NEET Chemistry - Chapter 2

Structure of Atom

Fresh NEET chemistry notes on atomic models, electromagnetic radiation, Bohr theory, hydrogen spectrum, de Broglie relation, uncertainty principle, quantum numbers, and electronic configuration.

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1. Atomic Models: Thomson → Rutherford → Bohr

Each model corrected its predecessor's fatal flaw. Understanding what each model explained and what it could not is pure NEET gold.

Scientist	Key Contribution	Fatal Flaw
Thomson (1897)	Discovered electron; plum-pudding model	Could not explain α-scattering results
Rutherford (1911)	Nucleus with positive charge; electrons orbit outside	Orbiting electrons should lose energy and spiral in (classical EM)
Bohr (1913)	Quantized stationary orbits for H-like species; no energy loss	Failed for multi-electron atoms; ignored wave nature

Bohr's postulates for hydrogen-like species ( $H, He^+, Li^{2+}$ ...):

mvr=\frac{nh}{2\pi}\qquad\text{(angular momentum quantisation)}

r_n=\frac{0.529\, n^2}{Z}\text{ Å},\qquad E_n=\frac{-13.6\,Z^2}{n^2}\text{ eV}

Numeric check: For H atom, ground state energy =

-13.6

eV (

n=1

). For

He^+

(

Z=2

), ground state energy =

-13.6\times4 = -54.4

eV. Ionisation energy =

13.6\,Z^2/n^2

eV.

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2. Electromagnetic Radiation, Photons, and Hydrogen Line Spectra

Light has a dual nature — wave and particle. The wave side gives wavelength $\lambda$ and frequency $\nu$ ; the particle side gives photon energy.

c=\lambda\nu\qquad(c=3\times10^8\text{ m s}^{-1})

E=h\nu=\frac{hc}{\lambda}\qquad(h=6.626\times10^{-34}\text{ J s})

\Delta E=E_{n_2}-E_{n_1}=13.6\,Z^2\left(\frac{1}{n_1^2}-\frac{1}{n_2^2}\right)\text{ eV}

Hydrogen line-spectrum series — memorise the series by $n_1$ and region:

Series	$n_1$	$n_2$	Region
Lyman	1	2,3,4…	Ultraviolet
Balmer	2	3,4,5…	Visible
Paschen	3	4,5,6…	Infrared
Brackett	4	5,6,7…	Infrared
Pfund	5	6,7…	Far Infrared

NEET tip: The number of spectral lines when an electron falls from level

n

to ground state =

\frac{n(n-1)}{2}

. From

n=4

: lines =

4\times3/2 = 6

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3. de Broglie Wave Nature, Heisenberg Uncertainty Principle

Every moving particle has an associated wavelength — this is the de Broglie hypothesis. It makes electrons (and all microscopic particles) wave-particle duals.

\lambda=\frac{h}{mv}=\frac{h}{p}

For an electron accelerated through potential $V$ : $\lambda=\frac{h}{\sqrt{2meV}}$ .

Heisenberg Uncertainty Principle: It is physically impossible to determine simultaneously the exact position and exact momentum of a microscopic particle.

\Delta x \cdot \Delta p \geq \frac{h}{4\pi}

Also: $\Delta E \cdot \Delta t \geq h/4\pi$ . Note: $\Delta x \cdot \Delta v \geq h/(4\pi m)$ for velocity uncertainty.

NEET trap: The uncertainty principle is a fundamental property of quantum objects — NOT a limitation of our measurement instruments. The act of measurement itself disturbs microscopic particles.

These ideas led to quantum mechanics where electrons occupy orbitals (3D probability regions) rather than Bohr's definite circular orbits.

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4. Quantum Numbers, Orbital Shapes, and Nodes

Four quantum numbers uniquely identify every electron in an atom. They arise naturally from Schrödinger's wave equation.

Quantum Number	Symbol	Allowed Values	Determines
Principal	$n$	1, 2, 3…	Shell size and energy
Azimuthal	$l$	0 to $n-1$	Subshell shape (s,p,d,f)
Magnetic	$m_l$	$-l$ to $+l$	Orbital orientation
Spin	$m_s$	$+\frac{1}{2}$ or $-\frac{1}{2}$	Electron spin direction

\text{Orbitals in shell }n = n^2,\quad\text{Max electrons in shell} = 2n^2

\text{Radial nodes} = n-l-1,\quad\text{Angular nodes} = l,\quad\text{Total nodes} = n-1

Subshell naming: $l=0$ (s), $l=1$ (p), $l=2$ (d), $l=3$ (f). A 3d orbital has $n=3$ , $l=2$ , so radial nodes = $3-2-1=0$ , angular nodes = 2, total = 2.

NEET quick counts: 2p: 3 orbitals; 3d: 5 orbitals; 4f: 7 orbitals. p orbitals are dumbbell-shaped; d orbitals mostly double-dumbbell or cloverleaf.

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5. Electronic Configuration: Rules, Exceptions, and NEET Traps

Three rules govern how electrons fill orbitals: Aufbau (fill lowest energy first), Pauli exclusion (max 2 electrons per orbital, opposite spins), and Hund's rule (maximum unpaired electrons in degenerate orbitals).

The Aufbau filling order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p…

Critical Exceptions (Half-filled and completely-filled stability)

Cr (Z=24): expected [Ar]3d $^4$ 4s $^2$ , actual [Ar]3d $^5$ 4s $^1$ (half-filled 3d is extra stable)
Cu (Z=29): expected [Ar]3d $^9$ 4s $^2$ , actual [Ar]3d $^{10}$ 4s $^1$ (completely filled 3d is extra stable)
Mo (Z=42), Pd (Z=46), Ag (Z=47) similarly deviate.

Cation configuration trap: When transition metals form cations, they lose ns electrons first, not (n-1)d electrons. Fe

^{2+}

is [Ar]3d

^6

(lost 4s

^2

), not [Ar]3d

^4

^2

Magnetic behaviour: Paramagnetic = has unpaired electrons. Diamagnetic = all paired. Count unpaired electrons from config: $Cu^+$ is [Ar]3d $^{10}$ — diamagnetic. $Fe^{3+}$ is [Ar]3d $^5$ — 5 unpaired, strongly paramagnetic.

NEET trap: Orbit (Bohr, circular path) ≠ Orbital (quantum mechanics, 3D probability region). Never say an electron "orbits" in quantum chemistry — it occupies an orbital.

Practice Tests

5 Chapter Tests of 25 Questions Each

Each test is original, NEET-aligned, and answer-backed. Use them as sectional revision instead of a single long mock so your weak subtopics become easier to identify quickly.

Test 1: Atomic Models

Subatomic particles, Thomson, Rutherford, photons, and Bohr basics.

Test 2: Bohr and Spectrum

Radius, energy, spectral series, hydrogen-like ions, and transitions.

Test 3: Quantum Mechanics

de Broglie relation, uncertainty principle, orbitals, quantum numbers, and nodes.

Test 4: Configuration Rules

Aufbau principle, Pauli, Hund, exceptions, magnetism, and cation configurations.

Test 5: Mixed NEET Drill

Integrated questions across models, spectra, quantum numbers, and configurations.

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