Class 9 Structure of the Atom — Models, Particles and Electron Distribution
Having learnt in the previous NCERT chapter that matter is made of atoms and molecules, the "Structure of the Atom" chapter goes one level deeper: what is inside an atom? This chapter's three atomic models, its subatomic particles, and the rules for arranging electrons in shells are the foundation for everything you will study about bonding and the periodic table from Class 10 onward. Here is the chapter worked through with the calculations examiners actually ask.
The three subatomic particles
| Particle | Discovered by | Charge | Relative mass | Location |
|---|---|---|---|---|
| Electron | J. J. Thomson | −1 | ~1/2000 of a proton | Outside the nucleus, in shells |
| Proton | E. Goldstein (as canal rays); identified as the nucleus's positive particle by Rutherford | +1 | 1 (reference) | Inside the nucleus |
| Neutron | James Chadwick | 0 | ~1 | Inside the nucleus |
Three models, in order
Thomson's model (1898): the atom is a positively charged sphere with electrons embedded in it, like seeds in a watermelon — often called the "plum pudding" model. It correctly proposed that atoms are electrically neutral overall, but had no concept of a dense nucleus.
Rutherford's model (1911): based on his gold-foil (alpha-particle scattering) experiment, in which most alpha particles passed straight through a thin gold foil but a very small fraction bounced back sharply. Rutherford concluded that an atom has a tiny, dense, positively charged nucleus at its centre, with electrons revolving around it, and that most of the atom is empty space. This model could not explain why orbiting electrons do not lose energy and spiral into the nucleus.
Bohr's model (1913): electrons revolve only in certain fixed, discrete orbits (called shells or energy levels, labelled K, L, M, N…), and as long as an electron stays in one of these permitted orbits it does not radiate energy. This fixed the stability problem in Rutherford's model and is the model used for electron distribution at this level.
Distribution of electrons in shells — the Bohr–Bury scheme
NCERT uses a simplified set of rules (the Bohr–Bury scheme) to fill electrons into shells:
| Shell | n | Maximum electrons (2n²) |
|---|---|---|
| K | 1 | 2 |
| L | 2 | 8 |
| M | 3 | 18 (but capped at 8 if it is the outermost shell) |
| N | 4 | 32 |
Worked example 1 — electron distribution and valency of chlorine
Chlorine has atomic number 17. Find its electron distribution and valency.
Fill shells in order: K = 2, L = 8, remaining electrons = 17 − 2 − 8 = 7, so
M = 7.
Electron distribution: 2, 8, 7 (three shells occupied).
The outermost shell has 7 electrons — one short of a stable octet — so chlorine's
valency is 1 (it needs to gain one electron).
Worked example 2 — why potassium is 2, 8, 8, 1 and not 2, 8, 9
Potassium has atomic number 19. Why isn't its M shell filled to 9 electrons before a new shell starts, since the M shell can hold up to 18?
The rule that the outermost occupied shell cannot exceed 8 electrons applies here: after
K = 2 and L = 8, only 9 electrons remain, but the M shell (not yet the outermost, since a new
shell must open) is capped at 8 as the outer shell condition kicks in early for the 19th
electron.
Electron distribution: 2, 8, 8, 1 — the 19th electron starts a new N shell
rather than making the M shell hold 9. This is exactly why potassium behaves chemically like
sodium and lithium (all end in a single outer electron), placing it correctly in Group 1.
Atomic number, mass number and finding the number of neutrons
Mass number (A) = number of protons + number of neutrons
Number of neutrons = A − Z
Worked example 3 — finding the number of neutrons
An atom of sodium has mass number 23 and atomic number 11. How many neutrons does it have?
Neutrons = A − Z = 23 − 11 = 12 neutrons
Isotopes and isobars
Isotopes are atoms of the same element (same atomic number) with different mass numbers, because they have different numbers of neutrons. Isobars are atoms of different elements that happen to share the same mass number.
Worked example 4 — isotopes vs isobars
Isotopes: carbon-12 (6 protons, 6 neutrons) and carbon-14 (6 protons, 8 neutrons) — same element (Z = 6 for both), different mass numbers, chemically identical because chemical behaviour depends on the number of electrons/protons, not neutrons.
Isobars: calcium-40 (Z = 20) and argon-40 (Z = 18) — different elements with different atomic numbers, but both have mass number 40, purely by coincidence of their proton and neutron totals adding to the same figure.
Common mistakes that cost marks
- Filling the outer shell beyond 8 electrons without checking whether it is genuinely the outermost shell — this is the single most common error in this chapter, as shown by the potassium example above.
- Confusing isotopes with isobars — isotopes share an atomic number, isobars share a mass number. They are opposite matches.
- Treating mass number as a fixed property of an element. Mass number varies between isotopes of the same element; only the atomic number is fixed for a given element.
- Attributing Rutherford's model to Bohr, or vice versa. Rutherford gave the nucleus; Bohr gave the fixed orbits that fixed the stability problem. Keep the attribution straight in written answers.
Where this chapter appears in exams
| Exam | Typical use |
|---|---|
| CBSE Class 9 | Electron distribution diagrams, valency from distribution, isotope/isobar identification |
| CBSE Class 10 | Directly used to explain atomic number/group/period position in the periodic table chapter |
| JEE/NEET foundation | Groundwork for quantum numbers and orbital theory in Class 11 |
Practise electron distribution alongside the periodic table. Use the free calculator suite while revising this chapter to check atomic numbers and formula masses.
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