Structure of the Atom – Case-based Questions with Answers

CBSE Class 9 Chemistry — Chapter 4: Structure of the Atom — 20 Case-Based Questions

Content Bank — Quick reference

Atomic number (Z) = number of protons; Mass number (A) = protons + neutrons.
Isotope notation: ^{A}_{Z}X. Neutrons = A − Z.
Electron shells (simple): K = 2, L = 8, M = 18 (for first 20 elements use K,L,M = 2,8, up to 8).
Avogadro's number: N_A = 6.022 × 10²³.

Case 1 — Isotope identification

A student finds two specimens of chlorine labelled as ³⁵Cl and ³⁷Cl in a lab. They have identical chemical reactions but slightly different masses.

1.1. What are these two specimens called and why do they show the same chemical behaviour?

They are isotopes of chlorine (same atomic number Z = 17 but different mass numbers A = 35 and 37). They show the same chemical behaviour because chemical properties depend on electron configuration (same Z) not on neutron number.

1.2. If natural chlorine contains 75% ³⁵Cl and 25% ³⁷Cl, explain why the atomic mass of chlorine in the periodic table is ≈35.5 (not a whole number).

The atomic mass is a weighted average of isotopic masses: 0.75×35 + 0.25×37 = 26.25 + 9.25 = 35.5. Natural abundance weighting produces fractional atomic masses.

Case 2 — Rutherford scattering

In a gold foil experiment, most α-particles pass straight through but a few deflect at large angles.

2.1. What conclusion about the atom's structure did Rutherford draw?

Rutherford concluded that atoms have a very small, dense, positively charged nucleus where most mass is concentrated, and most of the atom's volume is empty space, allowing most particles to pass through.

2.2. How does this explain why most of the matter is transparent to α-particles?

Because electrons occupy a large volume with low mass and the nucleus occupies tiny volume — α-particles mostly travel through empty space without collision; only occasionally they hit close to the nucleus causing large deflection.

Case 3 — Dalton vs modern atoms

A student reads Dalton's postulate that atoms are indivisible and identical for a given element.

3.1. Which modern discoveries contradict Dalton's postulates and why?

Discovery of subatomic particles (electrons, protons, neutrons) shows atoms are divisible. Discovery of isotopes shows atoms of same element can have different masses, contradicting Dalton's "identical mass" postulate.

Case 4 — Electronic configuration practice

A student has the element with atomic number 11 and must answer questions about its electrons.

4.1. Write the electronic distribution. 4.2. How many valence electrons? 4.3. Predict its common ion and valency.

For Z = 11 (Sodium): distribution = 2, 8, 1. Valence electrons = 1. Sodium commonly loses one electron to form Na⁺ (valency 1).

Case 5 — Mass number calculation

A sample of an element is represented as ^{31}_{15}X.

5.1. How many protons, neutrons and electrons are in a neutral atom?

Protons = Z = 15. Neutrons = A − Z = 31 − 15 = 16. Electrons (neutral atom) = protons = 15.

Case 6 — Using Avogadro's number

A student is asked: How many atoms are present in 0.01 mol of copper?

6.1. Compute the number of atoms.

Number = n × N_A = 0.01 × 6.022×10²³ ≈ 6.022×10²¹ atoms.

Case 7 — Bohr explanation of spectral lines

Hydrogen emission spectrum shows discrete lines (Balmer series).

7.1. Why does Bohr's model predict discrete spectral lines?

Bohr proposed quantized orbits: electrons occupy fixed energy levels. Emission occurs when an electron falls from higher to lower level, emitting a photon with energy equal to energy difference. Quantization yields discrete energies and hence discrete spectral lines.

Case 8 — Isotopic notation and neutrons

An element X has symbol ²⁷Al.

8.1. Write the isotope notation and compute neutrons if Z = 13.

Isotope is written as ^{27}_{13}Al. Neutrons = 27 − 13 = 14.

Case 9 — Comparing atoms and ions

Two species: Na and Na⁺.

9.1. Compare electrons and shells in Na and Na⁺. 9.2. Which is smaller in size and why?

Neutral Na (Z = 11) has 11 electrons: 2,8,1. Na⁺ has lost one electron → 10 electrons: 2,8. Na⁺ is smaller because losing the outer electron reduces electron–electron repulsion and the remaining electrons are pulled closer by the nucleus.

Case 10 — Nuclear vs electronic contributions

A student wonders which subatomic particles contribute to atomic mass and which to chemical reactions.

10.1. Which particles contribute most to mass? 10.2. Which determine chemical behaviour?

Protons and neutrons (nucleons) contribute most to mass (located in nucleus). Electrons determine chemical behaviour (valence electrons govern bonding and reactions).

Case 11 — Valency determination

Element Y has electronic distribution 2,8,7.

11.1. Identify the element and state its usual valency and common ion.

Distribution 2,8,7 corresponds to Chlorine (Z = 17). Valency typically 1 (gains one electron) and common ion is Cl⁻.

Case 12 — Empirical idea

A student is given atom masses and asked why chemists use moles rather than counting atoms.

12.1. Explain the mole concept briefly and its advantage.

A mole (6.022×10²³ entities) is a convenient bridge between atomic scale and lab-scale amounts. It allows chemists to relate mass to number of atoms/molecules using molar mass (g/mol), making calculations practical instead of counting enormous numbers of atoms.

Case 13 — Predicting compound formula

Magnesium (2,8,2) reacts with oxygen (2,6).

13.1. Predict the formula of the ionic compound formed and explain with electron transfer.

Mg tends to lose 2 electrons → Mg²⁺. O tends to gain 2 electrons → O²⁻. Charges balance 1:1 → formula MgO.

Case 14 — Atomic notation exercise

Student asked: Write isotope for an atom with 9 protons and 10 neutrons.

14.1. Provide isotope notation and name the element.

Protons Z = 9 → element = Fluorine (F). Neutrons = 10 → mass number A = 9 + 10 = 19. Notation: ^{19}_{9}F.

Case 15 — Electron shell filling reminder

A student memorizes "2,8,8,2" for first 20 elements and asks whether M shell ever holds more than 8.

15.1. Clarify why M shell capacity is 18 but often shown as 8 for first 20 elements.

Theoretical max for M (n=3) is 2n² = 18. But for the first 20 elements electrons fill K and L completely, and M starts filling up to 8 before d subshells come into play; hence practical shown as 8 for simplicity up to Ca.

Case 16 — Relative atomic mass concept

A student finds atomic mass of chlorine 35.45 and asks why not 35 or 37.

16.1. Explain the origin of fractional atomic mass value.

It is the weighted average of masses of naturally occurring isotopes (³⁵Cl and ³⁷Cl) weighted by their natural abundances, producing a fractional value (e.g., ≈35.45).

Case 17 — Mass of single atom estimation

Teacher asks: Estimate mass of one hydrogen atom in kg (atomic mass ≈ 1 u).

17.1. Show the method and answer approximately.

1 u ≈ 1.6605×10⁻²⁷ kg. Hydrogen atom ≈ 1 u → mass ≈ 1.66×10⁻²7 kg.

Case 18 — Chemical vs physical properties of isotopes

A blackboard shows two isotopes used in different applications: one is stable, one radioactive.

18.1. Explain why radioactive isotopes have special uses despite chemical similarity.

Radioactive isotopes undergo nuclear decay emitting radiation; this property (not chemical) is exploited in medical tracers, cancer therapy and dating. Chemical behaviour remains similar, allowing them to be incorporated into compounds for tracking.

Case 19 — Limitations of Bohr model

A student asks why Bohr model fails for larger atoms.

19.1. Give two reasons Bohr model is inadequate for multi-electron atoms.

(1) Electron–electron interactions in multi-electron atoms alter energy levels making Bohr's simple quantized orbits incorrect. (2) The wave nature and orbital shapes (s,p,d) require quantum mechanics; Bohr cannot predict fine spectral structure or chemical behaviour for many-electron atoms.

Case 20 — Exam-tip case

A student tends to mix up protons/neutrons/electrons when under time pressure.

20.1. Give three quick exam mnemonics or checks to avoid mixing them up.

(1) Remember: "P"roton = "P"ositive, located in the nucleus. (2) Electrons = "E"nergy/External (outside nucleus) and are negative. (3) Count check: Atomic number Z = protons = electrons (neutral atom). Neutrons = A − Z. Writing A and Z first helps avoid mistakes.