Atoms Notes for NEET and JEE Main 2027 PDF and Key Points

Atoms is Chapter 12 of NCERT Class 12 Physics Part II and part of Unit 18, Atoms and Nuclei, in both the NEET UG and JEE Main syllabi. These notes cover Thomson’s model, the Geiger-Marsden alpha-particle scattering experiment, Rutherford’s nuclear model, Bohr’s postulates, hydrogen energy levels and line spectra, and de Broglie’s explanation, with a free notes PDF.

Atoms Notes for NEET and JEE Main

Section-wise key points from NCERT Chapter 12, formula tables, high-yield facts and common mistakes for 2027 revision.

Atoms at a Glance

Item	Details
Syllabus unit	NEET UG 2026 and JEE Main 2026: Unit 18, Atoms and Nuclei
Exams that cover it	Both NEET UG and JEE Main (Paper 1); the topic is listed in the 2026 syllabus of each
NCERT textbook	Physics Part II, Class 12, Chapter 12 (rationalised book)
Other chapter in the unit	Nuclei (Chapter 13)
Syllabus topics	Alpha-particle scattering experiment; Rutherford’s model of atom; Bohr model, energy levels, hydrogen spectrum (the nucleus topics of the unit are in Nuclei, Chapter 13)
Official chapter	NCERT Chapter 12
Notes PDF	View PDF (AglaSem Class 12 notes for Dual Nature of Radiation and Matter and Atoms and Nuclei)

The topic list follows 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.

Early Models: Thomson’s Atom and Line Spectra

J. J. Thomson’s 1898 model pictured the atom as a uniform sphere of positive charge with electrons embedded in it like seeds in a watermelon, the plum pudding model, which later experiments disproved.

Atoms are electrically neutral, so the negative charge of electrons must be balanced by positive charge.
Solids, liquids and dense gases emit a continuous spectrum; rarefied gases (for example in a neon sign or mercury vapour lamp) emit only certain discrete wavelengths, a line spectrum.
Each element has a characteristic spectrum; hydrogen always gives lines with fixed relative positions.
Balmer (1885) found an empirical formula for the wavelengths of a group of hydrogen lines.

Alpha-particle Scattering and Rutherford’s Nuclear Model

The Geiger-Marsden experiment (1911) showed that the positive charge and most of the mass of an atom are concentrated in a tiny nucleus, which led to Rutherford’s nuclear (planetary) model.

Set-up: a beam of 5.5 MeV alpha-particles from a radioactive source, collimated by lead bricks, hit a gold foil 2.1 × 10⁻⁷ m thick. Scattered particles produced flashes on a zinc sulphide screen viewed through a microscope; the apparatus was in a vacuum chamber.
Most alpha-particles passed straight through. Only about 0.14% were scattered by more than 1°, and about 1 in 8000 by more than 90°.
Large-angle scattering needs a large repulsive force, possible only if the positive charge and most of the mass are concentrated at the centre.
Nucleus size: about 10⁻¹⁵ m to 10⁻¹⁴ m; atom size about 10⁻¹⁰ m, so the atom is 10,000 to 100,000 times larger and mostly empty space.
Gold has Z = 79 and its nucleus is about 50 times heavier than an alpha-particle, so it can be treated as stationary; the light electrons hardly affect the alpha-particles.
Impact parameter b: the perpendicular distance of the initial velocity of the alpha-particle from the centre of the nucleus. Small b gives large scattering; a head-on collision (minimum b) makes it bounce back (θ ≈ π); large b gives almost no deflection.
Distance of closest approach d = 2Ze²/(4πε₀K) from energy conservation. NCERT Example 12.2: for a 7.7 MeV alpha-particle and gold, d ≈ 3.0 × 10⁻¹⁴ m (30 fm), while the gold nucleus radius is about 6 fm.

Electron Orbits in the Rutherford Model

Quantity (hydrogen, orbit radius r)	Expression
Centripetal force = electrostatic attraction	mv²/r = e²/(4πε₀r²)
Kinetic energy	K = e²/(8πε₀r)
Potential energy	U = −e²/(4πε₀r)
Total energy	E = K + U = −e²/(8πε₀r) (negative, so the electron is bound)
Ground state of hydrogen (E = −13.6 eV)	r ≈ 5.3 × 10⁻¹¹ m, v ≈ 2.2 × 10⁶ m/s (NCERT Example 12.3)

Atomic Spectra: Emission and Absorption

A gas excited at low pressure emits an emission line spectrum of bright lines at specific wavelengths, and white light passed through the same gas shows dark absorption lines at exactly those wavelengths.

Emission line spectra act as a fingerprint for identifying a gas.
Absorption lines appear in the continuous spectrum at the same wavelengths as the emission lines of the gas.

Bohr Model of the Hydrogen Atom

Bohr (1913) kept Rutherford’s nucleus but added three quantum postulates: stationary orbits that do not radiate, quantised angular momentum L = nh/2π, and photon emission hν = E_i − E_f when an electron jumps to a lower orbit.

Problems with Rutherford’s model: an orbiting electron is accelerated, so by classical theory it should radiate, lose energy and spiral into the nucleus (atoms would be unstable), and it would emit a continuous spectrum instead of lines.
Classical theory would give an initial emitted frequency of about 6.6 × 10¹⁵ Hz for hydrogen (NCERT Example 12.4).
First postulate: electrons revolve in certain stable orbits (stationary states) without radiating; each state has a definite total energy.
Second postulate: only orbits with angular momentum L = nh/2π are allowed (n = 1, 2, 3, …).
Third postulate: when an electron jumps from a higher energy orbit to a lower one it emits a photon with hν = E_i − E_f.

Quantity	Bohr model result for hydrogen
Radius of nth orbit	r_n = (n²/m)(h/2π)²(4πε₀/e²), so r_n ∝ n²
Bohr radius (n = 1)	a₀ ≈ 5.3 × 10⁻¹¹ m
Energy of nth level	E_n = −me⁴/(8n²ε₀²h²) = −2.18 × 10⁻¹⁸/n² J = −13.6/n² eV
Ground state (n = 1)	−13.6 eV; ionisation energy of hydrogen = 13.6 eV
n = 2 (first excited state)	−3.40 eV; excitation energy from ground state 10.2 eV
n = 3	−1.51 eV; excitation energy from ground state 12.09 eV
n = ∞	0 eV (electron removed and at rest); above 0 eV the electron is free and energies form a continuum

The negative energy shows the electron is bound; energy rises (becomes less negative) in outer orbits.
Energy levels get closer together as n increases.
Since K = e²/(8πε₀r) and U = −e²/(4πε₀r), K = −E and U = 2E in any orbit.
At room temperature most hydrogen atoms are in the ground state; collisions or photon absorption can excite them.
Bohr received the Nobel Prize in 1922.

Line Spectra of the Hydrogen Atom

Each spectral line comes from a jump between two Bohr levels: hν_if = E_ni − E_nf, so hydrogen emits only discrete frequencies.

Emission lines: electron falls from a higher to a lower level and emits a photon. Absorption: the atom absorbs a photon of exactly the energy needed to go up a level, giving dark lines.
In Bohr’s model the frequency of the spectral line is not the frequency of revolution; it is the energy difference divided by h. The two agree only for transitions between very large n.

Spectral Series (syllabus Topic Beyond the Rationalised NCERT Text)

The syllabi list the hydrogen spectrum; the named series below come from the older NCERT edition, since the rationalised chapter no longer lists them. The wavelengths follow 1/λ = R(1/n_f² − 1/n_i²), with the Rydberg constant R ≈ 1.097 × 10⁷ m⁻¹.

Series	Final level n_f	Region
Lyman	1	Ultraviolet
Balmer	2	Visible
Paschen	3	Infrared
Brackett	4	Infrared
Pfund	5	Infrared

De Broglie’s Explanation and Limits of the Bohr Model

De Broglie explained Bohr’s second postulate: an electron orbit is allowed only if its circumference holds a whole number of electron wavelengths, 2πr_n = nλ, which gives mv_nr_n = nh/2π.

With λ = h/mv_n, the standing-wave condition 2πr_n = nλ leads directly to L = nh/2π (NCERT Fig. 12.8 shows n = 4).
Davisson and Germer verified the wave nature of electrons in 1927.
Limitation 1: the Bohr model works only for hydrogenic (one-electron) atoms such as H, He⁺ and Li²⁺; it cannot be extended even to helium, because it ignores electron-electron forces.
Limitation 2: it predicts the frequencies of hydrogen lines but cannot explain their relative intensities.
Both Thomson’s and Rutherford’s models are unstable: Thomson’s electrostatically, Rutherford’s because of electromagnetic radiation from the orbiting electron.
The orbit picture was replaced by quantum mechanics, where orbits become regions of high probability and a state needs four quantum numbers (n, l, m, s).

Most-asked Facts for NEET and JEE Main: Atoms

Revise these one-line facts and formulas from the NCERT chapter just before the exam.

Thomson’s plum pudding model (1898); Geiger-Marsden experiment (1911) at Rutherford’s suggestion.
Alpha scattering used 5.5 MeV alpha-particles on a 2.1 × 10⁻⁷ m gold foil; about 1 in 8000 deflected beyond 90°.
Nucleus about 10⁻¹⁵ to 10⁻¹⁴ m; atom about 10⁻¹⁰ m.
Small impact parameter means large scattering; head-on collision means rebound.
Distance of closest approach d = 2Ze²/(4πε₀K).
Bohr’s quantisation: L = nh/2π; emission hν = E_i − E_f.
E_n = −13.6/n² eV; ionisation energy of hydrogen 13.6 eV.
E₂ = −3.40 eV, E₃ = −1.51 eV; excitation energies 10.2 eV and 12.09 eV.
r_n ∝ n²; Bohr radius about 5.3 × 10⁻¹¹ m.
K = −E and U = 2E for an electron in a Bohr orbit.
de Broglie condition: 2πr_n = nλ gives L = nh/2π.
Bohr model works only for one-electron (hydrogenic) atoms and cannot explain line intensities.
Lyman series is ultraviolet, Balmer visible, Paschen infrared.

Common Mistakes to Avoid

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

Saying most alpha-particles bounced back. Most went straight through; only about 1 in 8000 deflected by more than 90°.
Treating the distance of closest approach as the nuclear radius. It is only an upper limit (30 fm for gold vs an actual radius of about 6 fm).
Writing r_n ∝ n or E_n ∝ 1/n. The correct relations are r_n ∝ n² and E_n ∝ 1/n².
Confusing ionisation energy (13.6 eV from the ground state) with the first excitation energy (10.2 eV).
Taking kinetic energy as negative. K is positive and equals −E; potential energy U = 2E is negative.
Applying the Bohr model to neutral helium. It works only for one-electron systems such as He⁺ and Li²⁺.
Thinking emitted frequency equals the orbital frequency. In Bohr’s model it is (E_i − E_f)/h.
Assuming energy levels are equally spaced. They crowd together as n increases.

Download the AglaSem notes PDF that covers Atoms (AglaSem Class 12 notes for Dual Nature of Radiation and Matter and Atoms and Nuclei: key notes and important questions with answers) free: View PDF

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Practice: NEET question papers and JEE Main question papers
Source text: NCERT Chapter 12 (NCERT, Class 12 Physics Part II)

Physics Notes for NEET and JEE Main: All Chapters

All chapters with key points and notes PDFs are on the Physics notes for NEET and JEE Main page.

Atoms Notes for NEET and JEE Main FAQs

Which NCERT chapter is Atoms?

It is Chapter 12 of the NCERT Class 12 Physics Part II textbook (rationalised edition). Nuclei (Chapter 13) completes Unit 18, Atoms and Nuclei.

Is Atoms in the NEET 2027 and JEE Main 2027 syllabus?

It is part of Unit 18, Atoms and Nuclei, in both the NEET UG 2026 syllabus (NMC, 22 December 2025) and the JEE Main 2026 syllabus (NTA, 31 October 2025). The 2027 syllabi have not been released yet.

What are Bohr’s postulates?

Electrons move in certain stable orbits without radiating; only orbits with angular momentum L = nh/2π are allowed; and a jump to a lower orbit emits a photon with hν = E_i − E_f.

What is the energy of the electron in the nth orbit of hydrogen?

E_n = −13.6/n² eV. So the ground state is −13.6 eV, n = 2 is −3.40 eV and n = 3 is −1.51 eV; the ionisation energy of hydrogen is 13.6 eV.

Are the hydrogen spectral series in the rationalised NCERT book?

No. The rationalised chapter explains line spectra through Bohr levels but no longer names the Lyman, Balmer and other series. The syllabi still list the hydrogen spectrum, so revise them from the older NCERT edition.

Atoms Notes for NEET and JEE Main 2027: Key Points and PDF