Atomic Structure Full Chapter for NEET 2026: Concepts, Formulas and Revision by PW

Atomic Structure forms the foundation for understanding how atoms are organised and how electrons behave at different energy levels. The chapter connects the discovery of subatomic particles with atomic models, electromagnetic radiation, atomic spectra, quantum numbers and orbitals.

For NEET 2026 preparation, it is important to understand the concepts behind the formulas instead of relying only on memorisation. PW’s revision approach focuses on connecting important concepts with formulas and numerical applications so that you can revise the chapter systematically.

Matter and Dalton’s Atomic Theory

Matter is anything that has mass and occupies space. Air is also matter because it has mass and occupies volume and contains molecules such as O₂ and N₂.

John Dalton proposed the Dalton Atomic Theory to explain the nature of matter. According to the theory:

• Matter consists of very small particles.

• These particles are called atoms.

• Atoms were considered the smallest particles of matter.

• Atoms were considered indivisible.

The idea that atoms are indivisible was later modified after the discovery of subatomic particles.

Discovery of Subatomic Particles

Experiments involving vacuum tubes, cathode rays and discharge tubes established that atoms contain smaller particles.

Cathode rays consist of electrons, whereas anode rays, also called canal or positive rays, consist of positively charged gaseous ions.

James Chadwick discovered the neutron by bombarding beryllium with alpha particles. An alpha particle is the nucleus of a helium atom.

Quick Revision: Thomson → Electron; Goldstein → Positive rays; Chadwick → Neutron.

Thomson’s Plum-Pudding Model

J. J. Thomson proposed the plum-pudding model, also known as the watermelon model.

According to this model:

• Positive charge is uniformly distributed throughout the atom.

• Electrons are embedded in the positively charged sphere.

• The atom as a whole is electrically neutral.

The model could not explain the results of later experiments and was eventually rejected.

Rutherford’s Nuclear Model

Rutherford proposed his nuclear model based on the gold-foil experiment. In this experiment, alpha particles were directed towards a thin gold foil and their paths were observed.

Rutherford’s Observations and Conclusions

Rutherford proposed that nearly all the positive charge and mass of an atom are concentrated in a small nucleus, while electrons revolve around it.

However, the model could not explain atomic stability. According to classical electromagnetic theory, an accelerating charged particle should continuously lose energy. An electron revolving around the nucleus would therefore lose energy and eventually fall into the nucleus.

Charge and Mass of the Electron

The charge of an electron was determined through Millikan’s oil-drop experiment, while Thomson determined the charge-to-mass ratio of the electron.

Important values are:

• Electron charge: −1.6 × 10⁻¹⁹ C

• Electron charge-to-mass ratio: 1.67 × 10¹¹ C kg⁻¹

• Electron mass: 9.1 × 10⁻³¹ kg

Quick Revision: Millikan → Charge; Thomson → Charge-to-mass ratio.

Electromagnetic Radiation

Electromagnetic radiation can be described in terms of wavelength, frequency and photon energy.

The basic relation is:

c = νλ

where:

• c = speed of light

• ν = frequency

• λ = wavelength

Therefore:

λ = c/ν

The energy of a photon is:

E = hν = hc/λ

where h is Planck’s constant.

When solving numerical questions, make sure that the units of frequency and wavelength are converted into the required SI units.

For example, if the frequency is 1368 kHz:

1368 kHz = 1368 × 10³ Hz

Using λ = c/ν, the wavelength is approximately:

λ ≈ 219.2 m

Photoelectric Effect and Work Function

The photoelectric effect refers to the emission of electrons from a metal surface when suitable electromagnetic radiation falls on it.

The minimum energy required to remove an electron from the surface of a metal is called the work function.

The photoelectric equation can be written as:

φ = hc/λ − K.E.

or

Photon Energy = Work Function + Kinetic Energy

Here:

• φ = work function

• hc/λ = photon energy

• K.E. = kinetic energy of the emitted electron

If kinetic energy is given per mole, first convert it into the energy of one electron by dividing it by Avogadro’s number. After calculating the work function per electron, it can be converted back to the required molar quantity.

Hydrogen Atomic Spectrum

When an atom absorbs energy, an electron can move to a higher energy level. When the electron returns to a lower energy level, energy is released in the form of electromagnetic radiation.

The major hydrogen spectral series are:

Quick Revision: Lyman → n = 1 → Ultraviolet; Balmer → n = 2 → Visible; Paschen → n = 3 → Infrared.

The energy of radiation can be calculated using:

E = hc/λ

For a wavelength of 45 nm, first convert the wavelength into metres before substituting it into the formula.

Bohr Model of the Atom

Bohr proposed that electrons revolve around the nucleus only in certain permitted orbits with definite energies.

The radius of the nth orbit is:

rₙ = 0.53 n²/Z Å

where:

• n = principal quantum number

• Z = atomic number

For He⁺ with n = 2 and Z = 2:

r₂ = 1.058 Å = 105.8 pm

For Li²⁺ with n = 3 and Z = 3:

r₃ = 1.587 Å = 158.7 pm

Therefore, the third orbit of Li²⁺ has a radius 1.5 times that of the second orbit of He⁺.

Important Length Conversion

1 Å = 10⁻¹⁰ m = 100 pm

The energy of an electron in the nth Bohr orbit is:

Eₙ = −2.178 × 10⁻¹⁸ Z²/n² J

The negative sign indicates that the electron is bound to the nucleus. A more negative energy value represents stronger binding.

Quantum Numbers

Quantum numbers describe the state of an electron in an atom. For a given principal quantum number n, the possible values of the other quantum numbers are:

• Principal quantum number: n = 1, 2, 3, ...

• Azimuthal quantum number: l = 0 to n − 1

• Magnetic quantum number: mₗ = −l to +l

• Spin quantum number: mₛ = +1/2 or −1/2

For example, when n = 4 and l = 3:

mₗ = −3, −2, −1, 0, +1, +2, +3

The value 0 must be included.

Subshells and Orbitals

The value of the azimuthal quantum number determines the subshell.

An orbital is described by the quantum numbers n, l and mₗ. The complete description of an electron also requires the spin quantum number mₛ.

For n = 3 and l = 1, the subshell is 3p.

A p subshell contains:

• 3 orbitals

• Maximum 6 electrons

The probability density of finding an electron is proportional to the square of the wave function:

Probability Density ∝ ψ²

Shapes of Orbitals

The d orbitals include:

• dxy

• dyz

• dzx

• dx²−y²

• dz²

The first four have similar double-dumbbell shapes, while dz² has two lobes along the z-axis with a doughnut-shaped region around the centre.

The electron density of dx²−y² is concentrated along the x- and y-axes.

The electron density of dz² is concentrated along the z-axis.

Nodes in Orbitals

A node is a region where the probability of finding an electron is zero.

The number of nodes can be determined using:

Angular nodes = l

Radial nodes = n − l − 1

For a 3s orbital:

• n = 3

• l = 0

• Angular nodes = 0

• Radial nodes = 3 − 0 − 1 = 2

Therefore, the 3s orbital has 2 radial nodes and 0 angular nodes.

Orbital Energy and the n + l Rule

In multi-electron atoms, orbital energy can be compared using 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.

• Therefore, when n + l is equal, the orbital with higher n has higher energy.

For example:

5f > 6p > 5p > 4d

For titanium, the increasing order of orbital energy is:

3s < 3p < 4s < 3d

The n + l rule should not be applied to hydrogen-like one-electron species because their energy depends only on the principal quantum number.

Hund’s Rule

Hund’s rule states that electrons occupy degenerate orbitals singly with parallel spins before pairing takes place.

For example, nitrogen has the electronic configuration:

1s² 2s² 2p³

The three electrons in the 2p subshell occupy the three degenerate p orbitals singly before any pairing occurs.

Hydrogen-Like Species

Hydrogen-like species contain only one electron. Examples include H, He⁺ and Li²⁺.

For hydrogen-like species, the energy of an electron depends only on the principal quantum number n.

Therefore:

E₂s = E₂p

This is different from multi-electron atoms, where orbital energies are affected by factors such as shielding and penetration and can be compared using the n + l rule.

Key Takeaways for Atomic Structure

• Matter consists of particles, and the discovery of subatomic particles modified Dalton’s original atomic theory.

• Thomson identified the electron, Goldstein worked with positive rays, and Chadwick discovered the neutron.

• Rutherford’s model established the presence of a small, dense, positively charged nucleus.

• Photon energy is given by E = hν = hc/λ.

• The photoelectric equation connects photon energy, work function and kinetic energy.

• Lyman, Balmer and Paschen series end at n = 1, 2 and 3, respectively.

• Bohr’s model gives expressions for the radius and energy of hydrogen-like species.

• Quantum numbers describe the state and location of electrons in atoms.

• Radial nodes = n − l − 1, while angular nodes = l.

• Hund’s rule explains the arrangement of electrons in degenerate orbitals.

• The n + l rule is used for comparing orbital energies in multi-electron atoms.

Atomic Structure is a foundational Chemistry topic that helps you understand the composition, behaviour and energy of atoms. Revise subatomic particles, atomic models, electromagnetic radiation, hydrogen spectrum, Bohr’s model, quantum numbers, orbitals, nodes and electronic configuration along with regular NCERT reading, formula revision and question practice. A consistent revision routine can help you strengthen your understanding of Atomic Structure and build a strong foundation for upcoming Chemistry chapters and NEET preparation. 

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