Quantum Numbers (Principal, Azimuthal, Magnetic and Spin) – Definition, Detailed Explanation, FAQs of Quantum Numbers.

February 11, 2025December 12, 2024 by Neeraj Anand, Param Anand

JEE Syllabus for Chapter - ATOMIC STRUCTURE

According to the JEE syllabus, the "Atomic Structure" chapter covers key concepts like: Nature of electromagnetic radiation, photoelectric effect, spectrum of the hydrogen atom, Bohr model of ahydrogen atom - its postulates, derivation of the relations for the energy of the electron and radii of the different orbits, limitations of Bohr's model, dual nature of matter, de Broglie's relationship, Heisenberguncertainty principle, elementary ideas of quantum mechanics, the quantum mechanical model of the atomand its important features, concept of atomic orbitals as one-electron wave functions, variation of Ψ and Ψ² with r for 1s and 2s orbitals, various quantum numbers (principal, angular momentumand magneticquantum numbers) and their significance, shapes of s, p and d - orbitals, electron spin and spin quantumnumber, rules for filling electrons in orbitals – Aufbau principle, Pauli's exclusion principle and Hund'srule, electronic configuration of elements and extra stability of half-filled and completely filled orbitals.

NEET Syllabus for Chapter - ATOMIC STRUCTURE

According to the NEET syllabus, the "Atomic Structure" chapter covers key concepts like: subatomic particles (protons, electrons, neutrons), atomic number and mass number, various atomic models (Dalton's, Thomson's, Rutherford's, Bohr's), quantum mechanical model (Schrödinger's equation, quantum numbers - principal, azimuthal, magnetic, and spin), shapes of orbitals (s, p, d), electronic configuration of elements based on Aufbau principle, Pauli exclusion principle, and Hund's rule; including the limitations of Bohr's model and the concept of dual nature of matter with de Broglie's relationship and Heisenberg's uncertainty principle. 

MCQs on Structure of Atom for Class 11 CBSE Board Exam

Here are some multiple-choice questions (MCQs) on Structure of Atom for Class 11 CBSE along with detailed explanations:

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1. Which of the following statements about the nucleus of an atom is correct?

A) It contains protons and neutrons
B) It has a negative charge
C) It occupies most of the volume of the atom
D) It is responsible for chemical properties of an atom

Answer: A) It contains protons and neutrons

Explanation:
The nucleus of an atom contains protons (positively charged) and neutrons (neutral). The electrons revolve around the nucleus in different energy levels. The nucleus has a positive charge due to the presence of protons. It occupies a very small volume but contributes to almost the entire mass of the atom.

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2. The total number of electrons that can be accommodated in the second shell (L-shell) is:

A) 2
B) 8
C) 18
D) 32

Answer: B) 8

Explanation:
The maximum number of electrons in a shell is given by 2n², where n is the shell number.
For the second shell (n = 2):
Max electrons = 2(2²) = 8.

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3. Which of the following is NOT a postulate of Bohr’s atomic model?

A) Electrons revolve around the nucleus in fixed orbits
B) Electrons emit or absorb energy when they jump between orbits
C) Energy levels are quantized
D) Electrons can have any random energy value

Answer: D) Electrons can have any random energy value

Explanation:
Bohr's model states that electrons revolve in fixed energy levels, and they cannot have arbitrary energy. Electrons only gain or lose energy when they transition between these discrete orbits.

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4. The isotope of hydrogen that contains one proton and two neutrons is:

A) Protium
B) Deuterium
C) Tritium
D) None of these

Answer: C) Tritium

Explanation:

• Protium (¹H) → 1 proton, 0 neutrons
• Deuterium (²H) → 1 proton, 1 neutron
• Tritium (³H) → 1 proton, 2 neutrons

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5. The wave nature of electrons was proposed by:

A) Bohr
B) Heisenberg
C) de Broglie
D) Rutherford

Answer: C) de Broglie

Explanation:
Louis de Broglie proposed that electrons exhibit both particle and wave nature (wave-particle duality). His equation λ = h/mv relates the wavelength (λ) of a moving particle to its momentum.

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6. The uncertainty principle was proposed by:

A) Bohr
B) Heisenberg
C) Rutherford
D) Schrodinger

Answer: B) Heisenberg

Explanation:
Heisenberg's Uncertainty Principle states that it is impossible to simultaneously determine the exact position and momentum of an electron.

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7. The quantum number that describes the shape of an orbital is:

A) Principal quantum number (n)
B) Azimuthal quantum number (l)
C) Magnetic quantum number (m)
D) Spin quantum number (s)

Answer: B) Azimuthal quantum number (l)

Explanation:
The Azimuthal quantum number (l) determines the shape of orbitals:

• s-orbital (l = 0) → Spherical
• p-orbital (l = 1) → Dumbbell
• d-orbital (l = 2) → Complex
• f-orbital (l = 3) → More complex

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8. Which of the following orbitals cannot exist?

A) 1s
B) 2p
C) 3f
D) 4d

Answer: C) 3f

Explanation:
For an orbital to exist, the Azimuthal quantum number (l) must satisfy:
l = 0 to (n-1), where n is the principal quantum number.
For n = 3, possible l values: 0 (s), 1 (p), 2 (d) → No f-orbital.

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9. Which of the following elements has the electronic configuration: 1s² 2s² 2p⁶ 3s¹?

A) Sodium (Na)
B) Magnesium (Mg)
C) Aluminium (Al)
D) Potassium (K)

Answer: A) Sodium (Na)

Explanation:

• 1s² 2s² 2p⁶ 3s¹ is the electronic configuration of sodium (Na) (Atomic number 11).
• Magnesium (Mg) = 1s² 2s² 2p⁶ 3s²
• Aluminium (Al) = 1s² 2s² 2p⁶ 3s² 3p¹

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10. The shape of an s-orbital is:

A) Spherical
B) Dumbbell
C) Double dumbbell
D) Complex

Answer: A) Spherical

Explanation:
The s-orbital is spherically symmetric around the nucleus. p-orbitals are dumbbell-shaped.

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11. The number of orbitals present in the third energy level (n = 3) is:

A) 3
B) 9
C) 18
D) 5

Answer: B) 9

Explanation:
Total orbitals in an energy level = n²
For n = 3, orbitals = 3² = 9
(1 s-orbital, 3 p-orbitals, 5 d-orbitals)

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12. If an electron has quantum numbers n = 3, l = 2, what type of orbital is it in?

A) 3s
B) 3p
C) 3d
D) 3f

Answer: C) 3d

Explanation:

• n = 3 (Third shell)
• l = 2 corresponds to d-orbital → 3d orbital

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13. The maximum number of electrons that can be accommodated in an orbital is:

A) 1
B) 2
C) 4
D) 6

Answer: B) 2

Explanation:
Each orbital can hold a maximum of 2 electrons with opposite spins, as per Pauli's exclusion principle.

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14. Which quantum number determines the energy of an electron in a hydrogen atom?

A) Principal quantum number (n)
B) Azimuthal quantum number (l)
C) Magnetic quantum number (m)
D) Spin quantum number (s)

Answer: A) Principal quantum number (n)

Explanation:
For hydrogen-like atoms, the energy of an electron depends only on n.

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15. The concept of orbitals was introduced by:

A) Bohr
B) Rutherford
C) Schrodinger
D) Heisenberg

Answer: C) Schrodinger

Explanation:
Schrodinger’s wave equation introduced orbitals (regions of high probability of finding an electron), replacing Bohr’s fixed orbits.

Assertion and Reason (A-R) type questions on Structure of Atom for Class 11 CBSE Board Exam

Here are some Assertion and Reason (A-R) type questions on Structure of Atom for Class 11 CBSE, along with detailed explanations.

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How to Answer Assertion-Reason Questions?

Each question consists of two statements:

• Assertion (A): A statement of fact.
• Reason (R): An explanation for the assertion.

You must choose the correct option:

1. Both A and R are true, and R is the correct explanation of A.
2. Both A and R are true, but R is NOT the correct explanation of A.
3. A is true, but R is false.
4. A is false, but R is true.

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1. Assertion (A): The nucleus of an atom contains protons and neutrons.

Reason (R): The electrons in an atom revolve around the nucleus in fixed orbits.

✅ Answer: Option (2) (Both A and R are true, but R is not the correct explanation of A.)

🔹 Explanation:

• The nucleus is composed of protons and neutrons, which was confirmed by Rutherford’s experiment.
• Electrons revolve around the nucleus in fixed orbits, as described in Bohr’s atomic model.
• However, the presence of electrons in orbits does not explain why the nucleus contains protons and neutrons.

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2. Assertion (A): The maximum number of electrons in a shell is given by 2n².

Reason (R): Electrons in an atom are arranged in different shells around the nucleus.

✅ Answer: Option (1) (Both A and R are true, and R is the correct explanation of A.)

🔹 Explanation:

• The Bohr-Bury rule states that the maximum number of electrons in a shell = 2n², where n is the shell number.
• Electrons are distributed in different shells, and this follows the energy level distribution principle.

Example:

• K shell (n = 1): 2(1²) = 2 electrons
• L shell (n = 2): 2(2²) = 8 electrons
• M shell (n = 3): 2(3²) = 18 electrons

Thus, R correctly explains A.

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3. Assertion (A): The mass of an atom is concentrated in the nucleus.

Reason (R): The electrons have negligible mass compared to protons and neutrons.

✅ Answer: Option (1) (Both A and R are true, and R is the correct explanation of A.)

🔹 Explanation:

• The nucleus contains protons and neutrons, which are heavy particles.
• Electrons are much lighter (mass = 1/1836 of a proton), so their contribution to atomic mass is negligible.
• Thus, the nucleus contains almost all the mass of an atom, and R explains A correctly.

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4. Assertion (A): The energy of electrons in an atom is quantized.

Reason (R): Electrons can exist at any random energy level.

✅ Answer: Option (3) (A is true, but R is false.)

🔹 Explanation:

• According to Bohr’s model, electrons can only occupy specific, discrete energy levels (quantized energy states).
• Electrons cannot have arbitrary energy values.
• Hence, A is true, but R is false because electrons follow quantized energy levels.

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5. Assertion (A): The Heisenberg Uncertainty Principle states that the exact position and momentum of an electron cannot be simultaneously determined.

Reason (R): Electrons move in fixed circular orbits around the nucleus.

✅ Answer: Option (3) (A is true, but R is false.)

🔹 Explanation:

• Heisenberg’s Uncertainty Principle states:
  Δx × Δp ≥ h/4π
  (where Δx = uncertainty in position, Δp = uncertainty in momentum)
• This means we cannot know both the exact position and momentum of an electron at the same time.
• Bohr’s model (fixed orbits) was later replaced by Schrodinger’s model (probability orbitals).
• R is false because electrons do not move in fixed orbits but in probabilistic regions (orbitals).

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6. Assertion (A): The azimuthal quantum number (l) determines the shape of an orbital.

Reason (R): The principal quantum number (n) determines the size of the orbital.

✅ Answer: Option (1) (Both A and R are true, and R is the correct explanation of A.)

🔹 Explanation:

• The Azimuthal quantum number (l) determines the shape of an orbital (s, p, d, f).
• The Principal quantum number (n) determines the size and energy of an orbital.
• Since both statements are correct and R correctly explains A, the answer is Option (1).

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7. Assertion (A): The 3f orbital exists.

Reason (R): The azimuthal quantum number (l) for an f-orbital is 3.

✅ Answer: Option (4) (A is false, but R is true.)

🔹 Explanation:

• l = 3 represents an f-orbital, but for an orbital to exist, n must be greater than l.
• For n = 3, the possible values of l are 0 (s), 1 (p), and 2 (d).
• f-orbital (l = 3) is possible only when n ≥ 4, so 3f does not exist.
• Hence, A is false, but R is true.

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8. Assertion (A): The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers.

Reason (R): An orbital can accommodate a maximum of two electrons with opposite spins.

✅ Answer: Option (1) (Both A and R are true, and R is the correct explanation of A.)

🔹 Explanation:

• Pauli's Exclusion Principle states that no two electrons in an atom can have identical quantum numbers.
• This is because each orbital can hold only two electrons with opposite spins.
• Since R explains A correctly, the answer is Option (1).

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9. Assertion (A): The probability of finding an electron is maximum near the nucleus.

Reason (R): The probability density of an electron is given by the wave function (ψ²).

✅ Answer: Option (1) (Both A and R are true, and R is the correct explanation of A.)

🔹 Explanation:

• The Schrodinger wave equation gives ψ², which represents the probability density of finding an electron in a region.
• For s-orbitals, the highest probability is near the nucleus.
• Since R explains A correctly, the answer is Option (1).

FAQs (Important Questions & Answers) on Structure of Atom – Class 11

1. What is the structure of an atom?

The structure of an atom consists of a nucleus containing protons and neutrons, surrounded by electrons moving in discrete energy levels or shells around the nucleus.

2. Who discovered the atom?

The concept of the atom was first proposed by John Dalton in his Atomic Theory (1808). Later, significant contributions were made by J.J. Thomson, Rutherford, and Bohr to explain its structure.

3. What are the fundamental particles of an atom?

The three fundamental particles of an atom are:

• Proton (p⁺) – Positively charged, found in the nucleus.
• Neutron (n⁰) – Neutral, found in the nucleus.
• Electron (e⁻) – Negatively charged, revolves around the nucleus.

4. What is Dalton’s Atomic Theory?

Dalton’s Atomic Theory (1808) states that:

• Matter is made of indivisible atoms.
• Atoms of an element are identical in mass and properties.
• Atoms combine in whole-number ratios to form compounds.
• Atoms cannot be created or destroyed in a chemical reaction.

5. What was J.J. Thomson’s Model of the Atom?

J.J. Thomson proposed the Plum Pudding Model (1897), where an atom was visualized as a positively charged sphere with negatively charged electrons embedded in it, like plums in a pudding. However, this model was later disproven by Rutherford’s experiment.

6. What is Rutherford’s Atomic Model?

Rutherford’s Gold Foil Experiment (1911) led to the discovery that:

• An atom has a dense positively charged nucleus.
• Electrons revolve around the nucleus.
• Most of the atom is empty space.
  This model could not explain atomic stability, which was later addressed by Bohr’s model.

7. What is Bohr’s Atomic Model?

Niels Bohr (1913) proposed that:

• Electrons revolve in fixed energy levels (shells).
• They do not lose energy while in stable orbits.
• Energy is absorbed or emitted when an electron jumps between energy levels.

8. What are Quantum Numbers?

Quantum numbers define the position and energy of an electron in an atom:

1. Principal Quantum Number (n) – Represents the main energy level.
2. Azimuthal Quantum Number (l) – Defines the shape of the orbital.
3. Magnetic Quantum Number (mₗ) – Indicates the orientation of the orbital.
4. Spin Quantum Number (mₛ) – Describes the spin of an electron (+½ or -½).

9. What is the Heisenberg Uncertainty Principle?

Proposed by Werner Heisenberg, it states that it is impossible to simultaneously determine the exact position and momentum of an electron in an atom.

10. What is the difference between an orbit and an orbital?

• Orbit – A fixed path in which electrons revolve (Bohr’s Model).
• Orbital – A 3D region in space where the probability of finding an electron is highest (Quantum Mechanical Model).

11. What is the Aufbau Principle?

The Aufbau Principle states that electrons fill atomic orbitals in order of increasing energy levels, i.e., lower-energy orbitals are filled first before higher ones.

12. What is Hund’s Rule?

Hund’s Rule states that in degenerate (equal energy) orbitals, electrons fill each orbital singly before pairing up.

13. What is Pauli’s Exclusion Principle?

Pauli’s Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. This means an atomic orbital can hold a maximum of two electrons with opposite spins.

14. What is the Quantum Mechanical Model of the Atom?

Developed by Schrödinger, this model describes electrons as wave-like particles with a probability distribution around the nucleus instead of fixed orbits. It introduces atomic orbitals (s, p, d, f) as probable electron locations.

15. What are the types of orbitals in an atom?

There are four types of orbitals:

• s-orbital – Spherical shape (holds max 2 electrons).
• p-orbital – Dumbbell shape (holds max 6 electrons).
• d-orbital – Complex shape (holds max 10 electrons).
• f-orbital – Complex shape (holds max 14 electrons).

16. What is the significance of the atomic number and mass number?

• Atomic Number (Z) – Number of protons in an atom.
• Mass Number (A) – Sum of protons and neutrons in an atom’s nucleus.

17. What are Isotopes, Isobars, and Isotones?

• Isotopes – Same atomic number, different mass number (e.g., Hydrogen: ¹H, ²H, ³H).
• Isobars – Same mass number, different atomic number (e.g., ¹⁴C and ¹⁴N).
• Isotones – Same number of neutrons, different atomic and mass numbers (e.g., ¹⁴C and ¹⁵N).

18. What is the Dual Nature of Electrons?

Proposed by de Broglie, it states that electrons exhibit both particle and wave-like properties, known as wave-particle duality.

19. What is the Electronic Configuration of an Atom?

The electronic configuration of an atom describes how electrons are distributed in different orbitals. It follows the Aufbau principle, Hund’s rule, and Pauli’s exclusion principle (e.g., Oxygen (O) = 1s² 2s² 2p⁴).

20. Why is the Bohr model still used despite its limitations?

Although the Bohr model fails for multi-electron atoms and does not explain fine spectral lines, it is still useful because it provides a simple and understandable representation of electron energy levels.

📚 CBSE Class 11 Chemistry: Structure of Atom – Complete Syllabus Overview

Are you a Class 11 CBSE student aiming to master "Structure of Atom"? Here’s a comprehensive breakdown of the chapter to help you focus your studies!

Case Study-Based MCQs on Structure of Atom – Class 11 CBSE

Case study-based questions are designed to test your analytical skills and conceptual understanding. Below are five case studies with multiple-choice questions (MCQs) and detailed explanations.

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Case Study 1: Discovery of Atomic Structure

The structure of an atom has been explored through various experiments. J.J. Thomson discovered electrons using a cathode ray tube experiment, while Goldstein discovered protons using canal rays. Rutherford’s alpha particle scattering experiment led to the discovery of a dense nucleus at the center of the atom. Later, Bohr’s model refined our understanding by proposing that electrons move in discrete energy levels.

1.1 What was the key observation in Rutherford’s experiment?

A) Most alpha particles were deflected back
B) Alpha particles passed through the gold foil without deflection
C) All alpha particles got absorbed by the gold foil
D) Electrons were ejected from the gold foil

✅ Answer: B) Alpha particles passed through the gold foil without deflection

🔹 Explanation: Rutherford’s experiment showed that most of the space in an atom is empty, as most alpha particles passed straight through the foil. However, a few were deflected, indicating the presence of a dense, positively charged nucleus.

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1.2 What did J.J. Thomson’s experiment conclude?

A) Atoms have a nucleus
B) Atoms are indivisible
C) Atoms contain negatively charged particles
D) Atoms are mostly empty space

✅ Answer: C) Atoms contain negatively charged particles

🔹 Explanation: J.J. Thomson’s cathode ray experiment showed the presence of negatively charged electrons, leading to the "plum pudding model" of the atom.

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Case Study 2: Bohr’s Atomic Model

Niels Bohr proposed that electrons move in fixed orbits (energy levels) around the nucleus without losing energy. He introduced the quantization of energy levels and explained the emission spectra of hydrogen.

2.1 According to Bohr’s model, what happens when an electron jumps from a higher to a lower energy level?

A) The atom becomes unstable
B) The electron absorbs energy
C) The electron loses energy in the form of radiation
D) The electron disappears

✅ Answer: C) The electron loses energy in the form of radiation

🔹 Explanation: When an electron moves from a higher energy level to a lower one, it releases energy in the form of light (photon), which creates the atomic emission spectrum.

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2.2 Which of the following is NOT a postulate of Bohr’s model?

A) Electrons revolve around the nucleus in circular orbits
B) Electrons can have any energy within an orbit
C) Energy levels are quantized
D) Electrons do not lose energy while revolving

✅ Answer: B) Electrons can have any energy within an orbit

🔹 Explanation: According to Bohr, electrons can only exist in specific, quantized energy levels, meaning they cannot have any arbitrary energy.

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Case Study 3: Quantum Mechanical Model

The quantum mechanical model, developed by Schrodinger, replaced Bohr’s model. It introduced the concept of orbitals, where electrons are found as a probability distribution rather than fixed orbits. The model is based on wave-particle duality and Heisenberg’s Uncertainty Principle.

3.1 What does Heisenberg’s Uncertainty Principle state?

A) Electrons are present in fixed orbits
B) We cannot determine both position and momentum of an electron simultaneously
C) Electrons revolve around the nucleus like planets
D) Electrons follow a predictable path

✅ Answer: B) We cannot determine both position and momentum of an electron simultaneously

🔹 Explanation: Heisenberg’s Uncertainty Principle states that it is impossible to simultaneously know the exact position and momentum of an electron, leading to the concept of probability orbitals.

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3.2 Which scientist introduced the wave equation to describe electron behavior?

A) Bohr
B) Heisenberg
C) Schrodinger
D) Rutherford

✅ Answer: C) Schrodinger

🔹 Explanation: Erwin Schrodinger developed the wave equation, which describes the probability of finding an electron in a given region of space, known as an orbital.

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Case Study 4: Quantum Numbers

Each electron in an atom is described by four quantum numbers:

1. Principal quantum number (n) → Determines the energy level
2. Azimuthal quantum number (l) → Determines the shape of orbitals
3. Magnetic quantum number (m) → Determines the orientation of orbitals
4. Spin quantum number (s) → Determines electron spin

4.1 What does the principal quantum number (n) determine?

A) Shape of the orbital
B) Orientation of the orbital
C) Size and energy of the orbital
D) Spin of the electron

✅ Answer: C) Size and energy of the orbital

🔹 Explanation: The principal quantum number (n) determines the size of the electron cloud and the energy level where the electron resides.

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4.2 How many orbitals are present in the third energy level (n = 3)?

A) 3
B) 9
C) 18
D) 5

✅ Answer: B) 9

🔹 Explanation: The number of orbitals in an energy level is given by n².
For n = 3 → 3² = 9 orbitals (1s, 3p, 5d).

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Case Study 5: Electronic Configuration and Periodicity

The electronic configuration of an atom follows the Aufbau principle, Pauli Exclusion Principle, and Hund’s Rule. Elements in the periodic table are arranged based on their atomic number and valence electron configuration.

5.1 Which rule states that electrons fill orbitals in order of increasing energy?

A) Hund’s Rule
B) Aufbau Principle
C) Pauli’s Exclusion Principle
D) Heisenberg’s Principle

✅ Answer: B) Aufbau Principle

🔹 Explanation: The Aufbau Principle states that electrons occupy the lowest energy orbitals first before filling higher ones.

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5.2 The electronic configuration of an element is 1s² 2s² 2p⁶ 3s². What is the element?

A) Magnesium (Mg)
B) Sodium (Na)
C) Aluminium (Al)
D) Oxygen (O)

✅ Answer: A) Magnesium (Mg)

🔹 Explanation:

• 1s² 2s² 2p⁶ 3s² corresponds to an atomic number of 12, which is magnesium (Mg).

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CBSE Board Exam Syllabus for Chapter 2: Structure of Atom

Key Topics

1. Discovery of Subatomic Particles
   • Electron, Proton, and Neutron
   • Experiments by J.J. Thomson and Rutherford
2. Atomic Models
   • Thomson’s Model – The "Plum Pudding" model
   • Rutherford’s Nuclear Model – Gold foil experiment
   • Bohr’s Model of the Hydrogen Atom
3. Dual Nature of Matter and Radiation
   • Photoelectric Effect – Einstein’s explanation
   • de Broglie’s Hypothesis – Matter waves
4. Heisenberg’s Uncertainty Principle
   • Limitation in determining position and momentum simultaneously
5. Quantum Mechanical Model of Atom
   • Introduction to Schrödinger’s Wave Equation
   • Concept of orbitals and shapes of s, p, and d orbitals
6. Quantum Numbers
   • Principal (n), Azimuthal (l), Magnetic (m), and Spin (s)
   • Significance and rules of electron filling
7. Electronic Configuration of Atoms
   • Aufbau Principle – Electrons fill lower energy orbitals first
   • Pauli’s Exclusion Principle – No two electrons can have identical quantum numbers
   • Hund’s Rule of Maximum Multiplicity
8. Hydrogen Spectrum
   • Explanation of line spectra based on Bohr’s theory
9. Limitations of Bohr’s Model
   • Transition to modern quantum mechanics

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Tips to Master The Chapter Structure of Atom

• Understand and visualize atomic models and orbitals for better retention.
• Focus on the photoelectric effect and Heisenberg’s principle, as they are conceptually important.
• Practice writing electronic configurations using rules like the Aufbau Principle and Hund’s Rule.
• Solve questions on quantum numbers to strengthen your grasp of the concept.

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💡 Structure of Atom forms the foundation for further chapters like Chemical Bonding and Periodic Properties. Mastering it will make future topics easier to understand.

Happy studying, and good luck with your preparations! 😊

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CBSE Class 11 Chemistry Syllabus

CBSE Class 11 Chemistry Syllabus is a vast which needs a clear understanding of the concepts and topics. Knowing CBSE Class 11 Chemistry syllabus helps students to understand the course structure of Chemistry.

Unit-wise CBSE Class 11 Syllabus for Chemistry

Below is a list of detailed information on each unit for Class 11 Students.

UNIT I – Some Basic Concepts of Chemistry

General Introduction: Importance and scope of Chemistry.

Nature of matter, laws of chemical combination, Dalton’s atomic theory: concept of elements,
atoms and molecules.

Atomic and molecular masses, mole concept and molar mass, percentage composition, empirical and molecular formula, chemical reactions, stoichiometry and calculations based on stoichiometry.

UNIT II – Structure of Atom

Discovery of Electron, Proton and Neutron, atomic number, isotopes and isobars. Thomson’s model and its limitations. Rutherford’s model and its limitations, Bohr’s model and its limitations, concept of shells and subshells, dual nature of matter and light, de Broglie’s relationship, Heisenberg uncertainty principle, concept of orbitals, quantum numbers, shapes of s, p and d orbitals, rules for filling electrons in orbitals – Aufbau principle, Pauli’s exclusion principle and Hund’s rule, electronic configuration of atoms, stability of half-filled and completely filled orbitals.

UNIT III – Classification of Elements and Periodicity in Properties

Significance of classification, brief history of the development of periodic table, modern periodic law and the present form of periodic table, periodic trends in properties of elements -atomic radii, ionic radii, inert gas radii, Ionization enthalpy, electron gain enthalpy, electronegativity, valency. Nomenclature of elements with atomic number greater than 100.

UNIT IV – Chemical Bonding and Molecular Structure

Valence electrons, ionic bond, covalent bond, bond parameters, Lewis structure, polar character of covalent bond, covalent character of ionic bond, valence bond theory, resonance, geometry of covalent molecules, VSEPR theory, concept of hybridization, involving s, p and d orbitals and shapes of some simple molecules, molecular orbital theory of homonuclear diatomic molecules(qualitative idea only), Hydrogen bond.

UNIT V – Chemical Thermodynamics

Concepts of System and types of systems, surroundings, work, heat, energy, extensive and intensive properties, state functions. First law of thermodynamics – internal energy and enthalpy, measurement of U and H, Hess’s law of constant heat summation, enthalpy of bond dissociation, combustion, formation, atomization, sublimation, phase transition, ionization, solution and dilution. Second law of Thermodynamics (brief introduction)
Introduction of entropy as a state function, Gibb’s energy change for spontaneous and nonspontaneous processes.
Third law of thermodynamics (brief introduction).

UNIT VI – Equilibrium

Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of mass action, equilibrium constant, factors affecting equilibrium – Le Chatelier’s principle, ionic equilibrium- ionization of acids and bases, strong and weak electrolytes, degree of ionization,
ionization of poly basic acids, acid strength, concept of pH, hydrolysis of salts (elementary idea), buffer solution, Henderson Equation, solubility product, common ion effect (with illustrative examples).

UNIT VII – Redox Reactions

Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, in terms of loss and gain of electrons and change in oxidation number, applications of redox reactions.

UNIT VIII – Organic Chemistry: Some basic Principles and Techniques

General introduction, classification and IUPAC nomenclature of organic compounds. Electronic displacements in a covalent bond: inductive effect, electromeric effect, resonance and hyper conjugation. Homolytic and heterolytic fission of a covalent bond: free radicals, carbocations, carbanions, electrophiles and nucleophiles, types of organic reactions.

UNIT IX – Hydrocarbons

Classification of Hydrocarbons
Aliphatic Hydrocarbons:
Alkanes – Nomenclature, isomerism, conformation (ethane only), physical properties, chemical reactions.
Alkenes – Nomenclature, structure of double bond (ethene), geometrical isomerism, physical properties, methods of preparation, chemical reactions: addition of hydrogen, halogen, water, hydrogen halides (Markovnikov’s addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition.
Alkynes – Nomenclature, structure of triple bond (ethyne), physical properties, methods of preparation, chemical reactions: acidic character of alkynes, addition reaction of – hydrogen, halogens, hydrogen halides and water.

Aromatic Hydrocarbons:

Introduction, IUPAC nomenclature, benzene: resonance, aromaticity, chemical properties: mechanism of electrophilic substitution. Nitration, sulphonation, halogenation, Friedel Craft’s alkylation and acylation, directive influence of functional group in monosubstituted benzene. Carcinogenicity and toxicity.

To know the CBSE Syllabus for all the classes from 1 to 12, visit the Syllabus page of CBSE. Meanwhile, to get the Practical Syllabus of Class 11 Chemistry, read on to find out more about the syllabus and related information in this page.

CBSE Class 11 Chemistry Practical Syllabus with Marking Scheme

In Chemistry subject, practical also plays a vital role in improving their academic scores in the subject. The overall weightage of Chemistry practical mentioned in the CBSE Class 11 Chemistry syllabus is 30 marks. So, students must try their best to score well in practicals along with theory. It will help in increasing their overall academic score.

CBSE Class 11 Chemistry Practical Syllabus

The experiments will be conducted under the supervision of subject teacher. CBSE Chemistry Practicals is for 30 marks. This contribute to the overall practical marks for the subject.

The table below consists of evaluation scheme of practical exams.

Evaluation Scheme	Marks
Volumetric Analysis	08
Salt Analysis	08
Content Based Experiment	06
Project Work	04
Class record and viva	04
Total	30

CBSE Syllabus for Class 11 Chemistry Practical

Micro-chemical methods are available for several of the practical experiments. Wherever possible such techniques should be used.

A. Basic Laboratory Techniques
1. Cutting glass tube and glass rod
2. Bending a glass tube
3. Drawing out a glass jet
4. Boring a cork

B. Characterization and Purification of Chemical Substances
1. Determination of melting point of an organic compound.
2. Determination of boiling point of an organic compound.
3. Crystallization of impure sample of any one of the following: Alum, Copper Sulphate, Benzoic Acid.

C. Experiments based on pH

1. Any one of the following experiments:

• Determination of pH of some solutions obtained from fruit juices, solution of known and varied concentrations of acids, bases and salts using pH paper or universal indicator.
• Comparing the pH of solutions of strong and weak acids of same concentration.
• Study the pH change in the titration of a strong base using universal indicator.

2. Study the pH change by common-ion in case of weak acids and weak bases.

D. Chemical Equilibrium
One of the following experiments:

1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of either of the ions.
2. Study the shift in equilibrium between [Co(H₂O)₆] ²⁺ and chloride ions by changing the concentration of either of the ions.

E. Quantitative Estimation
i. Using a mechanical balance/electronic balance.
ii. Preparation of standard solution of Oxalic acid.
iii. Determination of strength of a given solution of Sodium hydroxide by titrating it against standard solution of Oxalic acid.
iv. Preparation of standard solution of Sodium carbonate.
v. Determination of strength of a given solution of hydrochloric acid by titrating it against standard Sodium Carbonatesolution.

F. Qualitative Analysis
1) Determination of one anion and one cation in a given salt
Cations‐ Pb²⁺, Cu²⁺, As³⁺, Al³⁺, Fe³⁺, Mn²⁺, Ni²⁺, Zn²⁺, Co²⁺, Ca²⁺, Sr²⁺, Ba²⁺, Mg²⁺, NH₄ ⁺
Anions – (CO₃)^2‐ , S^2‐, NO₂^‐ , SO₃^2‐, SO^2‐ , NO ^‐ , Cl^‐ , Br^‐, I‐, PO₄^3‐ , C₂O²‐ ,CH₃COO^‐
(Note: Insoluble salts excluded)

2) Detection of ‐ Nitrogen, Sulphur, Chlorine in organic compounds.

G) PROJECTS
Scientific investigations involving laboratory testing and collecting information from other sources.

A few suggested projects are as follows:

• Checking the bacterial contamination in drinking water by testing sulphide ion
• Study of the methods of purification of water.
• Testing the hardness, presence of Iron, Fluoride, Chloride, etc., depending upon the regional
  variation in drinking water and study of causes of presence of these ions above permissible
  limit (if any).
• Investigation of the foaming capacity of different washing soaps and the effect of addition of
  Sodium carbonate on it.
• Study the acidity of different samples of tea leaves.
• Determination of the rate of evaporation of different liquids Study the effect of acids and
  bases on the tensile strength of fibres.
• Study of acidity of fruit and vegetable juices.

Note: Any other investigatory project, which involves about 10 periods of work, can be chosen with the approval of the teacher.

Practical Examination for Visually Impaired Students of Class 11

Below is a list of practicals for the visually impaired students.

A. List of apparatus for identification for assessment in practicals (All experiments)
Beaker, tripod stand, wire gauze, glass rod, funnel, filter paper, Bunsen burner, test tube, test tube stand,
dropper, test tube holder, ignition tube, china dish, tongs, standard flask, pipette, burette, conical flask, clamp
stand, dropper, wash bottle
• Odour detection in qualitative analysis
• Procedure/Setup of the apparatus

B. List of Experiments A. Characterization and Purification of Chemical Substances
1. Crystallization of an impure sample of any one of the following: copper sulphate, benzoic acid
B. Experiments based on pH
1. Determination of pH of some solutions obtained from fruit juices, solutions of known and varied
concentrations of acids, bases and salts using pH paper
2. Comparing the pH of solutions of strong and weak acids of same concentration.

C. Chemical Equilibrium
1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing
the concentration of eitherions.
2. Study the shift in equilibrium between [Co(H2O)6]2+ and chloride ions by changing the
concentration of either of the ions.

D. Quantitative estimation
1. Preparation of standard solution of oxalic acid.
2. Determination of molarity of a given solution of sodium hydroxide by titrating it against standard
solution of oxalic acid.

E. Qualitative Analysis
1. Determination of one anion and one cation in a given salt
2. Cations – NH+4
Anions – (CO3)2-, S2-, (SO3)2-, Cl-, CH3COO-
(Note: insoluble salts excluded)
3. Detection of Nitrogen in the given organic compound.
4. Detection of Halogen in the given organic compound.

Note: The above practicals may be carried out in an experiential manner rather than recording observations.

We hope students must have found this information on CBSE Syllabus useful for their studying Chemistry. Learn Maths & Science in interactive and fun loving ways with ANAND CLASSES (A School Of Competitions) App/Tablet.

Frequently Asked Questions on CBSE Class 11 Chemistry Syllabus

Q1

How many units are in the CBSE Class 11 Chemistry Syllabus?

There are 9 units in the CBSE Class 11 Chemistry Syllabus. Students can access various study materials for the chapters mentioned in this article for free at ANAND CLASSES (A School Of Competitions).

Q2

What is the total marks for practicals examination as per the CBSE Class 11 Chemistry Syllabus?

The total marks for the practicals as per the CBSE Class 11 Chemistry Syllabus is 30. It includes volumetric analysis, content-based experiment, salt analysis, class record, project work and viva.

Q3

Which chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry?

The organic chemistry chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry.