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| Electronic Configuration of Atoms – Complete CBSE Class 11 Chemistry Notes with MCQs, Assertion, Case Studies and Important Questions. |
Electronic Configuration of Atoms Class 11 Chemistry Notes PDF
CBSE Class 11 Chemistry – Electronic Configuration of Atoms
Question Bank with Answers (Based on NCERT)
A. Multiple Choice Questions (MCQs)
1. The distribution of electrons into orbitals of an atom is called:
a) Ionisation b) Electronic configuration c) Hybridisation d) Excitation
Answer: b) Electronic configuration
2. According to Pauli Exclusion Principle, an orbital can contain a maximum of:
a) 1 electron b) 2 electrons c) 4 electrons d) 6 electrons
Answer: b) 2 electrons
3. The electronic configuration of hydrogen is:
a) 1s² b) 2s¹ c) 1s¹ d) 2p¹
Answer: c) 1s¹
4. The electronic configuration of helium is:
a) 1s¹ b) 1s² c) 2s² d) 2p²
Answer: b) 1s²
5. Which orbital is filled after 3p?
a) 3d b) 4s c) 4p d) 5s
Answer: b) 4s
6. Which element has the electronic configuration [Ne]3s¹?
a) Mg b) Na c) Al d) Si
Answer: b) Na
7. Which of the following has a completely filled second shell?
a) Carbon b) Nitrogen c) Fluorine d) Neon
Answer: d) Neon
8. Chromium shows exceptional configuration because of:
a) Small atomic size b) Half-filled d-orbital stability c) Large nucleus d) Shielding effect
Answer: b) Half-filled d-orbital stability
9. Copper has the electronic configuration:
a) [Ar]3d⁹4s² b) [Ar]3d¹⁰4s¹ c) [Ar]3d⁸4s² d) [Ar]3d⁵4s¹
Answer: b) [Ar]3d¹⁰4s¹
10. The electrons present in the outermost shell are called:
a) Core electrons b) Inner electrons c) Valence electrons d) Spin electrons
Answer: c) Valence electrons
B. Very Short Answer Questions (1 Mark)
Q1. What is electronic configuration?
Answer: The arrangement of electrons in different orbitals of an atom is called electronic configuration.
Q2. Write the electronic configuration of hydrogen.
Answer: 1s¹
Q3. Write the electronic configuration of helium.
Answer: 1s²
Q4. Define core electrons.
Answer: Electrons present in completely filled inner shells are called core electrons.
Q5. Define valence electrons.
Answer: Electrons present in the outermost shell are called valence electrons.
Q6. Which orbital is filled before 3d?
Answer: 4s orbital
Q7. Which principle prevents more than two electrons in one orbital?
Answer: Pauli Exclusion Principle.
Q8. Name two elements showing exceptional electronic configuration.
Answer: Chromium and Copper.
Q9. Which configuration is more stable?
Answer: Half-filled and completely filled subshells.
Q10. Write the noble gas notation for sodium.
Answer: [Ne]3s¹
C. Short Answer Questions (2–3 Marks)
Q1. What are the two methods of representing electronic configuration?
Answer:
- Orbital notation (1s²2s²2p⁶)
- Orbital diagram (box-arrow representation)
Q2. Why does lithium not enter the third electron into 1s orbital?
Answer:
According to Pauli Exclusion Principle, an orbital can accommodate only two electrons with opposite spins. Hence the third electron enters the 2s orbital.
Q3. Why are Chromium and Copper exceptions?
Answer:
They attain extra stability due to half-filled (d⁵) and completely filled (d¹⁰) d-subshells.
Cr = [Ar]3d⁵4s¹
Cu = [Ar]3d¹⁰4s¹
Q4. Differentiate between core electrons and valence electrons.
| Core Electrons | Valence Electrons |
|---|---|
| Present in inner shells | Present in outermost shell |
| Do not participate in bonding | Participate in bonding |
Q5. Why is electronic configuration important?
Answer:
It helps to explain:
- Chemical bonding
- Reactivity
- Periodic properties
- Metallic and non-metallic nature
- Stability of atoms
D. Long Answer Questions (5 Marks)
Q1. Explain the filling of electrons from Hydrogen to Argon.
Answer:
- Hydrogen → 1s¹
- Helium → 1s²
- Lithium → 1s²2s¹
- Beryllium → 1s²2s²
- Boron to Neon → Filling of 2p orbitals
- Sodium to Argon → Filling of 3s and 3p orbitals
The filling follows:
- Aufbau Principle
- Pauli Exclusion Principle
- Hund's Rule
Q2. Explain why half-filled and fully-filled orbitals are more stable.
Answer:
Half-filled and completely filled orbitals possess:
- Symmetrical distribution of electrons
- Maximum exchange energy
- Minimum electron-electron repulsion
Examples:
- p³
- p⁶
- d⁵
- d¹⁰
- f⁷
- f¹⁴
Hence Chromium and Copper show exceptional configurations.
E. Assertion and Reason Questions
Q1.
Assertion (A): Chromium has electronic configuration [Ar]3d⁵4s¹.
Reason (R): Half-filled d-orbitals possess extra stability.
Answer: ✔ Both A and R are true and R is the correct explanation.
Q2.
Assertion (A): Helium is chemically inert.
Reason (R): Helium has a completely filled 1s orbital.
Answer: ✔ Both A and R are true and R explains A.
Q3.
Assertion (A): Sodium has one valence electron.
Reason (R): Sodium belongs to Group 1.
Answer: ✔ Both A and R are true and R explains A.
Q4.
Assertion (A): Copper has configuration [Ar]3d⁹4s².
Reason (R): Fully filled d-orbitals are stable.
Answer: ❌ Assertion is false, Reason is true.
Q5.
Assertion (A): Neon is reactive.
Reason (R): It has incomplete octet.
Answer: ❌ Both Assertion and Reason are false.
F. Fill in the Blanks
-
The arrangement of electrons is called electronic configuration.
-
An orbital can contain maximum 2 electrons.
-
The electronic configuration of hydrogen is 1s¹.
-
The electronic configuration of helium is 1s².
-
Sodium is represented as [Ne]3s¹.
-
Chromium has half-filled d-subshell.
-
Copper has fully-filled d-subshell.
-
Electrons in the outermost shell are called valence electrons.
-
The 4s orbital fills before 3d.
-
Neon has 10 electrons.
G. Case Study Questions
Case Study
A student studies the electronic configurations of Hydrogen to Copper. He observes that Chromium and Copper do not follow the expected electronic configuration. He also notices that half-filled and fully-filled orbitals possess greater stability.
Questions
1. Why does Chromium show exceptional configuration?
Answer: Because half-filled d-orbitals are more stable.
2. Write electronic configuration of Chromium.
Answer: [Ar]3d⁵4s¹
3. Write electronic configuration of Copper.
Answer: [Ar]3d¹⁰4s¹
4. Which principle limits two electrons in one orbital?
Answer: Pauli Exclusion Principle.
5. What is the significance of electronic configuration?
Answer: It explains chemical behaviour and bonding.
H. Statement-Based Questions
State whether True or False.
-
Helium has one electron. False
-
Neon has a completely filled second shell. True
-
Valence electrons take part in bonding. True
-
Core electrons are outermost electrons. False
-
Chromium has d⁵ configuration. True
-
Copper has d¹⁰ configuration. True
-
Sodium has one valence electron. True
-
The 3d orbital fills before 4s. False
-
Hydrogen configuration is 1s¹. True
-
Electronic configuration explains chemical behaviour. True
I. Match the Columns
| Column A | Column B |
|---|---|
| Hydrogen | 1s¹ |
| Helium | 1s² |
| Sodium | [Ne]3s¹ |
| Chromium | [Ar]3d⁵4s¹ |
| Copper | [Ar]3d¹⁰4s¹ |
Answers
- Hydrogen → 1s¹
- Helium → 1s²
- Sodium → [Ne]3s¹
- Chromium → [Ar]3d⁵4s¹
- Copper → [Ar]3d¹⁰4s¹
J. Important CBSE Competency-Based Questions
Q1. Why is Helium stable while Hydrogen is reactive?
Answer: Helium has a completely filled 1s orbital (duplet), making it chemically inert, whereas hydrogen has only one electron and tends to gain, lose, or share an electron to achieve stability.
Q2. Why do valence electrons determine chemical properties?
Answer: Valence electrons participate in chemical bonding and reactions, so they govern an element's chemical behaviour.
Q3. Why does the 4s orbital fill before the 3d orbital?
Answer: The 4s orbital has slightly lower energy than the 3d orbital in neutral atoms, so it is filled first according to the Aufbau principle.
Q4. Explain the importance of noble gas notation.
Answer: Noble gas notation provides a shorter way to write electronic configurations by replacing the inner-shell (core) electrons with the symbol of the nearest preceding noble gas, making configurations easier to read and write.
Q5. Explain the difference between orbital notation and orbital diagram.
Answer: Orbital notation uses symbols such as 1s² 2s² 2p⁶ to show electron distribution, whereas an orbital diagram uses boxes for orbitals and arrows (↑, ↓) to represent electrons and their spins, showing all four quantum numbers more clearly.
Internal Links
Structure of Atom Class 11 Notes
Quantum Numbers Class 11
Bohr's Atomic Model
Dual Nature of Matter
Aufbau Principle Explained
Hund's Rule with Examples
Pauli Exclusion Principle
Orbitals and Shapes of Orbitals
Periodic Table Class 11
Chemical Bonding Class 11 Notes
CBSE Class 11 Chemistry MCQs
NCERT Solutions for Class 11 Chemistry
Important Questions for Class 11 Chemistry
Sample Papers Class 11 Chemistry
Previous Year CBSE Chemistry Questions
Electronic Configuration Beyond Argon (NEET Notes)
1. Electronic Configuration from Sodium to Argon
Instead of writing the full electronic configuration every time, we use the noble gas shorthand.
[Ne] represents the electronic configuration of Neon.
- Na = [Ne]3s¹
- Mg = [Ne]3s²
- Al = [Ne]3s²3p¹
- Si = [Ne]3s²3p²
- P = [Ne]3s²3p³
- S = [Ne]3s²3p⁴
- Cl = [Ne]3s²3p⁵
- Ar = [Ne]3s²3p⁶
2. Core Electrons
- Core electrons are the electrons present in completely filled inner shells.
- They usually do not participate in chemical bonding.
Na = 1s²2s²2p⁶3s¹
- Core electrons = 1s²2s²2p⁶
- Core electrons = 10
3. Valence Electrons
- Valence electrons are the electrons present in the outermost shell.
- They determine the chemical properties of an element.
- They participate in chemical bonding.
Na = [Ne]3s¹
Valence electrons = 1
4. Potassium (K)
- Atomic Number = 19
- Next electron enters 4s orbital.
- 4s has lower energy than 3d.
Configuration: [Ar]4s¹
5. Calcium (Ca)
- Atomic Number = 20
- Second electron also enters 4s.
Configuration: [Ar]4s²
6. Beginning of Transition Elements
- Starts from Scandium (Sc).
- After 4s, electrons start filling 3d orbitals.
4s fills before 3d.
After 4s is filled, 3d fills before 4p.
7. First Transition Series
- Sc
- Ti
- V
- Cr
- Mn
- Fe
- Co
- Ni
- Cu
- Zn
8. Chromium Exception
Expected: [Ar]3d⁴4s²
Actual: [Ar]3d⁵4s¹
Reason: Half-filled d⁵ subshell is more stable.
9. Copper Exception
Expected: [Ar]3d⁹4s²
Actual: [Ar]3d¹⁰4s¹
Reason: Fully-filled d¹⁰ subshell is more stable.
10. Extra Stability
- Half-filled → p³, d⁵, f⁷
- Fully-filled → p⁶, d¹⁰, f¹⁴
11. After Zinc
- 3d becomes completely filled.
- Now 4p orbitals start filling.
12. Gallium to Krypton
Filling order:
4p¹ → 4p⁶
13. Rubidium to Xenon
Order of filling:
- 5s
- 4d
- 5p
14. Caesium and Barium
- Cs = 6s¹
- Ba = 6s²
15. Lanthanum to Mercury
- 4f orbitals fill.
- 5d orbitals fill.
16. After Mercury
- 6p
- 7s
- 5f
- 6d
17. Uranium and Beyond
- Called Transuranium elements.
- Artificially produced.
- Radioactive.
- Short-lived.
18. Importance of Electronic Configuration
- Explains chemical bonding.
- Explains reactivity.
- Explains periodic properties.
- Explains metals and non-metals.
- Foundation of modern chemistry.
19. Noble Gases
- He
- Ne
- Ar
Completely filled outer shell → Very stable.
20. Halogens
- F
- Cl
- Br
- I
Seven valence electrons → Highly reactive.
21. NEET Quick Revision
- Core electrons = Inner electrons.
- Valence electrons = Outermost electrons.
- 4s fills before 3d.
- 3d fills before 4p.
- Cr = [Ar]3d⁵4s¹.
- Cu = [Ar]3d¹⁰4s¹.
- Half-filled and fully-filled subshells are extra stable.
- Electronic configuration explains chemical behaviour.
Mind Map (Plain Text Tree)
Electronic Configuration (Beyond Argon)
│
├── Noble Gas Shorthand
│ ├── [Ne] = 1s²2s²2p⁶
│ ├── Na = [Ne]3s¹
│ ├── Mg = [Ne]3s²
│ └── Ar = [Ne]3s²3p⁶
│
├── Types of Electrons
│ ├── Core Electrons
│ │ ├── Inner shell electrons
│ │ ├── Completely filled shells
│ │ └── Not involved in bonding
│ │
│ └── Valence Electrons
│ ├── Outermost shell
│ ├── Chemical bonding
│ └── Chemical properties
│
├── After Argon
│ ├── K → [Ar]4s¹
│ ├── Ca → [Ar]4s²
│ └── 4s fills before 3d
│
├── First Transition Series
│ ├── Sc
│ ├── Ti
│ ├── V
│ ├── Cr
│ ├── Mn
│ ├── Fe
│ ├── Co
│ ├── Ni
│ ├── Cu
│ └── Zn
│
├── Exceptions
│ ├── Cr → 3d⁵4s¹
│ └── Cu → 3d¹⁰4s¹
│
├── Extra Stability
│ ├── Half-filled
│ │ ├── p³
│ │ ├── d⁵
│ │ └── f⁷
│ └── Fully-filled
│ ├── p⁶
│ ├── d¹⁰
│ └── f¹⁴
│
├── Filling Order
│ ├── 4p
│ ├── 5s
│ ├── 4d
│ ├── 5p
│ ├── 6s
│ ├── 4f
│ ├── 5d
│ ├── 6p
│ ├── 7s
│ ├── 5f
│ └── 6d
│
├── Uranium and Beyond
│ ├── Artificial
│ ├── Radioactive
│ └── Short-lived
│
├── Importance
│ ├── Chemical bonding
│ ├── Reactivity
│ ├── Periodic properties
│ ├── Metals & Non-metals
│ └── Modern Chemistry
│
└── NEET Facts
├── Core electrons
├── Valence electrons
├── 4s before 3d
├── Cr exception
├── Cu exception
└── Stable half-filled & fully-filled subshells
Electronic Configuration of Atoms
Definition:
Electronic configuration is the distribution of electrons into different orbitals of an atom.
Ways to Represent Electronic Configuration
1. Subshell Notation
Example:
- 1s¹
- 1s²
- 1s² 2s¹
Meaning:
- Number = Shell
- Letter = Subshell
- Superscript = Number of electrons
2. Orbital Diagram
Each orbital is represented by a box (□).
Electrons are represented by arrows.
- ↑ = Positive spin
- ↓ = Negative spin
Hydrogen (H)
Atomic Number = 1
Configuration = 1s¹
1s □ ↑
Helium (He)
Atomic Number = 2
Configuration = 1s²
1s □ ↑↓
Both electrons have opposite spins.
Lithium (Li)
Atomic Number = 3
Configuration = 1s² 2s¹
1s 2s □ ↑↓ □ ↑
Third electron enters 2s because 1s is already full.
Beryllium (Be)
Atomic Number = 4
Configuration = 1s² 2s²
1s 2s □ ↑↓ □ ↑↓
Filling of 2p Orbitals
| Element | Electronic Configuration |
|---|---|
| B | 1s² 2s² 2p¹ |
| C | 1s² 2s² 2p² |
| N | 1s² 2s² 2p³ |
| O | 1s² 2s² 2p⁴ |
| F | 1s² 2s² 2p⁵ |
| Ne | 1s² 2s² 2p⁶ |
Orbital Filling
B □ ↑ □ □ C □ ↑ □ ↑ □ N □ ↑ □ ↑ □ ↑ O □ ↑↓ □ ↑ □ ↑ F □ ↑↓ □ ↑↓ □ ↑ Ne □ ↑↓ □ ↑↓ □ ↑↓
Elements from Sodium to Argon
| Element | Electronic Configuration |
|---|---|
| Na | [Ne] 3s¹ |
| Mg | [Ne] 3s² |
| Al | [Ne] 3s²3p¹ |
| Si | [Ne] 3s²3p² |
| P | [Ne] 3s²3p³ |
| S | [Ne] 3s²3p⁴ |
| Cl | [Ne] 3s²3p⁵ |
| Ar | [Ne] 3s²3p⁶ |
Orbital Capacity
| Subshell | Orbitals | Maximum Electrons |
|---|---|---|
| s | 1 | 2 |
| p | 3 | 6 |
| d | 5 | 10 |
| f | 7 | 14 |
NEET Important Points
- Electronic configuration = Arrangement of electrons in orbitals.
- Lowest energy orbital fills first.
- Maximum 2 electrons in one orbital.
- Electrons in one orbital have opposite spins.
- Half-filled and fully-filled subshells are more stable.
- Hund's Rule: Fill empty orbitals before pairing.
- Neon shorthand: [Ne] = 1s²2s²2p⁶
- Orbital filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s
Mind Map
Electronic Configuration of Atoms
│
├── Definition
│ ├── Distribution of electrons in orbitals
│ └── Arrangement of electrons
│
├── Basic Rules
│ ├── Aufbau Principle
│ │ └── Lowest energy fills first
│ ├── Pauli Exclusion Principle
│ │ ├── Max 2 electrons/orbital
│ │ └── Opposite spins
│ └── Hund's Rule
│ ├── Fill singly first
│ └── Pair later
│
├── Representation
│ ├── Subshell Notation
│ │ ├── Number = Shell
│ │ ├── Letter = Subshell
│ │ └── Superscript = Electrons
│ │
│ └── Orbital Diagram
│ ├── □ = Orbital
│ ├── ↑ = Positive Spin
│ └── ↓ = Negative Spin
│
├── Filling
│ ├── 1s
│ │ ├── H
│ │ └── He
│ ├── 2s
│ │ ├── Li
│ │ └── Be
│ ├── 2p
│ │ ├── B
│ │ ├── C
│ │ ├── N
│ │ ├── O
│ │ ├── F
│ │ └── Ne
│ └── 3s & 3p
│ ├── Na
│ ├── Mg
│ ├── Al
│ ├── Si
│ ├── P
│ ├── S
│ ├── Cl
│ └── Ar
│
├── Orbital Capacity
│ ├── s = 2
│ ├── p = 6
│ ├── d = 10
│ └── f = 14
│
└── NEET Quick Facts
├── Max 2 electrons/orbital
├── Opposite spins
├── Stable half-filled subshell
├── Stable fully-filled subshell
└── Filling Order
1s
↓
2s
↓
2p
↓
3s
↓
3p
↓
4s
↓
3d
↓
4p
↓
5s
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