CBSE Class 11 Chemistry Revision Notes Chapter 2 Structure of Atom

CBSE Class 11 Chemistry Revision Notes Chapter 2 explain Structure of Atom through subatomic particles, atomic models, spectra and quantum numbers. For CBSE 2026 Chemistry, Structure of Atom builds the base for electronic configuration, orbitals and periodic properties.

Why did Dalton’s solid indivisible atom fail after cathode rays, canal rays and radioactivity experiments? The NCERT Class 11 Chemistry chapter Structure of Atom begins with this change in scientific thinking. Dalton’s theory explained laws of chemical combination, but it could not explain electrical charge, atomic spectra or the presence of electrons, protons and neutrons inside atoms.

The chapter then moves from discovery of subatomic particles to atomic models. Thomson explained neutrality, Rutherford discovered the nucleus, and Bohr explained hydrogen’s line spectrum. Later ideas such as wave-particle duality, Heisenberg uncertainty principle and Schrödinger’s quantum model replaced fixed electron paths with orbitals, probability and quantum numbers. These ideas help students write electronic configurations and understand atomic structure in CBSE 2026 Chemistry.

Key Takeaways

• Electron charge: Millikan’s oil drop experiment gave electron charge as −1.602176 × 10⁻¹⁹ C.
• Electron mass: The mass of an electron is 9.1094 × 10⁻³¹ kg.
• Nuclear size: Rutherford’s model gives atomic radius near 10⁻¹⁰ m and nuclear radius near 10⁻¹⁵ m.
• Photon energy: Planck’s quantum theory uses E = hν, where h = 6.626 × 10⁻³⁴ J s.

CBSE Class 11 Chemistry Revision Notes Chapter 2 Structure 2026

Subatomic Particles and the Failure of Dalton’s Atom

Cathode ray and canal ray experiments showed that atoms contain charged particles. This changed Dalton’s idea that atoms are indivisible particles of matter.

In CBSE Class 11 Chemistry Chapter 2 Structure of Atom, subatomic particles form the first base for atomic models. Electrons, protons and neutrons explain charge, mass and nuclear structure.

Discovery of electron

Cathode rays start from the cathode and move towards the anode. They travel in straight lines when no electric or magnetic field is applied.

In electric or magnetic fields, cathode rays behave like negatively charged particles. These particles were called electrons.

The characteristics of cathode rays do not depend on electrode material or gas inside the tube. This showed that electrons are basic constituents of all atoms.

Charge to mass ratio of electron

J.J. Thomson measured the charge to mass ratio of the electron in 1897. He used electric and magnetic fields in a cathode ray tube.

Formula value:

e/me = 1.758820 × 10¹¹ C kg⁻¹

The deflection of electrons depends on charge, mass and field strength. Lighter particles deflect more under the same field.

Charge and mass of electron

R.A. Millikan determined the charge on the electron through the oil drop experiment. The accepted charge is −1.602176 × 10⁻¹⁹ C.

Using Thomson’s e/me value and Millikan’s charge, electron mass was calculated.

Electron mass:

me = 9.1094 × 10⁻³¹ kg

Discovery of proton and neutron

Canal ray experiments led to the discovery of positively charged particles. The lightest positive ion came from hydrogen and was called proton.

Chadwick discovered neutron in 1932 by bombarding beryllium with alpha particles. Neutrons are electrically neutral and have mass slightly greater than protons.

Thomson and Rutherford Atomic Models

Thomson and Rutherford gave two early atomic models after the discovery of subatomic particles. Thomson explained atomic neutrality, while Rutherford explained the nuclear structure.

Class 11 chemistry chapter 2 structure of atom notes usually test these models through features, observations and drawbacks. The gold foil experiment is the key evidence for Rutherford’s model.

Thomson model of atom

J.J. Thomson proposed that an atom is a positively charged sphere with electrons embedded in it. The atom remains neutral because positive and negative charges balance.

This model is also called the plum pudding model, raisin pudding model or watermelon model. It assumed that mass and positive charge were uniformly spread in the atom.

Limitation of Thomson model

Thomson’s model could explain the neutrality of the atom. It could not explain later experimental results from alpha particle scattering.

The model failed because positive charge is not spread throughout the atom. Rutherford’s experiment showed that positive charge is concentrated in a tiny nucleus.

Rutherford alpha particle scattering experiment

Rutherford directed high-energy alpha particles at a thin gold foil. A fluorescent zinc sulphide screen detected the scattered particles.

The observations were:

1. Most alpha particles passed through undeflected.
2. Some particles deflected through small angles.
3. Very few particles bounced back by nearly 180°.

Rutherford nuclear model

Rutherford concluded that most of the atom is empty space. The positive charge and most atomic mass are concentrated in a tiny nucleus.

Electrons move around the nucleus in circular paths. Electrostatic attraction holds electrons and nucleus together.

Drawbacks of Rutherford model

Rutherford’s model could not explain atomic stability. According to classical theory, an accelerating electron should emit radiation and fall into the nucleus.

Calculations suggested that an electron would spiral into the nucleus in about 10⁻⁸ s. This does not happen in real atoms.

Rutherford’s model also failed to explain electron distribution and electron energies. Bohr’s model tried to solve this problem for hydrogen.

Atomic Number, Mass Number, Isotopes and Isobars

Atomic number is the number of protons in the nucleus. In a neutral atom, it is also equal to the number of electrons.

Mass number is the total number of protons and neutrons. These two numbers help write isotope and isobar symbols in Structure of Atom class 11 notes.

Atomic number and mass number

Atomic number is represented by Z.

Formula:

Z = Number of protons

For a neutral atom:

Z = Number of electrons

Mass number is represented by A.

Formula:

A = Number of protons + Number of neutrons

Number of neutrons = A − Z

Isotopes

Isotopes have the same atomic number but different mass numbers. They differ because they have different numbers of neutrons.

Examples:

¹H, ²D and ³T are isotopes of hydrogen.

¹²C, ¹³C and ¹⁴C are isotopes of carbon.

Isotopes of the same element show similar chemical behaviour because they have the same number of electrons.

Isobars

Isobars have the same mass number but different atomic numbers.

Examples:

¹⁴C and ¹⁴N are isobars.

They belong to different elements because their proton numbers are different.

Electromagnetic Radiation and Planck’s Quantum Theory

Electromagnetic radiation has electric and magnetic fields that oscillate perpendicular to each other. These waves do not need a medium and can travel through vacuum.

The NCERT Class 11 Chemistry chapter Structure of Atom uses radiation to explain spectra, photoelectric effect and Bohr’s model. Frequency, wavelength and energy become important formula terms.

Wavelength, frequency and speed of light

Electromagnetic radiation is described by wavelength and frequency. Wavelength is represented by λ, and frequency is represented by ν.

Formula:

c = νλ

Where:

c = 3.0 × 10⁸ m s⁻¹

ν = frequency

λ = wavelength

Wavenumber is the number of wavelengths per unit length. Its common unit is cm⁻¹.

Planck’s quantum theory

Planck proposed that energy is emitted or absorbed in small packets called quanta. The energy of one quantum depends on frequency.

Formula:

E = hν

Where:

E = energy

h = 6.626 × 10⁻³⁴ J s

ν = frequency

Energy cannot take every possible value in this model. It changes in discrete amounts.

Photoelectric effect

Photoelectric effect is the ejection of electrons from a metal surface when light of sufficient frequency falls on it. Hertz observed this phenomenon in 1887.

Einstein explained it using Planck’s quantum theory. Light behaves as photons when it interacts with matter.

Important facts:

1. Electrons are ejected without time lag.
2. Number of electrons depends on light intensity.
3. Kinetic energy depends on frequency.
4. Each metal has a threshold frequency.

Formula:

Kinetic energy = hν − hν0

Atomic Spectra and Hydrogen Line Spectrum

Atomic spectra show that atoms absorb or emit only specific frequencies of radiation. This gave evidence for quantised electronic energy levels.

In CBSE Class 11 Chemistry Chapter 2 Structure of Atom, spectra connect electromagnetic radiation with electron transitions. Hydrogen spectrum is the main example because hydrogen has only one electron.

Emission spectrum

An emission spectrum is produced when an excited atom emits radiation while returning to a lower energy state. Gas-phase atoms give line spectra.

Each element has a unique line emission spectrum. This helps identify unknown elements by spectroscopic methods.

Absorption spectrum

An absorption spectrum is formed when a sample absorbs specific wavelengths from continuous radiation. The missing wavelengths appear as dark lines.

An absorption spectrum is like the photographic negative of an emission spectrum. Both show specific energy transitions.

Hydrogen spectrum series

Hydrogen spectrum has several line series. These series correspond to electron transitions ending at different energy levels.

Important series:

1. Lyman series: n1 = 1, ultraviolet region
2. Balmer series: n1 = 2, visible region
3. Paschen series: n1 = 3, infrared region
4. Brackett series: n1 = 4, infrared region
5. Pfund series: n1 = 5, infrared region

Rydberg expression

The hydrogen spectral lines are expressed using the Rydberg formula.

Formula:

ν̅ = 109677 × (1/n1² − 1/n2²) cm⁻¹

Where:

ν̅ = wavenumber

n1 = lower energy level

n2 = higher energy level

n2 is greater than n1.

Bohr Model of Atom

Bohr explained hydrogen’s stability and line spectrum using quantised orbits. Bohr model of atom class 11 is mainly applied to hydrogen and hydrogen-like species.

Bohr’s model improved Rutherford’s model by giving fixed energy states. It also connected spectral lines with transitions between those states.

Postulates of Bohr model

Bohr proposed that an electron moves around the nucleus in circular paths of fixed radius and energy. These paths are called orbits or stationary states.

The energy of an electron does not change while it stays in a stationary orbit. Energy is absorbed or emitted only when the electron jumps between two orbits.

Angular momentum of an electron is quantised.

Formula:

mevr = nh/2π

Where:

me = mass of electron

v = velocity

r = orbit radius

n = principal quantum number

Radius and energy of hydrogen orbit

The radius of the nth orbit is:

rn = n²a0

Where:

a0 = 52.9 pm

The energy of the nth orbit of hydrogen is:

En = −RH × (1/n²)

Where:

RH = 2.18 × 10⁻¹⁸ J

The negative sign means the electron is bound to the nucleus. The most stable state is n = 1.

Bohr frequency rule

When an electron jumps between two stationary states, radiation is absorbed or emitted.

Formula:

ΔE = E2 − E1 = hν

If the electron moves to a higher energy level, energy is absorbed. If it moves to a lower energy level, energy is emitted.

Limitations of Bohr model

Bohr’s model explains hydrogen and hydrogen-like ions such as He⁺, Li²⁺ and Be³⁺. It fails for atoms with more than one electron.

It also fails to explain fine spectral details, Zeeman effect and Stark effect. It cannot explain chemical bonding.

Dual Behaviour of Matter and Heisenberg Uncertainty Principle

De Broglie proposed that matter shows both particle and wave-like behaviour. This idea applied wave character to electrons and other small particles.

Heisenberg uncertainty principle class 11 explains why fixed electron paths cannot exist. It also gives the reason Bohr’s circular orbit idea fails at the microscopic level.

De Broglie relation

De Broglie gave a relation between wavelength and momentum of a moving particle.

Formula:

λ = h/mv

Where:

λ = wavelength

h = Planck’s constant

m = mass

v = velocity

The wavelength of ordinary objects is too small to detect. Electron wavelength can be detected experimentally.

Heisenberg uncertainty principle

Heisenberg uncertainty principle states that exact position and exact momentum of an electron cannot be determined simultaneously.

Formula:

Δx × Δp ≥ h/4π

Or:

Δx × mΔv ≥ h/4π

If position is measured very accurately, velocity becomes highly uncertain. This makes a fixed path of electron impossible.

Why Bohr model fails

Bohr’s model treats an electron as a particle moving in a definite circular orbit. This ignores the wave nature of the electron.

A definite orbit also needs exact position and exact velocity at the same time. This contradicts Heisenberg uncertainty principle.

Quantum Mechanical Model of Atom

Quantum mechanical model of atom uses wave-particle duality and uncertainty principle. It replaces fixed electron orbits with orbitals and probability.

Quantum mechanics was developed independently by Werner Heisenberg and Erwin Schrödinger in 1926. Schrödinger’s equation gives possible energies and wave functions for electrons.

Meaning of orbital

An atomic orbital is the one-electron wave function ψ in an atom. It is characterised by three quantum numbers: n, l and ml.

The wave function ψ has no direct physical meaning. The square of the wave function, |ψ|², gives probability density.

Orbit and orbital difference

An orbit is a fixed circular path around the nucleus in Bohr’s model. An orbital is a region where the probability of finding an electron is high.

Bohr orbits have no real meaning under the uncertainty principle. Orbitals are part of the quantum mechanical model.

Features of quantum mechanical model

Important features are:

1. Electron energy in atoms is quantised.
2. Quantised energy levels arise from wave-like properties of electrons.
3. Exact position and velocity cannot be known together.
4. An orbital cannot contain more than two electrons.
5. Probability of finding an electron is proportional to |ψ|².

Quantum Numbers in Class 11 Chemistry

Quantum numbers describe the size, shape, orientation and spin of an electron in an atom. Each orbital is identified by n, l and ml.

Quantum numbers class 11 chemistry questions often ask allowed values and number of orbitals. The spin quantum number ms describes electron spin orientation.

Principal quantum number

Principal quantum number is represented by n. It describes the main shell and energy level.

Allowed values:

n = 1, 2, 3, 4...

A larger n means a larger orbital size and higher energy. The total number of orbitals in a shell is n².

Azimuthal quantum number

Azimuthal quantum number is represented by l. It describes the subshell and shape of the orbital.

Allowed values:

l = 0 to n − 1

Subshell symbols:

l = 0 means s

l = 1 means p

l = 2 means d

l = 3 means f

Magnetic quantum number

Magnetic quantum number is represented by ml. It describes orbital orientation in space.

Allowed values:

ml = −l to +l

For p subshell, l = 1.

So, ml = −1, 0, +1

This gives three p orbitals.

Spin quantum number

Spin quantum number is represented by ms. It describes the spin orientation of the electron.

Allowed values:

ms = +1/2 or −1/2

Two electrons in the same orbital must have opposite spins.

Shapes, Nodes and Energies of Orbitals

Orbital shape depends on the azimuthal quantum number. The s orbitals are spherical, while p orbitals are dumb-bell shaped.

The probability distribution of an electron also shows nodes. Nodes are regions where probability density becomes zero.

Shapes of s, p and d orbitals

s orbitals are spherical and non-directional. Their size increases from 1s to 2s to 3s.

p orbitals have two lobes and three orientations: px, py and pz. These three p orbitals have the same energy in hydrogen atom.

d orbitals have five orientations. Their shapes are more complex than s and p orbitals.

Nodes

Nodes are regions where probability density is zero. For ns orbitals, the number of nodes is n − 1.

Examples:

1s orbital has 0 nodes.

2s orbital has 1 node.

3s orbital has 2 nodes.

Energy of orbitals

In hydrogen atom, orbital energy depends only on n. In multi-electron atoms, orbital energy depends on both n and l.

For multi-electron atoms, the energy order is decided by the n + l rule. Lower n + l value means lower energy.

If two orbitals have the same n + l value, the orbital with lower n has lower energy.

Electronic Configuration of Atoms

Electronic configuration shows how electrons are arranged in orbitals. It follows Aufbau principle, Pauli exclusion principle and Hund’s rule.

Electronic configuration class 11 questions test the filling order of orbitals and the distribution of electrons. Correct configuration also explains valence electrons.

Aufbau principle

Aufbau principle states that electrons fill orbitals in increasing order of energy. The lower energy orbital fills before the higher energy orbital.

Common filling order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s

Pauli exclusion principle

Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers.

An orbital can contain a maximum of two electrons. These two electrons must have opposite spins.

Hund’s rule of maximum multiplicity

Hund’s rule states that pairing of electrons in degenerate orbitals begins only after each orbital gets one electron.

In p, d and f subshells, electrons first occupy empty orbitals singly. These singly filled electrons have parallel spins.

Examples of electronic configuration

Hydrogen, Z = 1

Electronic configuration: 1s¹

Carbon, Z = 6

Electronic configuration: 1s² 2s² 2p²

Sodium, Z = 11

Electronic configuration: 1s² 2s² 2p⁶ 3s¹

Chlorine, Z = 17

Electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁵

Important Formulas in Structure of Atom

The main formulas in CBSE Class 11 Chemistry Chapter 2 Structure of Atom connect charge, radiation, spectra, Bohr energy and electron wavelength. These formulas support most numerical questions from the chapter.

More formulas:

1. e/me = 1.758820 × 10¹¹ C kg⁻¹
2. A = Z + n
3. Number of neutrons = A − Z
4. ΔE = E2 − E1 = hν
5. En = −RH × (1/n²)
6. rn = n²a0
7. Δx × Δp ≥ h/4π
8. ν̅ = 109677 × (1/n1² − 1/n2²) cm⁻¹

Important Terms in Structure of Atom

Structure of Atom uses terms from experiments, atomic models, spectra and quantum mechanics. These terms help students answer one-mark and reasoning questions in CBSE Class 11 Chemistry Chapter 2.

Electron

Electron is a negatively charged subatomic particle with charge −1.602176 × 10⁻¹⁹ C.

Proton

Proton is the lightest positively charged particle and is found in the nucleus.

Neutron

Neutron is an electrically neutral particle found in the nucleus.

Atomic number

Atomic number is the number of protons in the nucleus.

Mass number

Mass number is the total number of protons and neutrons.

Isotope

Isotopes are atoms with the same atomic number but different mass numbers.

Orbital

Orbital is a region around the nucleus where the probability of finding an electron is high.

Quantum number

Quantum number describes the state, shape, orientation or spin of an electron.

Electronic configuration

Electronic configuration is the arrangement of electrons in atomic orbitals.

NCERT-Style Questions from Structure of Atom

In CBSE Class 11 Chemistry Chapter 2 Structure of Atom, NCERT-style questions usually test subatomic particles, atomic models, spectra, Bohr calculations and quantum numbers. Strong answers use the correct formula, value and reason.

Q1. Calculate the number of protons, neutrons and electrons in ⁸⁰₃₅Br.

The atom has 35 protons, 35 electrons and 45 neutrons.

Step 1:

Atomic number, Z = 35

So, number of protons = 35

Step 2:

Neutral atom has electrons = protons

Electrons = 35

Step 3:

Number of neutrons = A − Z

= 80 − 35

= 45

Q2. Why did Rutherford’s model fail?

Rutherford’s model failed because it could not explain atomic stability.

Explanation:

An electron moving in a circular orbit is an accelerating charged particle. According to electromagnetic theory, it should lose energy and fall into the nucleus.

Fact:

Real atoms are stable, so Rutherford’s explanation was incomplete.

Q3. What is the difference between orbit and orbital?

An orbit is a fixed circular path, while an orbital is a probability region around the nucleus.

Explanation:

Bohr used orbits for electrons in hydrogen. Quantum mechanics uses orbitals because exact electron paths cannot be known.

Fact:

An orbital is described by the wave function ψ.

Q4. What is the de Broglie relation?

The de Broglie relation connects wavelength with momentum of a moving particle.

Formula:

λ = h/mv

Explanation:

Electrons show wave character because their wavelength can be detected experimentally. Ordinary objects have wavelengths too small to observe.

Q5. How many orbitals are present for n = 3?

For n = 3, the total number of orbitals is 9.

Formula:

Number of orbitals in a shell = n²

Substitution:

n² = 3² = 9

These include 3s, 3p and 3d orbitals.