Journey Inside the Atom — Detailed Notes
All matter is made of atoms. This chapter traces how scientists discovered that atoms are themselves made of smaller subatomic particles, and how atomic models evolved.
8.1 Roots of atomic theory
Over 2000 years ago, Acharya Kanada (India) called the smallest indivisible particle parmanu; Democritus & Leucippus (Greece) called it atomos (“indivisible”). These were ideas, not experiments. In 1808 John Dalton gave the first scientific atomic theory — matter is made of indivisible atoms.
8.2 Atomic models
Discovery of the electron: J. J. Thomson (1897) studied cathode rays in a discharge tube and found them to be streams of negatively charged particles — electrons (charge −1.602 × 10⁻¹⁹ C, taken as −1) — present in all atoms.
Thomson’s model (plum-pudding / watermelon): a sphere of positive charge with electrons embedded in it, keeping the atom neutral.
Gold foil experiment (Geiger & Marsden under Rutherford, 1911): α-particles fired at thin gold foil — most passed straight through, some deflected, a few bounced back. This disproved Thomson’s model.
Rutherford’s (nuclear/planetary) model: the atom is mostly empty space; a tiny, dense, positively charged nucleus holds nearly all the mass; electrons revolve around it (nucleus ≈ 10⁻¹⁵ m, atom ≈ 10⁻¹⁰ m — about 10⁵ times smaller). The nucleus’s positive charge comes from protons (discovered/named by Rutherford; charge +1, much heavier than electrons). Limitation: a revolving (accelerating) electron should radiate energy and spiral into the nucleus — so the model could not explain atomic stability.
Bohr’s model (1913): electrons revolve only in fixed circular paths called shells / energy levels / stationary states (K, L, M, N… or n = 1, 2, 3, 4…). In a shell an electron does not lose energy (this explains stability). Energy increases with distance from the nucleus; electrons jump between shells by absorbing/releasing a fixed amount of energy.
8.3 The neutron
James Chadwick (1932) discovered the neutron — mass ≈ a proton, no charge (symbol n⁰). Neutrons occur in the nucleus of every atom except ordinary hydrogen. Mass of an atom comes mainly from protons + neutrons (the nucleons). Neutrons also reduce proton–proton repulsion and strengthen the nuclear force.
| Particle | Symbol | Relative charge |
|---|
| Electron | e⁻ | −1 |
| Proton | p⁺ | +1 |
| Neutron | n⁰ | 0 |
8.4–8.6 Symbols, atomic number, mass number
Elements have symbols (first letter capital, second small; some from Latin — Fe, Na, K, Cu, Au, Pb, Ag). Atomic number (Z) = number of protons (= number of electrons in a neutral atom); it identifies the element. Mass number (A) = protons + neutrons (nucleons). Notation: ZASymbol (e.g. carbon is 126C).
8.7 Electron distribution (Bohr–Bury rules)
Maximum electrons in a shell = 2n² (K = 2, L = 8, M = 18…). The outermost shell holds a maximum of 8 (first shell only 2). Shells fill in order K → L → M → N. This is the atom’s electronic configuration.
8.8 Valency
The outermost shell is the valence shell; its electrons are valence electrons. A complete octet (8 electrons, or 2 for helium) is stable/unreactive. Valency = electrons gained, lost or shared to complete the octet. If <4 valence electrons → tends to lose them (valency = that number); if >4 → gains electrons (valency = 8 − number); carbon (2,4) shares 4 (valency 4).
8.9 Isotopes, average atomic mass, isobars
Isotopes — atoms of the same element (same Z) with different mass numbers (different neutrons); same chemical properties (e.g. protium/deuterium/tritium; 12C, 13C, 14C). Uses: 235U (nuclear fuel), 60Co (cancer therapy), 131I (thyroid), 14C (carbon dating). Average (weighted) atomic mass accounts for the natural abundance of isotopes (e.g. chlorine = 75% 35Cl + 25% 37Cl → 35.5 u). Isobars — atoms of different elements (different Z) with the same mass number (e.g. Ar-40, K-40, Ca-40).
Key Terms
parmanu/atomos, Dalton, electron/cathode rays, plum-pudding model, gold foil (α-scattering), nucleus, proton, Rutherford model & stability, Bohr shells (K,L,M,N / energy levels), neutron, nucleons, atomic number (Z), mass number (A), 2n², octet, valency, isotopes, average atomic mass, isobars.
Revise, Reflect, Refine — All NCERT Exercise Answers
1Gold foil experiment — choose correct options and explain.
Correct: (ii) the results disproved the plum-pudding model and led to the idea of a central nucleus, and (iii) the large deflection of a few α-particles showed that the atom’s mass and positive charge are packed into a tiny centre. Incorrect: (i) the experiment did not reveal neutrons (discovered later by Chadwick), and (iv) it did not show how electrons move — it told us about the nucleus, not electron paths.
2Which statements are correct per Bohr’s model? Give reasons.
(iii) is correct — electrons revolve in orbits of fixed energy (stationary states) without losing energy, which is exactly Bohr’s key postulate explaining stability. The others are incorrect: (i) wrongly says electrons lose energy and fall in (that was Rutherford’s problem); (ii) and (iv) wrongly allow electrons anywhere/between levels — Bohr allowed only fixed shells with definite energies.
3Species X (p=18,n=19), Y (p=17,n=18), Z (p=17,n=20). Relate (i) Y & Z (ii) Z & X.
Mass numbers: X = 18+19 = 37, Y = 17+18 = 35, Z = 17+20 = 37. (i) Y and Z have the same number of protons (17) but different mass numbers (35 and 37) → they are isotopes. (ii) Z and X have the same mass number (37) but different atomic numbers (17 and 18) → they are isobars.
4What did Rutherford conclude about the atom’s positive part from the α-particles that bounced back / deflected widely?
That the positive charge (and almost all the mass) of an atom is not spread out but is concentrated in an extremely small, dense, positively charged region at the centre — the nucleus. Most α-particles passed straight through because the atom is mostly empty space; the few that came near this tiny dense nucleus were strongly repelled and deflected or bounced back.
5Arrange the atomic models in correct chronological order.
(iv) Dalton (atoms indivisible) → (ii) Thomson (plum-pudding, electrons in positive sphere) → (iii) Rutherford (dense central nucleus) → (i) Bohr (electrons in fixed energy orbits). Each model improved on the previous one as new experimental evidence appeared.
6Electrons orbit the nucleus — why don’t they fly away? What keeps them attracted?
The nucleus is positively charged (protons) and electrons are negatively charged. The electrostatic force of attraction between the oppositely charged nucleus and electrons holds the electrons in their orbits and stops them flying away. (In Bohr’s model they stay in fixed shells without losing energy, so the atom remains stable.)
7Assertion: discovery of subatomic particles helped understand atomic structure. Reason: the number of electrons equals the number of protons in an atom. Choose the option.
(ii) Both A and R are true, but R is not the correct explanation of A. It is true that discovering electrons, protons and neutrons helped us understand atomic structure, and it is also true that a neutral atom has equal protons and electrons — but that equality is not the reason the discoveries helped us understand the structure.
8Magnesium: mass number 24, atomic number 12. Find protons, neutrons, electrons; show the electron arrangement.
Protons = Z = 12; electrons = 12 (neutral atom); neutrons = A − Z = 24 − 12 = 12. Electronic configuration: K = 2, L = 8, M = 2 → 2, 8, 2 (2 valence electrons, so valency 2).
9For the atoms shown in Fig. 8.17, find name, symbol, total electrons, valence electrons, valency, protons, atomic number.
Method (apply to each diagram): total electrons = count all electrons shown = number of protons = atomic number (Z); the element is identified from Z (e.g. Z = 6 → carbon C, Z = 8 → oxygen O, Z = 11 → sodium Na, Z = 17 → chlorine Cl). Valence electrons = electrons in the outermost shell; valency = that number if ≤ 4 (electrons lost) or 8 − (that number) if > 4 (electrons gained). Read each diagram’s shells and fill in accordingly.
10Both Rutherford and Bohr have electrons orbiting the nucleus. Why did Rutherford’s model fail but Bohr’s succeed in explaining stability?
In Rutherford’s model an orbiting electron is constantly accelerating (circular motion), so by classical physics it should continuously radiate energy, spiral inward and crash into the nucleus — the atom could not be stable. Bohr fixed this by postulating stationary states: electrons occupy only certain fixed orbits (energy levels) in which they do not radiate/lose energy, so they stay in stable orbits and the atom does not collapse.
11An atom 70X has 31 electrons. How many neutrons are in its nucleus?
Neutral atom ⇒ protons = electrons = 31, so Z = 31. Mass number A = 70. Neutrons = A − Z = 70 − 31 = 39 neutrons.
12An atom has 79 protons and mass number 197. Find (i) neutrons (ii) electrons.
(i) Neutrons = A − Z = 197 − 79 = 118. (ii) Electrons = protons = 79 (neutral atom). (This element is gold, Au.)
13Complete the table (atomic number, mass number, neutrons, protons, electrons, name).
Row 1: Z = 5, n = 6 → A = 11, p = 5, e = 5, Boron. Row 2: A = 14, e = 7 → Z = 7, p = 7, n = 7, Nitrogen. Row 3: A = 24, p = 12 → Z = 12, n = 12, e = 12, Magnesium. Row 4: Z = 15, n = 16 → A = 31, p = 15, e = 15, Phosphorus. Row 5: A = 1, n = 0 → Z = 1, p = 1, e = 1, Hydrogen.
14Element X: mass number 35, 18 neutrons. (i) electrons & protons (ii) atomic number (iii) identify X (iv) electronic configuration (v) valence electrons (vi) mass number if 2 neutrons added (vii) relation with the new atom.
(i) Protons = A − n = 35 − 18 = 17, so electrons = protons = 17. (ii) Z = 17. (iii) X = Chlorine (Cl). (iv) Configuration = 2, 8, 7. (v) 7 valence electrons (valency 1). (vi) New mass number = 35 + 2 = 37. (vii) The new atom has the same Z (17) but a different mass number, so it is an isotope of X.
15An atom has 12 protons and 12 neutrons. If all electrons are replaced by particles of the same charge but 500× heavier, what happens to its (i) atomic number (ii) atomic mass (iii) mass number (iv) overall charge?
(i) Atomic number — unchanged = 12 (depends only on protons). (ii) Atomic mass — increases, because the replacement particles are much heavier, adding appreciable mass (normally electron mass is negligible). (iii) Mass number — unchanged = 24 (counts only protons + neutrons). (iv) Overall charge — unchanged: the atom stays neutral, since the new particles carry the same charge as electrons.