Class 9 Science: Journey Inside the Atom — Chapter Question Set

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parmanus; Vaisesika Sutras

1808; scientific/experimental

electron; cathode

pudding/watermelon; electrons (plums/seeds)

empty space

nucleus

10⁻¹⁰; 10⁻¹⁵

shells; stationary states (orbits)

2n²

James Chadwick; 1932

atomic; Z

mass; A

isotopes

isobars

valence

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Acharya Kanada proposed that if matter is divided again and again, one ultimately reaches indivisible particles, which he called parmanus. His ideas are recorded in the ancient Sanskrit text known as the Vaisesika Sutras.

Dalton proposed that all matter is made of tiny indivisible particles called atoms; atoms can neither be created nor destroyed in a chemical reaction. Atoms of a given element are identical, atoms of different elements differ, and atoms of different elements combine in small whole-number ratios to form compounds.

The discovery of radioactivity showed that certain elements spontaneously emit particles and energy from their atoms. This proved that atoms are divisible and are not the smallest, permanently indivisible units of matter, as Dalton's theory had assumed.

Thomson pictured the atom as a sphere of positive charge in which the negatively charged electrons are embedded, much like plums in a pudding or seeds scattered in a watermelon. In this model, the total positive charge balances the total negative charge of the electrons.

The three observations were: (i) most of the alpha particles passed straight through the gold foil without any deflection; (ii) some alpha particles were deflected by small angles; and (iii) a very few particles bounced back, deflected by very large angles of nearly 180°.

Rutherford concluded that: (i) most of the atom is empty space, since most alpha particles passed straight through; (ii) the positive charge and nearly all the mass are concentrated in a very tiny, dense nucleus at the centre, since a few particles bounced back; and (iii) the electrons revolve around the nucleus, like planets revolving around the Sun.

Any charged particle moving in a circular orbit, like a revolving electron, is accelerating and should radiate energy continuously. Such an electron would keep losing energy, spiral inward, and finally fall into the nucleus, collapsing the atom — but real atoms are stable, which this model could not explain.

Bohr postulated that electrons can revolve around the nucleus only in certain fixed shells, or stationary states, each having a definite energy. While moving in these allowed shells, electrons do not lose energy, so they never spiral into the nucleus and the atom remains stable.

The atomic number (Z) of an atom is the number of protons present in its nucleus. The mass number (A) is the total number of protons and neutrons (i.e. nucleons) present in the nucleus of the atom.

Nucleons are the particles present in the nucleus of an atom. This term includes both protons and neutrons, which together account for almost the entire mass of the atom.

The atom with atomic number 6 is carbon. In standard notation it is written as ¹²₆C, with the mass number 12 written at the top left and the atomic number 6 at the bottom left of the symbol C.

According to the Bohr-Bury rules: (i) the maximum number of electrons a shell can hold is 2n², where n is the shell number; (ii) the outermost shell cannot hold more than 8 electrons; and (iii) the shells are filled stepwise, starting from the K shell and moving outward to L, M and N shells.

Valency is the number of electrons an atom gains, loses or shares with other atoms in order to complete its octet, i.e. to have 8 electrons in its outermost shell. It therefore measures the combining capacity of the atom.

Elements whose atoms have a complete octet (8 electrons, or 2 electrons in the case of helium) in their outermost shell already have a stable configuration. Such atoms have no need to gain, lose or share electrons, which is why they are largely unreactive.

Isotopes are atoms of the same element that have the same atomic number but different mass numbers, i.e. they have equal protons but different neutrons. For example, carbon-12 and carbon-14 are isotopes of carbon, both having Z = 6 but different mass numbers.

Isobars are atoms of different elements that have the same mass number but different atomic numbers. For example, calcium-40 (Z = 20) and argon-40 (Z = 18) are isobars, since both have a mass number of 40.

Isotopes of an element have the same number of protons and hence the same number of electrons with identical electronic configurations. Since chemical properties depend mainly on the number of valence electrons, which is the same for all isotopes, they show identical chemical behaviour.

The three isotopes of hydrogen are protium (¹₁H), deuterium (²₁H) and tritium (³₁H). Each has one proton, but they differ in their number of neutrons — protium has none, deuterium has 1 and tritium has 2.

In a neutral atom, the number of protons equals the number of electrons, so it has 26 protons and 26 electrons. The number of neutrons = mass number − atomic number = 56 − 26 = 30.

Mass number A = number of protons + number of neutrons = 17 + 18 = 35. The atom is chlorine-35.

Introduction: Thomson proposed the first structural model of the atom after discovering the electron.
Key Points:
1. Model description — The atom is a sphere of positive charge with electrons distributed uniformly throughout it.
2. Analogy — It is compared to a plum pudding or a watermelon: pulp = positive charge, seeds = electrons.
3. Why discarded — It failed to explain the large-angle scattering of alpha particles in the gold foil experiment.
Conclusion: The gold foil experiment's results could not be explained by Thomson's model, requiring a new model.

Introduction: The gold foil experiment tested Thomson's model and reshaped the understanding of atomic structure.
Key Points:
1. Setup — A narrow beam of alpha particles was fired at an extremely thin gold foil.
2. Expected result — Thomson's model predicted alpha particles would pass through with only slight deflections.
3. Actual observations — Most passed undeflected, a few were deflected sharply, and some bounced straight back.
Conclusion: These surprising observations led to the nuclear model of the atom with a tiny dense nucleus.

Introduction: Rutherford interpreted the scattering pattern to reveal the atom's internal structure.
Key Points:
1. Empty space — Most alpha particles passed undeflected, so most of the atom is empty space.
2. Nucleus — Large deflections showed the positive charge and most mass are in a tiny dense nucleus.
3. Size ratio — The nucleus is about 10⁵ times smaller than the atom itself.
Conclusion: Rutherford thus proposed the planetary model with electrons revolving around a compact nucleus.

Introduction: Rutherford's model was a milestone but had a fundamental flaw regarding atomic stability.
Key Points:
1. Limitation — According to classical physics, a revolving electron accelerates and should lose energy continuously.
2. Predicted collapse — The electron would spiral inward and fall into the nucleus, collapsing the atom.
3. Bohr's remedy — Bohr postulated stationary states — electrons in fixed shells do not radiate energy at all.
Conclusion: Bohr's postulate of stationary states restored the stability that Rutherford's model could not explain.

Introduction: Bohr's 1913 model introduced quantum ideas to explain atomic structure and stability.
Key Points:
1. Fixed shells — Electrons revolve only in fixed circular paths called shells or stationary states, with definite energies.
2. No energy loss — While moving in a fixed shell, an electron does not lose energy.
3. Energy jumps — An electron shifts to another shell only by absorbing or releasing a fixed amount of energy.
4. Shell capacity — Each shell can hold only a certain number of electrons (2n²).
Conclusion: Bohr's model explained atomic stability and marked a major step in understanding the atom.

Introduction: The neutron completed the picture of the atom's nucleus and solved a mass puzzle.
Key Points:
1. Background — A helium atom has 2 protons but is about 4 times heavier than hydrogen — extra unexplained mass.
2. Discovery — James Chadwick (1932) found a neutral particle with mass nearly equal to the proton: the neutron.
3. Significance — Neutrons explained atomic masses and showed nuclei contain protons plus neutrons (except hydrogen).
Conclusion: The neutron's discovery completed the trio of subatomic particles and solved the mass-number puzzle.

Introduction: The Bohr-Bury scheme gives a simple rule for how many electrons each shell can hold.
Key Points:
1. The rule — The maximum number of electrons in the n-th shell is 2n².
2. K and L shells — K (n = 1): 2×1² = 2 electrons; L (n = 2): 2×2² = 8 electrons.
3. M shell — M (n = 3): 2×3² = 18 electrons; the outermost shell is capped at 8.
Conclusion: Electrons fill shells stepwise from K outward, giving each element a unique configuration.

Introduction: Electron distribution follows the Bohr-Bury rules of shell filling.
Key Points:
1. Sodium — Z = 11: configuration 2, 8, 1 — K-shell 2, L-shell 8, one valence electron in M.
2. Magnesium — Z = 12: configuration 2, 8, 2 — two valence electrons in the M-shell.
3. Chlorine — Z = 17: configuration 2, 8, 7 — seven valence electrons, one short of an octet.
Conclusion: These configurations determine each element's valency and chemical behaviour.

Introduction: Valency describes how many electrons an atom exchanges to achieve a stable octet.
Key Points:
1. Definition — Valency is the number of electrons gained, lost, or shared to complete the octet.
2. Sodium (2,8,1) — Losing its single valence electron gives an octet, so sodium's valency is 1.
3. Oxygen (2,6) and Carbon (2,4) — Oxygen gains 2 electrons (valency 2); carbon shares its 4 electrons (valency 4).
Conclusion: Thus the electronic configuration directly determines the valency of every element.

Introduction: Isotopes are atoms of the same element differing only in neutron number.
Key Points:
1. Definition — Isotopes have the same atomic number (Z) but different mass numbers (A).
2. Hydrogen isotopes — Protium (¹₁H, 0 neutrons), deuterium (²₁H, 1 neutron) and tritium (³₁H, 2 neutrons).
3. Carbon isotopes — Carbon-12, carbon-13 and carbon-14, each with 6 protons but 6, 7, 8 neutrons respectively.
Conclusion: Isotopes differ in physical properties but share chemical properties due to identical configurations.

Introduction: The average atomic mass must account for the natural abundance of each isotope.
Key Points:
1. Isotopic abundances — Chlorine occurs as about 75% Cl-35 and 25% Cl-37 in nature.
2. Weighted average — Average mass = (35 × 75/100) + (37 × 25/100) = 26.25 + 9.25.
3. Result — The weighted average atomic mass of chlorine is 35.5 u.
Conclusion: Note: no single chlorine atom has mass 35.5 u — it is only the natural weighted average.

Introduction: Basic atomic accounting identifies the element completely from A and neutron count.
Key Points:
1. Protons and Z — Protons = A − neutrons = 35 − 18 = 17; so atomic number Z = 17.
2. Identification — Z = 17 identifies the element as chlorine (Cl), with 17 electrons.
3. Configuration and valency — Configuration: 2, 8, 7; it needs 1 electron to complete its octet, so valency = 1.
Conclusion: Thus X is chlorine: Z = 17, configuration 2, 8, 7, valency 1.

Introduction: Isotopes and isobars are two important relationships between atomic species.
Key Points:
1. Isotopes — Same element, same Z, different A; e.g. ¹²₆C and ¹⁴₆C both have 6 protons.
2. Isobars — Different elements, same A, different Z; e.g. ⁴⁰₂₀Ca and ⁴⁰₁₈Ar both have A = 40.
3. Chemical relation — Isotopes share chemical properties; isobars belong to different elements entirely.
Conclusion: The two terms describe fundamentally different atomic relationships distinguished by Z and A.

Introduction: Counting the subatomic particles identifies the element and its electron arrangement.
Key Points:
1. Identity — 11 protons → atomic number 11 → the element is sodium (Na); mass number = 11 + 12 = 23.
2. Configuration — The 11 electrons fill as K = 2, L = 8, M = 1, i.e. configuration 2, 8, 1.
3. Valency — Losing its single M-shell electron completes the octet, so sodium's valency is 1.
Conclusion: The atom is sodium-23 with configuration 2, 8, 1 and valency 1, a soft reactive metal.

Introduction: Magnesium's subatomic composition follows directly from its Z and A values.
Key Points:
1. Protons and electrons — Z = 12 gives 12 protons and, for a neutral atom, 12 electrons.
2. Neutrons — Neutrons = A − Z = 24 − 12 = 12 in the nucleus.
3. Configuration and valency — 12 electrons fill as 2, 8, 2; losing 2 electrons completes the octet, so valency = 2.
Conclusion: Magnesium (2, 8, 2) with valency 2 explains its tendency to form Mg²⁺ ions in compounds.

(a) Since most alpha particles passed straight through, most of the atom must be empty space.
(b) The large-angle deflections and bouncing back revealed that the positive charge and nearly all the mass of the atom are concentrated in a very tiny, dense nucleus.
(c) Thomson's model spread the positive charge uniformly over the whole sphere, so it could never produce the strong repulsion needed for large deflections or bouncing back of alpha particles.
(d) A much thicker foil would give alpha particles more chances of encountering nuclei, so a larger number of particles would be scattered or deflected.

(a) The puzzle was that the nucleus of helium had extra mass that its two protons alone could not account for.
(b) The discovery of the neutron by James Chadwick in 1932 solved this puzzle.
(c) The neutron is electrically neutral (charge 0) and has a mass almost equal to that of a proton.
(d) Being uncharged, neutrons increase the separation between protons and reduce proton-proton repulsion, while strengthening the nuclear force that binds the nucleus together.

(a) Meena disagreed because a revolving electron is an accelerating charged particle and should radiate energy continuously, which would make the atom unstable.
(b) According to classical physics, it should continuously lose energy, spiral inward, and finally collapse into the nucleus, destroying the atom.
(c) Bohr resolved the problem by proposing stationary states — electrons revolve in fixed shells of definite energy and do not lose energy while in them, so they never spiral into the nucleus.
(d) The quantum mechanical model replaced Bohr's model, describing electrons as 'electron clouds' or probability regions rather than fixed orbits.

(a) Y and Z are isotopes — both have the same atomic number (17 protons) but different mass numbers (35 and 37).
(b) X and Z are isobars — they have different atomic numbers (18 and 17) but the same mass number (37).
(c) Mass number of X = 18 + 19 = 37; of Y = 17 + 18 = 35; of Z = 17 + 20 = 37.
(d) Y and Z share identical chemical properties, because as isotopes they have the same number of electrons and the same electronic configuration.

(a) X is sodium (Na), with mass number A = 23 and atomic number Z = 11.
(b) It has 11 protons and 11 electrons; the number of neutrons = 23 − 11 = 12.
(c) Its electronic configuration is 2, 8, 1 (K shell = 2, L shell = 8, M shell = 1).
(d) Its valency is 1 — by losing its single valence electron, the atom achieves the stable octet of 8 electrons in its L shell.

(a) It is called the planetary (nuclear) model of the atom; the nucleus contains all the positive charge (protons) and most of the mass of the atom.
(b) An atom is roughly 10⁻¹⁰ m across, while a nucleus is only about 10⁻¹⁵ m — so the nucleus is about 100000 (10⁵) times smaller than the atom.
(c) The nucleus would then compare to a tiny grain, like a black pepper seed of a few millimetres, placed at the centre of the ground.
(d) The model could not explain the stability of the atom — an accelerating revolving electron should radiate energy and collapse into the nucleus.

(a) Both isotopes have the same atomic number, 35 protons, but they differ in their number of neutrons (44 and 46 respectively) and hence in mass number.
(b) It is calculated as a weighted average of the masses of the isotopes according to their percentage abundances: (79 × 49.7 + 81 × 50.3)/100.
(c) Average atomic mass = 79 × 0.497 + 81 × 0.503 = 39.26 + 40.74 = 80.0 u approximately.
(d) No — both isotopes have the same number of electrons and the same valence-shell configuration, so their chemical reactions are identical.

(a) Number of protons = mass number − neutrons = 35 − 18 = 17; since the atom is neutral, it also has 17 electrons.
(b) The atomic number is 17, so the element is chlorine (Cl).
(c) Its electronic configuration is 2, 8, 7; hence it has 7 valence electrons.
(d) The new mass number is 35 + 2 = 37; the new atom (Cl-37) is an isotope of X, having the same atomic number but a different mass number.

(a) Uranium-235 (²³⁵₉₂U) is used as the fuel in nuclear reactors.
(b) Cobalt-60 (⁶⁰₂₇Co) is used in radiation therapy for the treatment of cancer.
(c) Iodine-131 (¹³¹₅₃I) is used in the treatment of goitre and thyroid cancer.
(d) Carbon-14 (¹⁴₆C) is used to determine the age of ancient fossils in archaeological and geological dating.

(a) The original atom has atomic number Z = 12 (magnesium); Z does not change after the replacement, because the atomic number depends only on the number of protons.
(b) The mass number does not change either, since it counts only protons and neutrons, and even electrons 500 times heavier add only about 0.5 u against the nucleus's 24 u.
(c) The atomic mass increases only negligibly — the heavier 'electrons' add a tiny fraction to the total mass.
(d) No, the overall charge does not change, because the replacing particles carry the same charge (−1 each) as electrons; 12 negative charges still balance the 12 protons, keeping the atom neutral.## Section I: 5 Mark Long Answer Questions

Introduction: Atomic models evolved step by step as new experiments revealed new evidence about atomic structure.
Key Points:
1. Dalton (1808) — Proposed atoms as indivisible particles — the first scientific theory of matter.
2. Thomson (1897) — Discovered electrons; proposed the plum pudding model with electrons in a positive sphere.
3. Rutherford (1911) — The gold foil experiment revealed the nucleus — mostly empty atoms with a dense positive core.
4. Bohr (1913) — Introduced fixed energy shells (stationary states) to explain atomic stability.
5. Modern model — The quantum mechanical model describes electrons as 'electron clouds' — probable regions, not fixed paths.
Conclusion: The story shows science advancing step by step through questioning, experimentation, and revision of ideas.

Introduction: Each atomic model represented progress built on the experiments and ideas of its time.
Key Points:
1. Thomson's model — Sphere of positive charge with embedded electrons; merit: first structural model explaining neutrality; limitation: failed at gold foil results.
2. Rutherford's model — Dense nucleus with orbiting electrons (planetary model); merit: explained scattering and discovered nucleus; limitation: could not explain atomic stability.
3. Bohr's model — Electrons in fixed energy shells without radiating; merit: explained stability and shell structure; limitation: later found incomplete — replaced by quantum mechanics.
4. Evidence-driven change — Each new model arose from experimental evidence the previous one could not explain.
5. Current status — The modern quantum mechanical model refines Bohr's idea with electron clouds of probability.
Conclusion: This progression illustrates how scientific models are provisional tools, improved by each new discovery.

Introduction: The gold foil experiment was the decisive test that revealed the atom's nuclear structure.
Key Points:
1. Setup — Alpha particles from a radioactive source were aimed at an extremely thin gold foil.
2. Observations — Most passed undeflected; some deflected at small angles; a few (about 1 in 8000) bounced back.
3. Conclusion — empty space — Undeflected passage proved most of the atom is empty space.
4. Conclusion — nucleus — Large deflections and rebounds proved a tiny, dense, positively charged centre holds most of the mass.
5. Model proposed — The planetary model: electrons revolve around a compact nucleus, like planets orbiting the Sun.
Conclusion: The experiment replaced Thomson's model and founded the nuclear picture of the atom that still guides chemistry.

Introduction: Bohr's model introduced quantised shells that finally explained why atoms do not collapse.
Key Points:
1. Shells and energies — Electrons revolve in fixed circular paths — shells/stationary states — each with definite energy.
2. No radiation — An electron in a fixed shell does not lose energy, resolving the collapse problem.
3. Energy jumps — Electrons absorb or release a fixed energy amount equal to the difference between two levels to jump shells.
4. Shell capacities — 2n² rule: K = 2, L = 8, M = 18; outermost shell capped at 8 electrons.
5. Example configurations — H: 1; He: 2; O: 2, 6; Ne: 2, 8; Na: 2, 8, 1; Ar: 2, 8, 8.
Conclusion: Bohr's shells explain both the stability of atoms and the systematic structure of the periodic table.

Introduction: The three subatomic particles were discovered over 35 years, each found through ingenious experiments.
Key Points:
1. Electron — Symbol e⁻, relative charge −1, negligible mass; discovered by J. J. Thomson (1897); orbits the nucleus in shells.
2. Proton — Symbol p⁺, relative charge +1, mass ≈ 1 u; discovered/named by Rutherford; resides in the nucleus.
3. Neutron — Symbol n⁰, relative charge 0, mass ≈ proton; discovered by Chadwick (1932); resides in the nucleus.
4. Neutral atom condition — Number of protons equals number of electrons, balancing total charge.
5. Mass contribution — Protons and neutrons (nucleons) carry nearly all the atom's mass; electrons are negligible.
Conclusion: Together these particles account for the charge, mass, and identity of every atom in nature.

Introduction: Atomic number and mass number are the two fundamental identifiers of any atom.
Key Points:
1. Atomic number (Z) — The number of protons in the nucleus; it uniquely identifies an element.
2. Mass number (A) — The total number of protons and neutrons (nucleons) in the nucleus.
3. Standard notation — A is written at top-left and Z at bottom-left of the element's symbol, e.g. ¹²₆C.
4. Carbon example — ¹²₆C: 6 protons, 6 electrons, 12 − 6 = 6 neutrons.
5. Sodium and chlorine — ²³₁₁Na: 11 p, 11 e, 12 n; ³⁵₁₇Cl: 17 p, 17 e, 18 n.
Conclusion: From A and Z alone, the complete proton-neutron-electron count of any neutral atom can be found.

Introduction: The Bohr-Bury rules determine how electrons occupy the shells of every atom.
Key Points:
1. Rule 1 — capacity — The n-th shell holds at most 2n² electrons: K = 2, L = 8, M = 18, N = 32.
2. Rule 2 — outer limit — The outermost shell can hold at most 8 electrons (the first shell only 2).
3. Rule 3 — filling order — Shells fill stepwise from the nucleus outward: K, then L, then M, then N.
4. Configurations (Li, C, O) — Lithium (3): 2, 1; Carbon (6): 2, 4; Oxygen (8): 2, 6.
5. Configurations (Na, Si, Ar) — Sodium (11): 2, 8, 1; Silicon (14): 2, 8, 4; Argon (18): 2, 8, 8.
Conclusion: The configuration — especially the valence shell — decides an element's chemical behaviour.

Introduction: Valency links an atom's electron arrangement to its combining capacity in compounds.
Key Points:
1. Definition — Valency is the number of electrons gained, lost or shared to complete the octet.
2. Losing electrons — Elements with fewer than 4 valence electrons lose them: Na (2,8,1) has valency 1; Mg (2,8,2) has valency 2.
3. Sharing electrons — Carbon (2,4) shares its 4 electrons, giving valency 4.
4. Gaining electrons — Elements with more than 4 gain electrons: O (2,6) has valency 2; Cl (2,8,7) has valency 1.
5. Zero valency — Elements with complete octets (Ne: 2,8; Ar: 2,8,8) are stable and have valency 0.
Conclusion: Thus the valence electron count alone determines whether an atom loses, gains, or shares electrons.

Introduction: Isotopes and isobars describe two distinct ways atomic species can be related.
Key Points:
1. Isotopes — Atoms of the same element with the same Z but different A, e.g. ¹²₆C, ¹³₆C, ¹⁴₆C.
2. Hydrogen isotopes — Protium (0 neutrons), deuterium (1 neutron) and tritium (2 neutrons) — all Z = 1.
3. Isobars — Atoms of different elements with the same A but different Z, e.g. Ca-40 (Z = 20) and Ar-40 (Z = 18).
4. Abundance of Cl — Natural chlorine is about 75% Cl-35 and 25% Cl-37.
5. Weighted average — Average atomic mass = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5 u.
Conclusion: The weighted average accounts for natural abundance, unlike a simple mean, which would wrongly give 36 u.

Introduction: Isotopes have practical applications spanning energy, medicine, and historical research.
Key Points:
1. Nuclear energy — Uranium-235 is used as fuel in nuclear reactors to generate electricity in power plants.
2. Cancer treatment — Cobalt-60 is used in radiation therapy for treating cancer patients.
3. Thyroid treatment and dating — Iodine-131 treats goitre and thyroid cancer; carbon-14 dates ancient fossils and artefacts.
4. Chemical similarity — Isotopes have identical electron counts and configurations; chemistry depends on valence electrons.
5. Physical differences — Different neutron numbers change mass, so melting/boiling points and densities differ.
Conclusion: Isotopes thus serve humanity in diverse fields while behaving identically in chemical reactions.