Structure of Atom Notes for NEET and JEE Main 2027: Key Points, PDF

Structure of Atom is Unit 2 of NCERT Class 11 Chemistry Part I and the Atomic Structure unit of the NEET UG and JEE Main syllabus. These notes cover subatomic particles, Thomson and Rutherford models, Planck’s theory, the photoelectric effect, hydrogen spectrum, Bohr model, de Broglie and Heisenberg, quantum numbers, orbital shapes and electronic configuration, with a free PDF.

Structure of Atom Notes for NEET and JEE Main

Section-wise key points from NCERT Unit 2, formula and comparison tables, high-yield facts and common mistakes for 2027 revision.

Structure of Atom at a Glance

The topic lists follow the NEET UG 2026 syllabus released by the National Medical Commission (NMC) on 22 December 2025 and the JEE Main 2026 syllabus released by NTA on 31 October 2025, which stay the reference until the 2027 syllabi are out. Full lists: NEET syllabus and JEE Main syllabus. The Physics syllabus of both exams also covers the photoelectric effect, de Broglie relation and Bohr model under Dual Nature of Matter and Atoms and Nuclei.

Discovery of Electron, Proton and Neutron

Experiments on electrical discharge through gases showed that atoms are made of electrons, protons and neutrons.

• Cathode rays travel from cathode to anode in straight lines, behave like negatively charged particles (electrons) in electric and magnetic fields, and do not depend on the electrode material or the gas.
• J.J. Thomson (1897) measured the charge to mass ratio of the electron: e/me = 1.758820 × 10¹¹ C kg⁻¹.
• R.A. Millikan’s oil drop experiment (1906-14) gave the electron’s charge; the accepted value is −1.602176 × 10⁻¹⁹ C, which gives me = 9.1094 × 10⁻³¹ kg.
• Canal rays are positive gaseous ions whose mass and e/m depend on the gas. The smallest, from hydrogen, is the proton (characterised in 1919).
• Chadwick (1932) discovered the neutron by bombarding a thin sheet of beryllium with α-particles.

Thomson and Rutherford Models of the Atom

Thomson spread the positive charge and mass evenly over the atom; Rutherford’s α-particle scattering showed they are concentrated in a tiny nucleus.

• Rutherford fired α-particles at gold foil about 100 nm thick: most passed undeflected, a small fraction deflected by small angles, and about 1 in 20,000 bounced back by nearly 180°.
• Conclusions: most of the atom is empty space; the positive charge sits in a very small nucleus. Atomic radius is about 10⁻¹⁰ m and nuclear radius about 10⁻¹⁵ m.
• Drawback: an orbiting electron is accelerating, so by Maxwell’s theory it should radiate energy and spiral into the nucleus in about 10⁻⁸ s, which does not happen.
• Atomic number Z = number of protons = number of electrons in a neutral atom. Mass number A = protons + neutrons; neutrons = A − Z for atoms and ions alike.
• Isotopes: same Z, different A (protium 99.985%, deuterium 0.015%, tritium in traces; 35Cl and 37Cl). Isobars: same A, different Z (14C and 14N). Isotopes show the same chemical behaviour.
• α-particles are helium nuclei (charge +2, mass 4 u). Penetrating power: β is about 100 times and γ about 1000 times that of α.

Electromagnetic Radiation and Planck’s Quantum Theory

Light has a wave nature (diffraction, interference) and a particle nature (black-body radiation, photoelectric effect), so electromagnetic radiation shows dual behaviour.

• Electromagnetic waves need no medium; the electric and magnetic fields are perpendicular to each other and to the direction of travel.
• c = νλ, with c = 3.0 × 10⁸ m s⁻¹ in vacuum. Wavenumber ν̄ = 1/λ (SI unit m⁻¹; cm⁻¹ is common).
• Visible light runs from violet, 400 nm (7.50 × 10¹⁴ Hz), to red, 750 nm (4.00 × 10¹⁴ Hz).
• A black body absorbs and emits all frequencies; as temperature rises, the intensity maximum shifts to shorter wavelength.
• Planck (1900): energy is emitted or absorbed only in discrete quanta, E = hν, where h = 6.626 × 10⁻³⁴ J s. Allowed energies are 0, hν, 2hν … nhν.

Photoelectric Effect

When light above a threshold frequency strikes certain metals, electrons are ejected at once; Einstein (1905) explained this with photons.

• Observed by H. Hertz (1887) with metals such as potassium, rubidium and caesium.
• Electrons are ejected with no time lag; their number is proportional to the intensity of light.
• Below the threshold frequency ν0 no electrons are ejected, however bright the light. Kinetic energy rises with frequency, not intensity.
• Einstein’s equation: hν = hν0 + ½mev², where hν0 is the work function W0.
• Threshold frequency of potassium is 5.0 × 10¹⁴ Hz: red light does not eject electrons, but even weak yellow light does.

Line Spectrum of Hydrogen

Atoms in the gas phase give line (atomic) spectra with bright lines at specific wavelengths; each element’s line spectrum is unique.

• An emission spectrum comes from excited atoms giving up energy; an absorption spectrum is its photographic negative, with dark lines on a continuous background.
• Rubidium, caesium, thallium, indium, gallium and scandium were discovered by spectroscopy, and helium was discovered in the sun.
• Rydberg equation: ν̄ = 109,677 (1/n1² − 1/n2²) cm⁻¹, where 109,677 cm⁻¹ is the Rydberg constant for hydrogen and n2 > n1.

Bohr’s Model of the Hydrogen Atom

Niels Bohr (1913) applied Planck’s quantisation to the hydrogen atom and explained its line spectrum quantitatively.

• The electron moves in circular orbits (stationary states) of fixed radius and energy; energy does not change while it stays in an orbit.
• Energy is absorbed or emitted only when the electron jumps between orbits: ν = ΔE/h (Bohr’s frequency rule).
• Angular momentum is quantised: mevr = nh/2π, n = 1, 2, 3 …
• A free electron at rest (n = ∞) has zero energy; bound electrons have negative energy, most negative at n = 1, the ground state.

• Frequency of a line: ν = 3.29 × 10¹⁵ (1/ni² − 1/nf²) Hz; wavenumber: 1.09677 × 10⁷ (1/ni² − 1/nf²) m⁻¹.
• Absorption: nf > ni, energy absorbed. Emission: ni > nf, energy released.
• Electron velocity increases with nuclear charge and decreases with n.
• Limitations: cannot explain the doublets (fine lines) of the hydrogen spectrum, spectra of multi-electron atoms such as helium, splitting in a magnetic field (Zeeman effect) or electric field (Stark effect), or the formation of chemical bonds.

Dual Behaviour of Matter and the Uncertainty Principle

de Broglie and Heisenberg showed why the Bohr picture of fixed orbits had to be replaced.

• de Broglie (1924): every moving particle has a wavelength λ = h/mv = h/p. It was confirmed when an electron beam was diffracted, and it underlies the electron microscope (magnification about 15 million times).
• Wavelengths of everyday objects are too small to detect; those of electrons can be measured.
• Heisenberg (1927): the exact position and exact momentum of an electron cannot be known at the same time: Δx × Δp ≥ h/4π.
• The principle rules out definite paths or trajectories for electrons; it matters for microscopic particles and is negligible for macroscopic objects.
• Bohr’s model ignores the wave character of the electron and contradicts the uncertainty principle, which is why it failed.

Quantum Mechanical Model and Quantum Numbers

Schrödinger’s equation, Ĥψ = Eψ, gives the allowed energies and wave functions (orbitals) of an electron; |ψ|² gives the probability density.

• Quantum mechanics was developed independently in 1926 by Heisenberg and Schrödinger.
• ψ itself has no physical meaning; |ψ|² is the probability density and is always positive.
• An orbit (Bohr) is a definite path that cannot be demonstrated; an orbital is a one-electron wave function ψ defined by n, l and ml.
• An orbital holds at most two electrons, with opposite spins.

Shapes of Orbitals and Nodes

s orbitals are spherical, p orbitals have two lobes, and four of the five d orbitals share one shape while dz² is different.

• For 1s, probability density is highest at the nucleus and falls sharply; for 2s it falls to zero (a node), rises to a small maximum and then decreases.
• Boundary surface diagrams enclose about 90% probability; a 100% boundary is impossible because |ψ|² never becomes exactly zero at finite distance.
• Size and energy of s orbitals: 4s > 3s > 2s > 1s; of p orbitals: 4p > 3p > 2p.
• The three 2p orbitals (px, py, pz) have identical size, shape and energy and differ only in orientation; there is no simple link between ml values and x, y, z.
• d orbitals start at n = 3: dxy, dyz, dxz, dx²−y² and dz². All five 3d orbitals have equal energy.

Energies of Orbitals and Filling of Electrons

In hydrogen, orbital energy depends only on n; in multi-electron atoms it depends on n and l because of shielding by inner electrons.

• Hydrogen: 1s < 2s = 2p < 3s = 3p = 3d < 4s = 4p = 4d = 4f (orbitals with the same n are degenerate).
• Multi-electron atoms: within a shell, s < p < d < f. The (n + l) rule: lower (n + l) means lower energy; for equal (n + l), lower n is lower in energy.
• Effective nuclear charge (Zeff) felt by an electron falls with increasing l in a shell, so s electrons are held most tightly. Energy of the same orbital falls with Z: E2s(H) > E2s(Li) > E2s(Na) > E2s(K).
• Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s.
• Aufbau principle: orbitals fill in order of increasing energy.
• Pauli exclusion principle (1926): no two electrons in an atom have the same set of four quantum numbers, so an orbital holds two electrons of opposite spin and a shell holds 2n².
• Hund’s rule of maximum multiplicity: pairing in p, d or f starts only after each orbital is singly occupied, that is with the 4th, 6th and 8th electron.
• Exceptions: Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹, because half-filled and completely filled subshells are extra stable (symmetrical distribution and maximum exchange energy).
• Electrons in completely filled inner shells are core electrons; those in the outermost shell are valence electrons.

Most-asked Facts for NEET and JEE Main: Structure of Atom

Revise these one-line facts from the NCERT unit just before the exam.

• e/me = 1.758820 × 10¹¹ C kg⁻¹ (Thomson); e = −1.602176 × 10⁻¹⁹ C (Millikan); me = 9.1094 × 10⁻³¹ kg.
• Neutron: Chadwick, 1932, by bombarding beryllium with α-particles.
• About 1 in 20,000 α-particles bounced back from gold foil.
• 14C and 14N are isobars; 35Cl and 37Cl are isotopes.
• E = hν; h = 6.626 × 10⁻³⁴ J s; c = νλ.
• hν = hν0 + ½mev²; photoelectron KE depends on frequency, number on intensity.
• Balmer series is the only visible series of hydrogen; Lyman is ultraviolet.
• Bohr radius 52.9 pm; E1 of hydrogen = −2.18 × 10⁻¹⁸ J.
• rn ∝ n²/Z and En ∝ −Z²/n² for hydrogen-like species.
• λ = h/mv (de Broglie); Δx·Δp ≥ h/4π (Heisenberg).
• Number of orbitals in a shell = n²; maximum electrons = 2n²; orbitals in a subshell = 2l + 1.
• Radial nodes = n − l − 1; angular nodes = l; total nodes = n − 1.
• Cr: [Ar]3d⁵4s¹; Cu: [Ar]3d¹⁰4s¹.

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this unit.

• Using orbit and orbital as the same word. An orbit is Bohr’s fixed path; an orbital is a wave function.
• Thinking brighter light gives faster photoelectrons. Intensity changes only the number of electrons; frequency changes their kinetic energy.
• Applying Bohr’s formulas to helium atoms. They work only for one-electron species such as H, He⁺, Li²⁺ and Be³⁺.
• Forgetting Z² in the energy of hydrogen-like ions, or using Z² instead of Z in the radius.
• Writing Cr as 3d⁴4s² or Cu as 3d⁹4s².
• Mixing up radial and angular nodes. A 2p orbital has one angular node and no radial node.
• Allowing l = n. The maximum value of l is n − 1, so there is no 1p or 2d orbital.
• Saying hydrogen 2s is lower in energy than 2p. In hydrogen they are degenerate; the s < p order applies to multi-electron atoms.
• Assuming isotopes have different chemical properties. Chemistry depends on electrons, so isotopes behave the same.

Structure of Atom Notes PDF and Related Pages

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• Some Basic Concepts of Chemistry
• Structure of Atom
• Chemical Kinetics
• Equilibrium
• Organic Chemistry: Some Basic Principles and Techniques
• Electrochemistry
• Amines
• Solutions
• Thermodynamics
• Coordination Compounds
• Haloalkanes and Haloarenes
• p-Block Elements
• Alcohols, Phenols and Ethers
• Hydrocarbons
• Aldehydes, Ketones and Carboxylic Acids

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