How to Write Electronic Configuration Using Aufbau Principle Hund Rule and Pauli Exclusion Principle
Have you ever considered how atoms are represented? The number of orbitals for atoms with a high atomic number is also high. For instance, the atomic number of mercury (Hg) is 80. For example, mercury (Hg) has an atomic number of 80. The filling of electrons is somewhat complicated as it has 80 electrons to fill around the nucleus. We must deal with a few electronic configuration prediction rules in order to specify the position of an electron in an atom.
Electronic configuration is defined as the layered arrangement of shells, where a specific number of electrons orbit around the nucleus. An atom has a central core where the entire positive charge concentrates; this is the nucleus. The nucleus contains neutrons and protons huddled closely right at the centre of the atom. Atomic shells have a layered arrangement. Each shell has a fixed number of electrons revolving around the nucleus.
A scientist named Charles Bury gave a set of simple rules that made it very easy to understand how electrons are arranged in Bohr's atomic model. This scheme of electron arrangement is known as the Bohr-Bury electron arrangement.
Rules for Filling Electrons
Rules for filling electrons based on Bohr-Bury electron arrangement are given as follows:
An electron will always fill an orbit or a shell with lower energy and then occupy the higher energy ones. Therefore, electron filling starts from the K shell, then the L, M, and N accordingly.
Depending on the orbit number, each orbit can only hold a finite number of electrons. The formula can determine this.
Number of electrons = 2n2
where n is the orbit number. Thus, this rule is frequently referred to as the "2n2 rule". The number of electrons that each orbit can hold increases with respect to its 'n' value.
Shells and Subshells
The shells are the paths around the nucleus where the electrons move. The shell nearest to the nucleus has lower energy. The energy increases as it travels to the outermost shell. It can be denoted as K, L, M, N, etc. Subshells are seen within the shells, which serve as the paths of electrons. The names of the subshells are based on the quantum number of angular momentum. A shell has four different subshell types: s, p, d, and f. The lowest energy subshell is the s, followed by p, d, and f. The formula 2(2l+1) can be used to determine the maximum number of electrons that can be occupied by each subshell. As a result, the s, p, d, and f subshells should be able to hold 2, 6, 10, and 14 electrons, respectively.
Number of Electrons Occupied in Each Shell
The maximum number of electrons that the outermost shell of an atom can possess is eight. This is called the octet rule. Having the eight electrons in the outermost shell makes the atom stable. The outermost shell is also termed as the valence shell. The electrons occupied in the valence shell are called valence electrons.
Electronic Configuration of Atoms
Now let's put these rules to work and figure out the electronic configurations for some elements. For example, magnesium has an atomic number of 12. According to the first rule, the K shell is filled first as it is the lowest in energy. Only 2 out of 12 electrons are filled in the K shell. As a result, the second rule is also followed, because the K shell can only hold a fixed number of electrons, namely two. Now there is a balance of 10 electrons. As per the third rule, the valence shell, which is the L shell here, cannot hold more than 8 electrons. So out of 10, 8 electrons enter the L shell and the remaining 2 electrons enter the M shell. Hence, the M shell becomes the valence shell.
Electronic Configuration of Some Elements
Electronic Configuration of Ions
The electronic configuration for ions is similar to that of atoms. The electronic configuration of a cation can be found out by removing electrons from the valence shells. Similarly, the electronic configuration of anions can be found by adding electrons to the valence shells.
For example, Mg has an atomic number of 12. The electronic configuration of Mg is
Mg: 1s2 2s2 2p6 3s2
For Mg2+, the electronic configuration can be written by removing 2 electrons from the valence shell, that is, 3s. Hence, the electronic configuration becomes
Mg2+: 1s2 2s2 2p6
Similarly, for Cl-, the atomic number is 17, with the electron configuration is
Cl: 1s2 2s2 2p6 3s2 3p5
To write the electronic configuration of cl-anion, we have to add one electron to the outermost shell, 3p.
Cl-: 1s2 2s2 2p6 3s2 3p6.
In this way, we can predict the electronic configuration of any ion.
Electronic Configuration Diagram
Electrons are arranged into energy levels or shells around the nucleus of an atom. The orbital radius increases as the energy level increases. We depict the shell by drawing a circle. A dot or a cross represents each electron and represents the nucleus by the chemical symbol. Each electron in an atom is in a particular shell, and the electron must occupy the lowest available shell nearest the nucleus. So, when we draw the electronic configuration, we have to fill up each shell in turn, starting with the lowest.
Let’s take lithium as an example.
Lithium Electron Structure
Lithium has an atomic number of 3. It has 3 electrons and 3 protons - first electron into the first shell. And the second electron goes to the same shell. However, this shell can only contain a maximum of 2 electrons. So, the third element must fill in the next shell. This same process of filling electrons applies to even larger atoms too. The filling of electrons should be from the lowest to the highest energy levels.
Interesting Facts
Have you ever wondered why atomic shells are represented by the letters K, L, M, N, etc. instead of using A, B, C, etc. This is because of the fact that, at that point of time, only some orbitals were discovered. Scientists wanted to be certain that there was room to add more orbits inside and outside of the existing ones.
Key Features
An electronic configuration is the arrangement of electrons in an atom in a particular manner in the shells.
It follows the Bohr-Bury scheme to fill the electrons in the orbits.
A lower-energy orbit or shell will always be filled by an electron before moving on to a higher-energy orbit or shell.
According to the 2n2 rule, only a fixed number of electrons can be filled in each orbit, where n is the number of the orbit.
The maximum number of electrons that the outermost shell of an atom can contain is 8.
FAQs on Electronic Configuration of Atoms and Ions in Chemistry
1. What is electronic configuration of atoms and ions?
The electronic configuration of atoms and ions is the arrangement of electrons in different energy levels, sublevels, and orbitals around the nucleus. It shows how electrons are distributed in shells (n = 1, 2, 3…), subshells (s, p, d, f), and orbitals according to specific rules. For example, the electronic configuration of sodium (Na, Z = 11) is 1s2 2s2 2p6 3s1. For ions, electrons are added or removed from the outermost shell, such as Na+: 1s2 2s2 2p6.
2. How do you write the electronic configuration of an atom?
To write the electronic configuration of an atom, fill orbitals in order of increasing energy using the Aufbau principle.
- Step 1: Find the atomic number (number of electrons in a neutral atom).
- Step 2: Fill orbitals in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s…
- Step 3: Follow the Pauli exclusion principle (maximum 2 electrons per orbital with opposite spins).
- Step 4: Apply Hund’s rule (electrons occupy degenerate orbitals singly first).
3. What are the rules for writing electronic configuration?
The rules for writing electronic configuration are the Aufbau principle, Pauli exclusion principle, and Hund’s rule.
- Aufbau principle: Electrons fill lower energy orbitals first.
- Pauli exclusion principle: No two electrons in an atom have the same four quantum numbers; each orbital holds maximum 2 electrons with opposite spins.
- Hund’s rule: Electrons occupy degenerate orbitals singly before pairing.
4. What is the electronic configuration of common ions?
The electronic configuration of common ions is obtained by adding or removing electrons from the outermost shell of the atom.
- Na+ (Z = 11): 1s2 2s2 2p6
- Cl- (Z = 17): 1s2 2s2 2p6 3s2 3p6
- Mg2+ (Z = 12): 1s2 2s2 2p6
5. Why do atoms form ions based on electronic configuration?
Atoms form ions to achieve a more stable noble gas electronic configuration with a full valence shell.
- Metals lose electrons to form cations (e.g., Na → Na+ + e-).
- Non-metals gain electrons to form anions (e.g., Cl + e- → Cl-).
6. What is the difference between the electronic configuration of atoms and ions?
The difference between the electronic configuration of atoms and ions is that atoms are neutral while ions have gained or lost electrons.
- Atom: Number of electrons = atomic number (e.g., Al: 1s2 2s2 2p6 3s2 3p1).
- Ion: Electrons are removed or added (e.g., Al3+: 1s2 2s2 2p6).
7. How do you write the electronic configuration of transition metal ions?
For transition metal ions, electrons are removed first from the ns orbital before the (n−1)d orbital.
- Example: Fe (Z = 26) = 1s2 2s2 2p6 3s2 3p6 4s2 3d6
- Fe2+: remove two 4s electrons → [Ar] 3d6
- Fe3+: remove one more electron from 3d → [Ar] 3d5
8. What is noble gas configuration in electronic configuration?
A noble gas configuration is an electron arrangement identical to that of a noble gas, with a completely filled valence shell.
- Example: Neon (Ne) = 1s2 2s2 2p6
- Na+ and Mg2+ both have the same configuration as Ne.
9. What are the exceptions to the Aufbau principle in electronic configuration?
Some elements, especially transition metals, show exceptions to the Aufbau principle due to extra stability of half-filled and fully filled subshells.
- Chromium (Cr, Z = 24): Expected 4s2 3d4, actual [Ar] 3d5 4s1
- Copper (Cu, Z = 29): Expected 4s2 3d9, actual [Ar] 3d10 4s1
10. How is electronic configuration related to valency and chemical properties?
Electronic configuration determines valency and chemical properties by showing the number of electrons in the outermost shell.
- Valency depends on electrons lost, gained, or shared to complete the octet.
- Elements with similar valence configurations show similar chemical behavior (e.g., Group 1: ns1 configuration).
