NCERT Notes Class 11 Chemistry (Part-I) Chapter 2: Structure of Atom (Free PDF)

Atoms are considered the building blocks of matter. Everything we observe in the physical world is composed of atoms. In an earlier chapter, we explored the physical world and the concept of matter. Now, this unit on the structure of the atom will provide a deeper understanding of these building blocks of matter, including their properties, various atomic theories, quantum theories of the atom, and the principles of electronic configuration.

Explore Notes of Class 11 Chemistry

Chapter 1

Chapter 3

Chapter 4

Chapter 5

Chapter 6

Download PDF of NCERT Notes Class 11 Chemistry Chapter 2: Structure of Atom

Discovery of Sub-Atomic Particles

Until the late 19th century, atoms were believed to be indivisible. However, experiments on electric discharge through gases revealed the existence of fundamental particles — electrons, protons, and neutrons, which are now known as sub-atomic particles.

Discovery of the Electron

The discovery of the electron is discussed below.

• Experiment
  In the mid-1850s, discharge tube experiments (by J.J. Thomson and others) were conducted using a glass tube with electrodes at low pressure.
• Observations
  • A stream of rays emerged from the cathode (negative electrode) and moved towards the anode.
  • These rays produced a greenish glow on the glass coated with fluorescent material (zinc sulphide).
  • When a perforated anode was used, rays passed through and created a spot on the glass behind.
• Characteristics of Cathode Rays
  • Travel in straight lines.
  • Produce mechanical and heating effects.
  • They are deflected by electric and magnetic fields.
  • Carry a negative charge and mass.
    
  • Independent of the nature of the gas and the electrodes.
• Conclusion
  Cathode rays are made of negatively charged particles, later named electrons by J.J. Thomson in 1897.

Charge to Mass Ratio of an Electron

Measured by J.J. Thomson in 1897 using electric and magnetic fields. Charge to mass ratio of an electron:  e/m = 1.758 × 10¹¹ C/kg. This value remained the same for all gases and cathodes used, proving that electrons are present in all atoms.

Charge on the Electron

Below, we have provided a summary of information on the charge on the electron.

• Experiment: Millikan’s Oil Drop Experiment (1906–14)
• Principle
  • Oil droplets acquire a negative charge.
  • Their motion is observed under gravity and an electric field.
  • The charge on a single drop was measured.
• Result: Charge on electron, e = –1.602 × 10⁻¹⁹ coulomb
• Combining with Thomson’s e/m ratio: Mass of electron, m = e ÷ (e/m) = 9.109 × 10⁻³¹ kg

Discovery of Protons and Neutrons

The discovery of protons and neutrons is discussed below.

Protons (Discovered via Canal Rays)

• Positive rays observed in discharge tubes with perforated cathodes.
• Move in the opposite direction to the cathode rays.
• Properties depend on the gas used.
• From hydrogen, the lightest positive particle was obtained, named a proton.
• Charge of proton = +1.602 × 10⁻¹⁹ C
• Mass of proton ≈ 1.672 × 10⁻²⁷ kg

Neutrons (Discovered by James Chadwick, 1932):

• Beryllium bombarded with alpha particles released neutral radiation.
• These were named neutrons — particles with no charge but mass nearly equal to protons.
• Charge of neutron = 0
• Mass of neutron ≈ 1.675 × 10⁻²⁷ kg

Thomson’s Model of the Atom (1898)

Also known as the “Plum Pudding Model” or “Watermelon Model”

• The atom is a uniformly positive sphere with negatively charged electrons embedded in it.
• The positive charge balances the negative electrons so that the atom remains electrically neutral.
• Like seeds (electrons) embedded in a watermelon (positive sphere), or plums in a pudding.

Advantages:

• Explained the overall neutrality of atoms.
• Introduced the concept of the internal structure of atoms.

Drawbacks

• Could not explain the results of Rutherford’s alpha scattering experiment.
• Failed to explain the stability of atoms and spectral lines.

Also Read: Best Books for UPSC Optional Psychology

Rutherford’s Nuclear Model of the Atom (1911)

Based on the Alpha Particle Scattering Experiment

• Conducted by Hans Geiger and Ernest Marsden under Rutherford.
• Setup: A thin gold foil (~1000 atoms thick) bombarded with fast-moving alpha particles (He²⁺ ions).
• Observations:
  1. Most alpha particles passed through without deflection.
  2. Some were deflected at small angles.
  3. Very few (1 in 20,000) bounced back or deflected at large angles.

Conclusions

1. Most of the atoms have empty spaces, allowing particles to pass through.
2. Positive charge and most mass are concentrated in a tiny central region — the nucleus.
3. Electrons revolve around the nucleus like planets around the sun.

Postulates of Rutherford’s Model:

• An atom has a dense, positively charged nucleus.
• Electrons move around the nucleus in circular orbits.
• The size of the nucleus is very small compared to the size of the atom.

Limitations

• Could not explain why electrons do not spiral into the nucleus (classical theory suggests orbiting electrons should lose energy).
• No explanation for atomic spectra — failed to describe line spectra of hydrogen or other atoms.

Also Read: NCERT Notes Class 11 Geography Fundamentals of Geography Chapter 6 Landforms and their Evolution (Free PDF)

Atomic Number and Mass Number

Atomic number and mass number are discussed below.

Atomic Number (Z):

• Defined as the number of protons in the nucleus of an atom.
• It also equals the number of electrons in a neutral atom.
• Symbol: Z

Example: For carbon, Z = 6 (6 protons)

Mass Number (A)

• The sum of the number of protons and neutrons in the nucleus.
• Formula:
  • A = Z + Number of neutrons
• Example: For carbon-12, A = 12 (6 protons + 6 neutrons

Isobars and Isotopes

Isobars and isotopes are discussed below

Isotopes

• Atoms of the same element (same atomic number) but different mass numbers.
• Example: ₁H¹, ₁H² (deuterium), ₁H³ (tritium). 
• All are hydrogen isotopes with 1 proton, but different neutrons.

Isobars

• Atoms of different elements with the same mass number but different atomic numbers.
  Example:
• ₁₈Ar⁴⁰ and ₂₀Ca⁴⁰
• Both have a mass number of 40, but different Z.

Wave Nature of Electromagnetic Radiation

Electromagnetic radiation includes light, X-rays, UV, infrared, radio waves, etc. It travels through space as oscillating electric and magnetic fields perpendicular to each other and to the direction of wave motion.

Key Terms

1. Wavelength (λ)
   • Distance between two crests or troughs.
   • Unit: metre (m); commonly used: nanometre (nm), angstrom (Å).
   • 1 nm = 10⁻⁹ m, 1 Å = 10⁻¹⁰ m
2. Frequency (ν)
   • Number of wave cycles per second.
   • Unit: hertz (Hz) or s⁻¹
3. Velocity of light (c)
   • Speed of electromagnetic waves in vacuum.
   • c = 3.0 × 10⁸ m/s
4. Wave equation
   • c = ν × λ
5. Wavenumber (ν̅)
   • Number of wavelengths per unit length (usually per cm).
   • ν̅ = 1 / λ (in cm⁻¹)

Electromagnetic Spectrum

• The range of all types of electromagnetic radiation.
• Visible light is a small portion (400 nm to 750 nm).
• Shorter wavelength → higher frequency → more energy.

Planck’s Quantum Theory (1900)

• Energy is not continuous but emitted or absorbed in discrete packets called quanta.
• For light, each quantum is called a photon.
• Energy of one quantum: E = h × ν

Photoelectric Effect (Einstein’s Explanation, 1905)

When light of a certain frequency falls on a metal surface, electrons are ejected instantly.

Observations

1. Emission occurs only if the light frequency is above a threshold frequency (ν₀).
2. The number of electrons ejected depends on light intensity.
3. The kinetic energy of ejected electrons increases with increasing frequency.

Einstein’s Equation

K.E. = hν – hν₀

• hν = energy of incoming photon
• hν₀ = minimum energy to eject an electron (work function)

Bohr’s Model for the Hydrogen Atom

To overcome the limitations of Rutherford’s atomic model and explain hydrogen’s line spectrum, Niels Bohr in 1913 proposed a new model based on quantized electron orbits.

Postulates of Bohr’s Model

1. Electrons revolve around the nucleus in fixed orbits (called stationary states or energy levels) without radiating energy.
   • These orbits have fixed energy and are denoted as n = 1, 2, 3,… (called K, L, M shells).
2. The energy of an electron in an orbit is quantized.
   • As long as the electron stays in an orbit, its energy remains constant.
3. Energy is emitted or absorbed only when an electron jumps from one orbit to another.
   • If the electron jumps from a higher to a lower orbit → energy is emitted.
   • If the electron jumps from a lower to a higher orbit → energy is absorbed.
   • Energy absorbed or emitted:
     ΔE = E₂ – E₁ = hν

Expression for Radius of nth Orbit (rₙ)

For a hydrogen atom:
rₙ = 0.529 × n² Å
(where 1 Å = 10⁻¹⁰ m)

The smallest orbit (n = 1) has radius r₁ = 0.529 Å (called Bohr radius)

Expression for Velocity of Electron in nth Orbit (vₙ)

vₙ = 2.18 × 10⁶ × (1/n) m/s

Expression for Energy of Electron in nth Orbit (Eₙ):

For a hydrogen atom:
Eₙ = –13.6 eV / n²

• Energy is negative, meaning the electron is bound to the nucleus.
• As n increases, Eₙ becomes less negative (electron becomes less tightly bound).
• When n → ∞, E = 0 (electron is completely free → ionization).

Total Energy of the Electron

E = Kinetic Energy + Potential Energy = –13.6 eV / n²

Explanation of the Line Spectrum of Hydrogen

• When an excited hydrogen atom returns to a lower energy level, it emits radiation corresponding to the energy difference between the two levels.
• This energy is seen as a line in the hydrogen spectrum.

Frequency of emitted radiation

• ν = (E₂ – E₁) / h
• Or in terms of wavenumber:
• ν̅ = R × (1/n₁² – 1/n₂²)
  Where:
  • ν̅ = wavenumber in cm⁻¹
  • R = Rydberg constant = 109,677 cm⁻¹
  • n₂ > n₁

This formula explains the series of lines observed in the hydrogen spectrum:

• Lyman series: n₁ = 1 (ultraviolet region)
• Balmer series: n₁ = 2 (visible region)
• Paschen series: n₁ = 3 (infrared)
• Brackett series: n₁ = 4
• Pfund series: n₁ = 5

Dual Behavior of Matter

The dual behaviour of matter is discussed below.

de Broglie’s Hypothesis (1924)

• Proposed that matter (like electrons) also exhibits wave-particle duality, similar to light.
• A moving particle (e.g., an electron) has an associated wavelength.

de Broglie Equation

λ = h / (mv)
Where:

• λ = wavelength of matter wave (in meters)
• h = Planck’s constant = 6.626 × 10⁻³⁴ Js
• m = mass of particle (in kg)
• v = velocity of the particle (in m/s)

Heisenberg’s Uncertainty Principle (1927)

Proposed by Werner Heisenberg, it states that: It is impossible to simultaneously determine the exact position and exact momentum of a microscopic particle like an electron.

Mathematical Expression

Δx × Δp ≥ h / (4π)

Where:

• Δx = uncertainty in position
• Δp = uncertainty in momentum (mass × velocity)
• h = Planck’s constant

Implications:

• The more accurately we know the position, the less accurately we can know the momentum, and vice versa.
• Makes it impossible to define a precise path (orbit) for electrons.

Significance of ψ and ψ²

Quantum mechanics replaces fixed orbits with probability distributions.

Wave Function (ψ)

• Introduced by Erwin Schrödinger, ψ represents the amplitude of an electron wave.
• It has no direct physical meaning.

Probability Interpretation

• ψ² gives the probability density — the likelihood of finding an electron in a given region.
• The higher the value of ψ², the more probable the presence of an electron at that point.
• The region in space where the probability of finding the electron is maximum is called an orbital.

Also Read: NCERT Notes Class 11 Political Science Indian Constitution at Work Chapter 6: Judiciary (Free PDF)

Quantum Mechanical Model of the Atom

Developed using Schrödinger’s wave equation and the uncertainty principle, this model describes atoms using orbitals (regions of probability) rather than defined paths for electrons.

• Applicable to multi-electron systems.
• Describes electronic structure in terms of energy levels, sub-levels, and orbitals.

Orbitals and Quantum Numbers

An orbital is a region in space around the nucleus where the probability of finding an electron is maximum (90–95%). To specify the location and properties of an electron in an atom, four quantum numbers are used:

1. Principal Quantum Number (n)

• Represents the main energy level or shell.
• Values: n = 1, 2, 3, 4… (K, L, M, N…)
• Determines:
  • Size of the orbital
  • Energy of orbital
• Higher n → larger size → higher energy

2. Azimuthal Quantum Number (l)

• Also called the angular momentum quantum number
• Defines a subshell or shape of an orbital
• Values: l = 0 to (n – 1)
  
  • l = 0 → s-orbital
  • l = 1 → p-orbital
  • l = 2 → d-orbital
  • l = 3 → f-orbital

3. Magnetic Quantum Number (m)

• Describes the orientation of orbitals in space
• Values: m = –l to +l, including 0
• For example:
  • If l = 1 (p-orbital), m = –1, 0, +1 → 3 orientations (pₓ, pᵧ, p_z)

4. Spin Quantum Number (s)

• Describes the spin direction of an electron
• Values: +½ or –½
  • +½ → clockwise spin
  • –½ → anti-clockwise spin

Shapes of Orbitals

Orbitals have different shapes depending on the value of l:

s-Orbitals (l = 0)

• Shape: Spherical
• Electron density is the same in all directions
• Number of orientations: 1

p-Orbitals (l = 1)

• Shape: Dumbbell-shaped
• Three orientations: pₓ, pᵧ, p_z
• Located along the x, y, and z axes

d-Orbitals (l = 2)

• Shape: Complex, cloverleaf-like
• Five orientations: dₓᵧ, d_yz, d_zx, d_z², dₓ²–y²

f-Orbitals (l = 3)

• Shape: Highly complex
• Seven orientations

Filling of Orbitals in Atoms

The way electrons are distributed among various orbitals in atoms is governed by a set of principles and rules. This arrangement is called the electronic configuration of atoms.

Aufbau Principle

Aufbau is a German word meaning “building up.” Electrons are filled into orbitals in order of increasing energy, starting from the lowest energy orbital.

Key Rule (n + l Rule)

• The energy of an orbital depends on the sum of the principal quantum number (n) and the azimuthal quantum number (l).
• Orbitals with lower (n + l) are filled first.
• If two orbitals have the same (n + l), the one with lower n is filled first.

Order of Filling (Based on Energy Levels)

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s…

Pauli Exclusion Principle

Proposed by Wolfgang Pauli (1926). No two electrons in the same atom can have all four quantum numbers identical. An orbital can hold a maximum of two electrons with opposite spins. This explains why electrons pair with opposite spins in orbitals.

Hund’s Rule of Maximum Multiplicity

Electrons occupy degenerate orbitals (same energy) singly first, with parallel spins, before pairing begins.

Electronic Configuration of Atoms

Electronic configuration describes how electrons are distributed among the various shells and subshells in an atom.

Shell-wise vs Subshell-wise Distribution

Shell-wise configuration lists total electrons in each shell (K, L, M…). Subshell-wise configuration follows the order of filling based on the Aufbau principle and gives a more detailed view of electron arrangement.

Important Formulas in NCERT Notes Class 11 Chemistry (Part-I) Chapter 2: Structure of Atom

Here are the important formulas and equations from NCERT Class 11 Chemistry Chapter 2: Structure of Atom:

• Charge-to-mass ratio of an electron
  e/m = 1.758820 × 10¹¹ C/kg
• Charge on an electron
  e = 1.6022 × 10⁻¹⁹ C
• Speed of light and wavelength-frequency relation
  c = ν × λ
  (where c = 3.0 × 10⁸ m/s)
• Wavenumber
  ν̅ = 1 / λ
• Energy of a photon
  E = h × ν = (h × c) / λ
  (h = 6.626 × 10⁻³⁴ J·s)
• Energy of one mole of photons
  E = Nₐ × h × ν
  (Nₐ = 6.022 × 10²³ mol⁻¹)
• Photoelectric equation (Einstein’s equation)
  K.E. = hν − W₀
  (W₀ = work function of the metal)
• de Broglie wavelength
  λ = h / (m × v)
• Energy of an electron in the nth orbit (Bohr’s model)
  Eₙ = −2.18 × 10⁻¹⁸ / n² J
• Radius of the nth orbit
  rₙ = 0.529 × n² Å
• Wavenumber for spectral lines (Hydrogen-like atoms)
  ν̅ = R × (1/n₁² − 1/n₂²)
  (R = 1.097 × 10⁷ m⁻¹)
• Heisenberg’s uncertainty principle
  Δx × Δp ≥ h / (4π)

Explore Notes of Class 11 Chemistry

Chapter 1

Chapter 3

Chapter 4

Chapter 5

Chapter 6

Download the Solutions of Other Chapters of Class 11 Chemistry

Chapter 1

Chapter 2

Chapter 3

Chapter 4

Chapter 5

Related Reads

NCERT Solutions Class 8 Geography	NCERT Notes Class 11 English Snapshots Chapter: 3 Mother’s Day (Free PDF)
NCERT Notes Class 11 Sociology Understanding Society Chapter 4: Introducing	NCERT Solutions Class 11 Sociology Understanding Society Chapter 4: Introducing Western Sociologists (Free PDF)
Western Sociologists (Free PDF) NCERT Solutions Class 11 Sociology Understanding Society Chapter 4: Introducing Western Sociologists (Free PDF) Best Books for UPSC Optional Psychology	NCERT Solutions Class 11 English Snapshots Chapter: 3 Mother’s Day (Free PDF)

Source: Ashu Ghai 11 & 12

Explore notes on other subjects in the NCERT Class 11

English

Sociology

Geography

Political Science

Psychology

For more topics, follow LeverageEdu NCERT Study Material today!