.

The excited state is an unstable state. So, the atom is most reactive in this state as it wants to revert back to its ground state by losing energy.

Standard notation for writing electronic configuration

The standard notation for writing the electronic configuration of atoms uses:

• Principle quantum number (1,2,3,…) to represent the atomic shell number. For example,1 refers to shell no. 1, 2 denotes shell no. 2, and so on and so forth.
• Symbols (s,p,d,f) for representing subshells. For example, 1s represents the first subshell of shell no. 1 (it has only one subshell), 2s stands for the first subshell of shell no. 2 while 2p denotes the second subshell of shell no.2.
• Numbers (1,2,3,…) in the superscript next to s,p,d,f denotes the number of electrons present in a specific subshell of an atomic element.

Ground state electronic configuration of Oxygen (⁸O)

The Periodic Table of elements tells us that there are a total of 8 electrons present in oxygen in its ground state.

These 8 electrons can be arranged using the standard notation as 1s² 2s² 2p⁴.

While placing electrons in their respective orbitals, we must follow the following three rules:

• Aufbau principle: The electrons are filled in an increasing energy order. The closer a subshell is to the atomic nucleus, the lower its energy. This is because the positively charged nucleus will have a stronger hold on the negatively charged electrons present in that subshell. So, the electrons from the oxygen atom are first placed in 1s then 2s followed by 2p orbitals.

• Hund’s rule: The electrons are first singly filled in the orbitals and then pairing occurs. For example, the p subshell has 3 degenerate atomic orbitals. So, the 4 remaining electrons of oxygen (after filling the lower energy 1s and 2s orbitals), are first placed in each of the 3 orbitals one by one and then the fourth is paired up (see the figure below).
• Pauli exclusion principle: Each orbital can accommodate a maximum of two electrons only and these two electrons will be situated in an anti-parallel spin to minimize electronic repulsion.

The ground state electronic configuration of Oxygen can thus be represented as shown in the figure below.

The electronic configuration shows that the outer shell of oxygen (i.e., shell no 2) has a total of 6 electrons.  So, oxygen has 6 valence electrons respectively.

Now, that we know its ground state electronic configuration, let’s find out how will we write its electronic configuration in the excited state.

Excited state electronic configuration of Oxygen (⁸O)

The electrons jump from the lower energy orbitals to higher energy orbitals by absorbing energy from an external source. As a result, the atoms achieve an excited state. The arrangement of electrons in the atom in this excited state is called the excited state electronic configuration.

As we found out earlier that the ground state electronic configuration of Oxygen is 1s² 2s² 2p⁴. A half-filled p orbital (all the 3 orbitals occupying a single electron only) is more stable than a partially filled p orbital.

The paired electrons in 2p⁴ experience a high electronic repulsive effect. So, on gaining energy, the paired electrons get unpaired.

As a result, one of the two electrons is shifted from 2p to the higher energy 3s orbital. Thus, the electronic configuration of oxygen in the excited state becomes 1s² 2s² 2p³ 3s¹. The total electron count stays the same.

We should note that on excitation the electrons can jump from a low-energy atomic orbital to any of the higher-energy orbitals.

But, as a general rule of thumb, it is usually the least stable electron that undergoes excitation, and it jumps to an energy level such that the energy difference (∆E) between parent and host orbitals corresponds to the energy absorbed from an external source.

Here is another example for you.

Example

The ground state electronic configuration of Carbon (⁶C) is:

1s² 2s² 2p²

One 2p orbital of carbon is still empty so upon excitation, one of the two 2s electrons is unpaired and shifted to the empty 2p orbital. So, the electronic configuration of ⁶C in its excited state becomes:

1s² 2s¹ 2p³         

Why is electronic excitation important?

As the atoms are usually unstable in their excited state, that means they are also the most reactive in this state. So, for example, in CO₂ formation, when an electron jumps from the lower energy 2s to the higher energy 2p orbital, the 2s and one of the 2p orbitals hybridize to form two sp hybrid orbitals of equal energy. These sp hybrid orbitals then form chemical bonds with oxygen.

Electronic excitation followed by de-excitation is also very important for an atom’s identification and quantification using atomic spectroscopy. A specific amount of energy is absorbed and/or released called a quantum/quanta.

The d-block, transition metals display different colors because their electrons are filled in the higher energy orbitals before completely filling inner subshell orbitals. The excitation and de-excitation of these electrons release energy in the visible region of the electromagnetic spectrum.

Different excitation states of an atom

If an electron undergoes a transition from its ground state to the lowest unoccupied orbital then it is known as the first excitation state.

It corresponds to the lowest energy required for an electronic excitation.

If a ground state electron jumps to the second lowest unoccupied orbital then it is called the second excitation state.

The excited electron can further jump to a higher energy level. The farther away it jumps, the greater the amount of energy that needs to be absorbed. These energy differences lie in different regions of the electromagnetic spectrum as per their wavelength and frequency. Energy (E) is related to frequency (f) by the Max Planck equation.

Bohr’s hydrogen model shows different excitation states of a hydrogen atom. The ground state electronic configuration of ¹H is 1s¹. The only valence electron of hydrogen is present in the principle quantum number, n=1.

It can undergo excitation by absorbing energy. If it jumps from n=1 to n=2, it absorbs energy equal to ∆E. This is called the first excitation energy. If it jumps directly from n=1 to n=3, it absorbs energy equal to ∆E’ where ∆E’> ∆E. This is called second excitation energy.

Similarly, the de-excitation of electrons from n=3 to n=1 releases more energy than the energy released by electronic de-excitation from n=3 to n=2.

The transition from a higher energy level to n=1 requires energy in the ultraviolet region (more energetic radiations). This is called the Lyman series. Contrarily, the transition of electrons from a higher energy level to n=2 releases less energetic visible radiations. This is called the Balmer series. Similarly, you may have a look at other transition series from Bohr’s hydrogen model given below.

FAQ

What is the difference between the ground state and excited state electronic configurations?

Ground state The ground state is the lowest energy electronic configuration of an atom. It can be determined by identifying the element from the Periodic Table. The electrons are filled in the atomic orbitals following the Aufbau principle, Hund’s rule, and Pauli exclusion principle respectively. Excited state The atoms achieve an excited state electronic configuration when electrons absorb energy and are shifted from a lower energy orbital to a higher energy atomic orbital. It is usually an unstable electronic configuration. Consequently, the electrons fall back to the ground state by releasing a specific amount of energy.

What is the electronic configuration of nitrogen in the first excited state?

The ground state electronic configuration of nitrogen (7N) is 1s2 2s2 2p3. The 2p orbitals are not completely filled. There is still some space available. So, the first excitation (by minimum energy absorption) makes a 2s electron jump to a 2p orbital. Thus, the electronic configuration of nitrogen in its first excited state is 1s2 2s1 2p4 which is a very unstable configuration.

What is the excited state electronic configuration of sulfur?

The ground state electronic configuration of sulfur (16S) is 1s2 2s2 2p6 3s2 3p4 . Excitation by energy absorption leads to an electron jump from the 3p to the higher energy 3d orbital. So, its electronic configuration becomes 1s2 2s2 2p6 3s2 3p3 3d1. Upon further energy absorption, one of the 3s electrons also shifts to a 3d (the 3p electrons do not jump further because they are very stable in a half-filled 3p3 state). So, the excited state electronic configuration of sulfur becomes 1s2 2s2 2p6 3s1 3p3 3d2

How do you write an excited state electronic configuration for sodium?

The ground state electronic configuration of sodium (11Na) is 1s2 2s2 2p6 3s1. Upon excitation, the 3s electron shifts to its adjacent 3p orbital so the electronic configuration of sodium in this excited state becomes 1s2 2s2 2p6 3s0 3p1. This is a very unstable configuration so the 3p electron readily moves back to 3s, releasing energy equal to 580 nm. This wavelength corresponds to the yellow color of a sodium vapor lamp.

Which of the following electronic configurations represents an excited state for a potassium atom? 1s2 2s2 2p6 3s2 3p5 4s1 1s2 2s2 2p6 3s2 3p5 4s2 1s2 2s2 2p6 3s2 3p6 4s1 1s2 2s2 2p6 3s2 3p6 4s2

The ground state electronic configuration of potassium ( 19K ) is 1s2 2s2 2p6 3s2 3p6 4s1 There is a space for one more electron in the 4s orbital. So, upon excitation, one of the 3p electrons absorbs energy and jumps to the adjacent 4s orbital. Thus, the electronic configuration of a potassium atom in its first excited state is 1s2 2s2 2p6 3s2 3p5 4s1 which is option ii.

Which of the following electronic configuration shows that the atom is in an excited state? 1s2 2s1 1s2 2s2 2p4 1s2 2s2 1s2 2s2 2p6 3s2 1s2 2s2 3s1

Option 5 is the correct answer. There is a 2p orbital before a 3s orbital. If an electron shifts from a 2p to a 3s orbital, the ground state electronic configuration 1s2 2s2 2p1 of Boron (5B) becomes 1s2 2s2 2p0 3s1 or simply 1s2 2s2 3s1.

Summary

• The ground state electronic configuration of an atom refers to the arrangement of electrons in the most stable, least reactive, lowest energy state.
• Upon energy absorption, electrons can jump from a lower energy orbital to a higher energy orbital.
• The arrangement of electrons in this excited state differs from its ground state electronic configuration.
• There is no specific excited state electronic configuration, it depends entirely on the atom involved and the energy difference between the two concerning levels.
• The 1s electron of hydrogen can jump from n=1 to n=2 to n=3, as explained by Bohr’s atomic model.
• First excitation energy is always less than the energy required for subsequent electronic excitations.
• An atom is unstable in its excited state, so the electron falls back to the ground state by spontaneous energy emission.