Electronic Configuration
Complete Guide to Electron Arrangement in Atoms
Electronic configuration is one of the most important concepts in atomic structure. It describes how electrons are arranged in different energy levels and orbitals around the nucleus of an atom. Understanding electronic configuration helps explain chemical bonding, reactivity, periodic trends, and many physical properties of elements.
In this article, we will explore the concept of electronic configuration, the rules used to determine it, and the electronic configurations of the first few elements in the periodic table.
What is Electronic Configuration?
Electronic configuration refers to the distribution of electrons in the orbitals of an atom. Each electron occupies a specific orbital depending on its energy level.
Electrons fill orbitals in a specific sequence based on energy. Lower-energy orbitals fill first before higher-energy orbitals.
The order in which orbitals are filled follows the Aufbau Principle, which provides a systematic method for determining electronic configurations.
Order of Orbital Filling (Aufbau Principle)
According to the Aufbau principle, electrons occupy orbitals starting from the lowest energy level and gradually move to higher energy levels.
The filling order of orbitals is:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
This order is often remembered using the diagonal rule diagram shown in the image.
Major Principles of Electronic Configuration
To correctly write the electronic configuration of elements, three main principles are followed.
1. Aufbau Principle
The Aufbau principle states that electrons fill orbitals beginning with the lowest energy orbital.
For example:
1s fills before 2s
2s fills before 2p
3p fills before 4s
This explains why the filling sequence is not always in simple numerical order.
2. Hund’s Rule
Hund’s rule explains how electrons fill orbitals within the same subshell.
It states:
Electrons fill each orbital singly before pairing.
All electrons in singly occupied orbitals have the same spin.
For example, the p subshell has three orbitals.
If three electrons are present, they will occupy separate orbitals like this:
↑ ↑ ↑
instead of pairing immediately.
Incorrect arrangement:
↑↓ ↑ _
Correct arrangement:
↑ ↑ ↑
This arrangement reduces electron repulsion and makes the atom more stable.
3. Pauli Exclusion Principle
The Pauli Exclusion Principle states that:
No two electrons in an atom can have the same set of four quantum numbers.
Each orbital can hold maximum two electrons only.
These two electrons must have opposite spins.
Representation:
↑↓
One electron spins upward and the other spins downward.
Electronic Configuration of the First 28 Elements
Below is the electronic configuration of elements from Hydrogen to Nickel based on the Aufbau principle.
1. Hydrogen (H) – Atomic Number 1
Electronic configuration:
1s¹
2. Helium (He) – Atomic Number 2
Electronic configuration:
1s²
3. Lithium (Li) – Atomic Number 3
Electronic configuration:
1s² 2s¹
4. Beryllium (Be) – Atomic Number 4
Electronic configuration:
1s² 2s²
5. Boron (B) – Atomic Number 5
Electronic configuration:
1s² 2s² 2p¹
6. Carbon (C) – Atomic Number 6
Electronic configuration:
1s² 2s² 2p²
7. Nitrogen (N) – Atomic Number 7
Electronic configuration:
1s² 2s² 2p³
8. Oxygen (O) – Atomic Number 8
Electronic configuration:
1s² 2s² 2p⁴
9. Fluorine (F) – Atomic Number 9
Electronic configuration:
1s² 2s² 2p⁵
10. Neon (Ne) – Atomic Number 10
Electronic configuration:
1s² 2s² 2p⁶
Neon has a completely filled outer shell which makes it chemically stable.
11. Sodium (Na) – Atomic Number 11
Electronic configuration:
1s² 2s² 2p⁶ 3s¹
12. Magnesium (Mg) – Atomic Number 12
Electronic configuration:
1s² 2s² 2p⁶ 3s²
13. Aluminium (Al) – Atomic Number 13
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p¹
14. Silicon (Si) – Atomic Number 14
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p²
15. Phosphorus (P) – Atomic Number 15
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p³
16. Sulfur (S) – Atomic Number 16
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁴
17. Chlorine (Cl) – Atomic Number 17
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁵
18. Argon (Ar) – Atomic Number 18
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶
Argon has a fully filled outer shell which makes it an inert gas.
19. Potassium (K) – Atomic Number 19
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
20. Calcium (Ca) – Atomic Number 20
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
21. Scandium (Sc) – Atomic Number 21
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
22. Titanium (Ti) – Atomic Number 22
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²
23. Vanadium (V) – Atomic Number 23
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
24. Chromium (Cr) – Atomic Number 24
Chromium shows an exception:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵
This occurs because half-filled orbitals provide extra stability.
25. Manganese (Mn) – Atomic Number 25
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵
26. Iron (Fe) – Atomic Number 26
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
27. Cobalt (Co) – Atomic Number 27
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷
28. Nickel (Ni) – Atomic Number 28
Electronic configuration:
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸
Example Problem
Find the electronic configuration of Oxygen (Atomic number = 8).
Step-by-step filling:
1s → 2 electrons
2s → 2 electrons
2p → remaining 4 electrons
Final configuration:
1s² 2s² 2p⁴
Practice Questions
1. Write the electronic configuration of Nitrogen (Atomic number 7).
2. What is the electronic configuration of Magnesium (Atomic number 12)?
3. Why does Chromium show an exceptional electronic configuration?
4. State Hund’s rule with an example.
Conclusion
Electronic configuration explains how electrons are distributed among orbitals in an atom. This arrangement determines the chemical behavior and physical properties of elements.
To correctly determine electronic configuration, three main principles must always be followed:
Aufbau Principle – electrons fill lowest energy orbitals first
Hund’s Rule – electrons occupy orbitals singly before pairing
Pauli Exclusion Principle – each orbital holds a maximum of two electrons with opposite spins
Understanding these rules helps students grasp advanced topics such as chemical bonding, periodic trends, and spectroscopy.
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