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Electronic Configuration: A Complete Guide for Chemistry Students
In the world of chemistry, understanding how electrons are arranged around the nucleus of an atom is key to unlocking the mysteries of chemical bonding, atomic behavior, and the periodic table. This arrangement is known as electronic configuration.
The electronic configuration of an element describes the distribution of electrons in atomic orbitals. It provides insights into the chemical properties of elements and plays a crucial role in predicting chemical reactions and trends in the periodic table.
Whether you’re a class 11 or 12 student, or preparing for NEET or JEE, mastering electronic configuration is a foundational skill.
What is Electronic Configuration?
Electronic configuration is the representation of electrons in atomic orbitals. It follows a specific order based on the energy levels of orbitals. An electron configuration typically looks like this: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
Why is Electronic Configuration Important?
- Predicts chemical reactivity of elements
- Determines valency and bonding behavior
- Explains the position of elements in the periodic table
- Helps in understanding ion formation
- Crucial for JEE/NEET chemistry questions
Key Concepts Behind Electronic Configuration
Atomic Orbitals and Subshells
Electrons reside in regions around the nucleus called orbitals. These orbitals are grouped into four subshells:
| Subshell | Maximum Electrons | Shape |
|---|---|---|
| s | 2 | Spherical |
| p | 6 | Dumbbell |
| d | 10 | Clover-leaf |
| f | 14 | Complex |
Quantum Numbers
Each electron in an atom is described using four quantum numbers:
- Principal quantum number (n): Energy level (1, 2, 3…)
- Azimuthal quantum number (l): Subshell type (s=0, p=1, d=2, f=3)
- Magnetic quantum number (m): Orientation of orbital
- Spin quantum number (s): Spin of electron (+½ or -½)
These numbers help locate an electron precisely in an atom.
Rules for Writing Electronic Configuration
Electrons fill orbitals starting from the lowest energy level first. The order of filling is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …
Pauli Exclusion Principle
No two electrons in an atom can have the same set of four quantum numbers. Hence, one orbital can hold a maximum of two electrons with opposite spins.
Hund’s Rule
When filling degenerate (same energy) orbitals, electrons will first occupy all orbitals singly before pairing up. This minimizes repulsion and increases stability.
Electronic Configuration of First 30 Elements
| Atomic Number | Element | Electronic Configuration |
|---|---|---|
| 1 | Hydrogen (H) | 1s¹ |
| 2 | Helium (He) | 1s² |
| 3 | Lithium (Li) | 1s² 2s¹ |
| 4 | Beryllium (Be) | 1s² 2s² |
| 5 | Boron (B) | 1s² 2s² 2p¹ |
| 6 | Carbon (C) | 1s² 2s² 2p² |
| 7 | Nitrogen (N) | 1s² 2s² 2p³ |
| 8 | Oxygen (O) | 1s² 2s² 2p⁴ |
| 9 | Fluorine (F) | 1s² 2s² 2p⁵ |
| 10 | Neon (Ne) | 1s² 2s² 2p⁶ |
| 11 | Sodium (Na) | 1s² 2s² 2p⁶ 3s¹ |
| 12 | Magnesium (Mg) | 1s² 2s² 2p⁶ 3s² |
| 13 | Aluminium (Al) | 1s² 2s² 2p⁶ 3s² 3p¹ |
| 14 | Silicon (Si) | 1s² 2s² 2p⁶ 3s² 3p² |
| 15 | Phosphorus (P) | 1s² 2s² 2p⁶ 3s² 3p³ |
| 16 | Sulfur (S) | 1s² 2s² 2p⁶ 3s² 3p⁴ |
| 17 | Chlorine (Cl) | 1s² 2s² 2p⁶ 3s² 3p⁵ |
| 18 | Argon (Ar) | 1s² 2s² 2p⁶ 3s² 3p⁶ |
| 19 | Potassium (K) | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ |
| 20 | Calcium (Ca) | 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² |
| 21 | Scandium (Sc) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s² |
| 22 | Titanium (Ti) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d² 4s² |
| 23 | Vanadium (V) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d³ 4s² |
| 24 | Chromium (Cr) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ |
| 25 | Manganese (Mn) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s² |
| 26 | Iron (Fe) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s² |
| 27 | Cobalt (Co) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁷ 4s² |
| 28 | Nickel (Ni) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s² |
| 29 | Copper (Cu) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹ |
| 30 | Zinc (Zn) | 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² |
Electronic Configuration of s, p, d, and f Block Elements
s-Block Elements (Groups 1 and 2)
- Location: Leftmost part of the Periodic Table
- General Configuration:
ns¹–² - Electrons fill the s-orbital
| Element | Atomic Number | Electronic Configuration |
|---|---|---|
| Hydrogen (H) | 1 | 1s¹ |
| Helium (He) | 2 | 1s² |
| Lithium (Li) | 3 | 1s² 2s¹ |
| Beryllium (Be) | 4 | 1s² 2s² |
| Sodium (Na) | 11 | 1s² 2s² 2p⁶ 3s¹ |
| Magnesium (Mg) | 12 | 1s² 2s² 2p⁶ 3s² |
🧪 2. p-Block Elements (Groups 13 to 18)
- Location: Right side of Periodic Table
- General Configuration:
ns² np¹–⁶ - Electrons fill the p-orbital
| Element | Atomic Number | Electronic Configuration |
|---|---|---|
| Boron (B) | 5 | 1s² 2s² 2p¹ |
| Carbon (C) | 6 | 1s² 2s² 2p² |
| Nitrogen (N) | 7 | 1s² 2s² 2p³ |
| Oxygen (O) | 8 | 1s² 2s² 2p⁴ |
| Fluorine (F) | 9 | 1s² 2s² 2p⁵ |
| Neon (Ne) | 10 | 1s² 2s² 2p⁶ |
d-Block Elements (Transition Elements – Groups 3 to 12)
- Location: Center of Periodic Table
- General Configuration:
ns² (n-1)d¹–¹⁰ - Electrons fill d-orbitals
- Exceptions: Chromium (24) → 4s¹ 3d⁵, Copper (29) → 4s¹ 3d¹⁰
| Element | Atomic Number | Electronic Configuration |
|---|---|---|
| Scandium (Sc) | 21 | [Ar] 4s² 3d¹ |
| Titanium (Ti) | 22 | [Ar] 4s² 3d² |
| Vanadium (V) | 23 | [Ar] 4s² 3d³ |
| Chromium (Cr) | 24 | [Ar] 4s¹ 3d⁵ |
| Manganese (Mn) | 25 | [Ar] 4s² 3d⁵ |
| Iron (Fe) | 26 | [Ar] 4s² 3d⁶ |
| Cobalt (Co) | 27 | [Ar] 4s² 3d⁷ |
| Nickel (Ni) | 28 | [Ar] 4s² 3d⁸ |
| Copper (Cu) | 29 | [Ar] 4s¹ 3d¹⁰ |
| Zinc (Zn) | 30 | [Ar] 4s² 3d¹⁰ |
f-Block Elements (Lanthanides & Actinides)
- Location: Bottom of the Periodic Table
- General Configuration:
- Lanthanides:
6s² 4f¹–¹⁴ - Actinides:
7s² 5f¹–¹⁴ - Electrons fill the f-orbital
- Often written as exceptions due to complex behavior
| Element | Atomic Number | Electronic Configuration |
|---|---|---|
| Cerium (Ce) | 58 | [Xe] 6s² 4f¹ 5d¹ |
| Gadolinium (Gd) | 64 | [Xe] 6s² 4f⁷ 5d¹ |
| Lutetium (Lu) | 71 | [Xe] 6s² 4f¹⁴ 5d¹ |
| Thorium (Th) | 90 | [Rn] 7s² 6d² |
| Uranium (U) | 92 | [Rn] 7s² 5f³ 6d¹ |
| Lawrencium (Lr) | 103 | [Rn] 7s² 5f¹⁴ 7p¹ |
Top Electron Configuration Exceptions (With Reason)
Electron configurations usually follow the Aufbau Principle, but a few elements violate this rule due to extra stability in half-filled or fully filled orbitals.
✅ Key Reason:
- Half-filled (e.g. d⁵) and fully filled (e.g. d¹⁰) sublevels are more stable.
- Some electrons are shifted from the s-orbital to the d-orbital to gain stability.
🔍 Electron Configuration Exceptions (Important for NEET, JEE, Class 11–12)
| Atomic No. | Element | Expected Configuration | Actual Configuration | Reason |
|---|---|---|---|---|
| 24 | Chromium (Cr) | [Ar] 4s² 3d⁴ | [Ar] 4s¹ 3d⁵ | Half-filled d⁵ is more stable |
| 29 | Copper (Cu) | [Ar] 4s² 3d⁹ | [Ar] 4s¹ 3d¹⁰ | Fully filled d¹⁰ is more stable |
| 42 | Molybdenum (Mo) | [Kr] 5s² 4d⁴ | [Kr] 5s¹ 4d⁵ | Half-filled d⁵ stability |
| 46 | Palladium (Pd) | [Kr] 5s² 4d⁸ | [Kr] 4d¹⁰ | No 5s electron; d¹⁰ full stability |
| 47 | Silver (Ag) | [Kr] 5s² 4d⁹ | [Kr] 5s¹ 4d¹⁰ | Full d¹⁰ preferred |
| 79 | Gold (Au) | [Xe] 6s² 4f¹⁴ 5d⁹ | [Xe] 6s¹ 4f¹⁴ 5d¹⁰ | Full d orbital stability |
Frequently Asked Questions (FAQs) – Electronic Configuration
1. What is electronic configuration?
Electronic configuration is the arrangement of electrons in an atom’s orbitals (s, p, d, f). It shows how electrons are distributed across energy levels and sub-levels based on the atomic number.
2. Why is electronic configuration important?
It helps in understanding chemical properties, bonding behavior, valency, and the placement of elements in the periodic table. It’s also crucial in solving NEET, JEE, and board exam questions.
3. How do I find the electronic configuration of an element?
You can find it by following the Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule — or use our Free Online Electronic Configuration Tool to get instant results!
4. What are spdf orbitals?
The orbitals are:
- s (holds 2 electrons)
- p (6 electrons)
- d (10 electrons)
- f (14 electrons)
They represent different sub-shells within electron shells.
5. What are the exceptions in electronic configuration?
Some elements like Chromium (24) and Copper (29) show irregular configurations due to stability of half-filled or fully-filled orbitals. Our tool highlights these exceptions automatically.
6. How accurate is this tool for electronic configuration?
Our tool uses standard rules of quantum mechanics and includes known exceptions, so it’s 99.9% accurate for all 118 elements.
7. Can I use this tool for NEET or JEE preparation?
Absolutely! This tool is built for school and competitive exam students. It helps you learn configurations quickly and understand atomic structure concepts easily.
8. Is this tool really free and unlimited?
Yes, it’s 100% free, unlimited, and does not require any login or sign-up. Just enter an atomic number (1–118) and get instant results!
9. Does the tool show the valence electrons too?
Currently, it shows the complete configuration. In the next update, we’ll also display valence shell configuration, block, group, and period.
10. Who should use this Electronic Configuration Tool?
This tool is ideal for:
- Students of Class 9–12
- NEET/JEE aspirants
- Teachers
- Chemistry enthusiasts
Relationship Between Electronic Configuration and the Periodic Table
Periodic Table is Based on Electronic Configuration
The modern periodic table is structured according to the electronic configurations of elements. This is why elements with similar properties fall into the same groups/columns.
Periods (Horizontal Rows) and Electron Shells
Each period in the periodic table corresponds to a principal energy level (n) or shell.
Example:
- Hydrogen (H) and Helium (He) are in Period 1 → Their electrons fill the 1st shell.
- Sodium (Na) is in Period 3 → Its electrons fill up to the 3rd shell.
Groups (Vertical Columns) and Valence Electrons
- Elements in the same group have the same number of valence electrons, which is why they show similar chemical properties. Example:
- Group 1 (Li, Na, K…) all have 1 electron in their outermost shell → Highly reactive metals.
Blocks of the Periodic Table and Subshells
The periodic table is divided into blocks based on which subshell (s, p, d, f) is being filled:
- s-block: Groups 1 & 2 → s-orbital filling
- p-block: Groups 13–18 → p-orbital filling
- d-block: Transition metals (Groups 3–12) → d-orbital filling
- f-block: Lanthanides & Actinides → f-orbital filling
Reactivity Trends Explained by Configuration
- Alkali metals (Group 1): 1 valence electron → Very reactive
- Noble gases (Group 18): Full outer shell → Very stable
- Halogens (Group 17): Need 1 electron → Very reactive non-metals
Example: Sodium (Na)
- Atomic number: 11
- Electronic configuration: 1s² 2s² 2p⁶ 3s¹
- It is in:
- Period 3 → 3 energy levels
- Group 1 → 1 electron in the outermost shell (3s¹)
- s-block → Filling s-orbital