Atomic Structure — NEET Class 11 revision
The full NEET Atomic Structure unit in one place: concept notes from the Bohr model to quantum numbers, orbitals and electronic configuration, differentiate tables, common doubts, quick-recall facts — plus an interactive 3D Bohr atom you can rotate and tap through.
Explore it visually
Tap through the interactive models for this unit — each part explained in plain English with an exam-accurate NEET note.
Key concepts
Electromagnetic radiation and the photoelectric effect reveal the quantised nature of energy. The hydrogen spectrum led to Bohr's model and discrete energy levels. de Broglie's relation and Heisenberg's principle then led to the quantum-mechanical model, where electrons occupy orbitals described by wave functions and quantum numbers. Aufbau, Pauli and Hund determine electronic configurations, with extra stability for half-filled and completely filled subshells.
Electromagnetic radiation
Electromagnetic radiation is described by wavelength and frequency, related by the speed of light. Each photon carries an energy proportional to its frequency.
- c = νλ
- E = hν = hc/λ
Photoelectric effect
Electrons are emitted from a metal only when the incident light frequency exceeds a threshold. Increasing intensity increases the number of emitted electrons, while frequency controls the maximum kinetic energy.
- Below threshold frequency → no photoemission
- K_max = hν − hν₀
Hydrogen emission spectrum
Excited hydrogen atoms emit discrete spectral lines because electrons can occupy only specific energy levels. A line appears when an electron moves from a higher to a lower level.
- 1/λ = R_H(1/n₁² − 1/n₂²), n₂ > n₁
- Lyman → n₁=1; Balmer → n₁=2
Bohr model of hydrogen
Bohr allowed only stationary circular orbits with quantised angular momentum. For a hydrogen-like species, orbit radius grows as n² and electron energy varies as −1/n².
- r_n = a₀ n²/Z
- E_n = −13.6 Z²/n² eV
- mvr = nh/2π
Limits of the Bohr model
Bohr's model works well for hydrogen-like one-electron species but cannot fully explain many-electron spectra, fine structure, Zeeman/Stark effects or the wave nature of electrons.
- Best for H, He⁺, Li²⁺ …
- Fixed classical paths are replaced by orbitals in quantum mechanics
Dual nature of matter
de Broglie proposed that moving particles have a wavelength. The wavelength becomes important for microscopic particles such as electrons.
- λ = h/p = h/mv
Heisenberg uncertainty principle
Position and momentum cannot both be known with unlimited precision. This is a fundamental quantum limit, not an experimental defect.
- Δx·Δp ≥ ħ/2
Quantum-mechanical orbitals
An orbital is described by a one-electron wave function ψ; ψ² gives the probability density. It is a probability distribution, not a fixed path.
- ψ → wave function
- ψ² → probability density
- 2s has one radial node; 1s has none
Quantum numbers
Four quantum numbers specify an electron's shell, subshell, orbital orientation and spin.
- n = 1,2,3,…
- l = 0 to n−1
- m_l = −l to +l
- m_s = +½ or −½
Shapes and capacities of orbitals
s orbitals are spherical, p orbitals are dumb-bell shaped, and d orbitals have characteristic multi-lobed shapes. A subshell contains 2l+1 orbitals.
- s: 1 orbital, 2 e⁻
- p: 3 orbitals, 6 e⁻
- d: 5 orbitals, 10 e⁻
- f: 7 orbitals, 14 e⁻
Aufbau, Pauli and Hund
Electrons fill lower-energy orbitals first, no two electrons in an atom have all four quantum numbers identical, and degenerate orbitals are singly occupied with parallel spins before pairing.
- Aufbau: increasing n+l; lower n first if tied
- Pauli: max 2 e⁻ per orbital with opposite spins
- Hund: singly fill degenerate orbitals first
Electronic configuration and extra stability
Electronic configurations follow the filling rules, but half-filled and completely filled subshells have extra stability. This explains familiar exceptions such as chromium and copper.
- Cr: [Ar] 3d⁵4s¹
- Cu: [Ar] 3d¹⁰4s¹
- Half-filled and filled subshells are especially stable
Differentiate the three electron-filling rules
Every configuration obeys all three together.
The filling rules
| Aufbau | Fill the lowest-energy orbital first (n + l order) |
| Pauli | ≤ 2 electrons per orbital, with opposite spins |
| Hund | Singly fill degenerate orbitals before pairing (parallel spins) |
Concepts students confuse
Tap the pairs that decide tricky NEET questions.
Orbit vs orbital ★
Orbit (Bohr)
- Fixed circular path
- Definite radius
- Outdated for many-electron atoms
Orbital (quantum)
- Probability distribution
- No definite path
- 3-D region of high probability
Aufbau vs Hund's rule ★
Aufbau
- Which orbital fills first
- By increasing energy (n+l)
- Sets the order
Hund's rule
- How to fill a subshell
- Singly, parallel spins, then pair
- Within equal-energy orbitals
de Broglie vs Heisenberg
de Broglie
- Particles have wavelength
- λ = h/mv
- Wave nature of matter
Heisenberg
- Position & momentum limit
- Δx·Δp ≥ ħ/2
- Fundamental, not instrumental
Subshells: s vs p vs d
s (l=0)
- 1 orbital
- Holds 2 electrons
- Spherical
p (l=1) / d (l=2)
- 3 / 5 orbitals
- Hold 6 / 10 electrons
- 2l+1 orbitals each
Common NEET doubts
Do I use the mass number to count electrons?
No. Use the atomic number Z (protons). In a neutral atom the number of electrons equals Z, not the mass number.
Should the M shell get 18 electrons before 4s fills?
No. 4s fills before 3d, so potassium and calcium are 2,8,8,1 and 2,8,8,2 — the M shell is not filled to 18 first.
Is helium in group 2 because it has two electrons?
No. Its K shell is complete, so helium behaves as a noble gas and is placed in group 18.
Is an orbital a circular path the electron follows?
No. An orbital is a three-dimensional probability distribution — the region where the electron is most likely to be found, not a fixed path.
Can the azimuthal quantum number l equal n?
No. l ranges from 0 to n − 1, and mₗ ranges from −l to +l. l can never equal or exceed n.
Can degenerate orbitals be paired before each holds one electron?
No. Hund's rule says fill each degenerate orbital singly (parallel spins) before any pairing.
Is the uncertainty principle just a limit of our instruments?
No. It is a fundamental property of nature, not a measurement error — you cannot know exact position and momentum simultaneously even in principle.
Does higher light intensity overcome a frequency below the photoelectric threshold?
No. Below the threshold frequency, photoemission does not occur however intense the light is.
Quick recall
Q. What does the atomic number Z equal?
The number of protons — and, in a neutral atom, the number of electrons.
Q. State the filling order for the first twenty elements.
1s, 2s, 2p, 3s, 3p, 4s (4s before 3d).
Q. State the Aufbau principle.
Electrons fill the lowest-energy orbital first, ordered by increasing n + l.
Q. State the Pauli exclusion principle.
An orbital holds at most two electrons, and they must have opposite spins.
Q. State Hund's rule.
Degenerate orbitals are filled singly with parallel spins before any pairing.
Q. How many electrons can a subshell of azimuthal number l hold?
2(2l + 1): s = 2, p = 6, d = 10, f = 14.
Q. Write the de Broglie wavelength.
λ = h / p = h / mv.
Q. What are the allowed values of l for shell n?
0 to n − 1 (s, p, d, f for l = 0, 1, 2, 3).
Q. Write the photon energy and wave relation for electromagnetic radiation.
E = hν = hc/λ and c = νλ.
Q. Write the Bohr energy of a hydrogen-like species.
E_n = −13.6 Z²/n² eV.
Q. What does ψ² represent for an atomic orbital?
The probability density of finding the electron at a point in space.
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