Structure of the Atom – Short Answer Type Questions
CBSE Class 9 Chemistry — Chapter 4: Structure of the Atom
A topic-wise collection of 50 short-answer questions and clear, NCERT-aligned answers — ideal for CBSE board revision.
CBSE Board Examinations — Short Answer Questions (systematic order by topic)
Subatomic Particles (Q1–Q8)
Q1
Name the three main subatomic particles and give one property of each.
Electrons (negative charge, very small mass), Protons (positive charge, ~1 u mass), Neutrons (neutral, ~1 u mass).
Q2
Which subatomic particle determines the identity of an element?
Number of protons (atomic number Z) determines the element's identity.
Q3
Where is most of the mass of an atom concentrated?
In the nucleus (protons and neutrons), which contains nearly all atomic mass.
Q4
Which particle has negligible mass compared to protons and neutrons?
The electron has negligible mass (~1/1836 of a proton) compared to nucleons.
Q5
Give the charge on proton and electron.
Proton: +1 elementary charge; Electron: −1 elementary charge.
Q6
What is a neutral atom?
An atom with equal numbers of protons and electrons so total charge is zero.
Q7
State one role of neutrons in the nucleus.
Neutrons add mass and help stabilize the nucleus by reducing proton–proton repulsion.
Q8
How are electrons held in an atom?
By electrostatic attraction to the positively charged nucleus; they occupy energy shells.
Atomic Number & Mass Number (Q9–Q16)
Q9
Define atomic number (Z).
Z is the number of protons in the nucleus of an atom.
Q10
Define mass number (A).
A is the total number of protons and neutrons in the nucleus (A = Z + neutrons).
Q11
How do you calculate the number of neutrons from A and Z?
Neutrons = A − Z.
Q12
Write isotope notation for an atom with 11 protons and 12 neutrons.
Protons 11 → Z = 11; neutrons 12 → A = 23. Notation:
^{23}_{11}Na.Q13
If an atom has A = 40 and Z = 20, how many neutrons are present?
Neutrons = 40 − 20 = 20 neutrons.
Q14
What does the notation
^{14}_{6}C tell you?Carbon atom with Z = 6 (protons) and A = 14 (total nucleons), so 8 neutrons.
Q15
True/False: Two atoms with the same mass number must be the same element.
False — different elements can have the same mass number (isobars).
Q16
Why is atomic number more fundamental than mass number?
Atomic number defines element identity and chemical behaviour; mass number varies with isotopes.
Isotopes & Related Concepts (Q17–Q24)
Q17
Define isotopes.
Isotopes are atoms of the same element (same Z) with different mass numbers (different neutrons).
Q18
Give two examples of isotopes.
Examples:
^{12}C and ^{13}C; ^{35}Cl and ^{37}Cl.Q19
Do isotopes have the same chemical properties? Why?
Yes, because chemical properties depend on electrons/atomic number, which are same for isotopes.
Q20
State one practical use of isotopes.
Medical tracers (radioisotopes) or carbon dating (¹⁴C) are common uses.
Q21
What are isobars?
Isobars are atoms of different elements with the same mass number (A) but different Z.
Q22
Are isotopes identical in mass? Explain.
No, isotopes differ in mass because they have different numbers of neutrons.
Q23
How does isotopic composition affect atomic mass shown in periodic table?
Atomic masses are weighted averages based on natural isotopic abundances, so fractional atomic masses appear.
Q24
Which of these are isotopes:
^{35}Cl and ^{37}Cl — explain.Yes — both are chlorine isotopes (same Z = 17) with different A (35 and 37).
Atomic Models & Key Experiments (Q25–Q32)
Q25
Summarize Dalton's atomic theory in one sentence.
Dalton proposed that matter is made of indivisible atoms, identical for each element and combining in fixed ratios to form compounds.
Q26
What was J.J. Thomson's contribution to atomic theory?
He discovered the electron and proposed the plum pudding model with electrons in a positive sphere.
Q27
Describe Rutherford's gold foil experiment result briefly.
Most α-particles passed through foil, but some were deflected strongly, indicating a small, dense, positive nucleus and mostly empty space.
Q28
State one limitation of the plum pudding model.
It couldn't explain the deflection of α-particles observed in Rutherford's experiment or the concentrated nucleus.
Q29
What key idea did Bohr introduce?
Bohr introduced quantized electron orbits/energy levels; electrons can jump between fixed levels emitting/absorbing specific quanta.
Q30
Why is Bohr model still taught despite limitations?
Because it explains hydrogen spectral lines simply and introduces energy-level concepts useful at Class 9 level.
Q31
What did Rutherford conclude about the size of the nucleus relative to the atom?
The nucleus is extremely small compared to the overall size of the atom; most of the atom is empty space.
Q32
Give one example of evidence for subatomic particles.
Cathode ray experiments (electrons) and Rutherford's scattering (nucleus) provided direct evidence.
Electronic Configuration & Valency (Q33–Q40)
Q33
What are the K, L and M shells? Give their maximum classical capacities.
K (n=1) can hold 2 electrons, L (n=2) 8 electrons, M (n=3) up to 18 electrons (but usually shown as up to 8 for first 20 elements).
Q34
Write electronic configuration of sodium (Z = 11).
Sodium: 2, 8, 1.
Q35
How is valency related to electronic configuration?
Valency is the number of electrons an atom gains, loses or shares to attain a stable (noble gas) configuration; it depends on outer-shell electrons.
Q36
State the valency of oxygen and explain briefly.
Oxygen has valency 2: electronic config 2,6 → needs 2 electrons to complete octet.
Q37
What is the valency of magnesium (Z = 12)?
Magnesium: 2,8,2 → valency 2 (loses two electrons to form Mg²⁺).
Q38
Explain why noble gases are chemically inert in terms of shells.
Noble gases have fully filled outer shells (stable electron configuration), so they have little tendency to gain/lose electrons.
Q39
Give electronic configuration for chlorine (Z = 17).
Chlorine: 2, 8, 7.
Q40
How many valence electrons does carbon have and what is its usual valency?
Carbon: 2,4 → 4 valence electrons; usual valency = 4 (shares four electrons to complete octet).
Calculations & Atomic Mass Concepts (Q41–Q46)
Q41
What is the atomic mass unit (u) definition used in NCERT?
1 u is defined as 1/12th the mass of one carbon-12 atom.
Q42
What is Avogadro's number and its significance?
N_A = 6.022 × 10²³; it is the number of atoms/molecules in one mole of a substance.
Q43
How many atoms are there in 2 moles of helium (He)?
Atoms = 2 × N_A = 2 × 6.022×10²³ ≈ 1.2044×10²⁴ atoms.
Q44
If sodium has atomic mass ≈23 g mol⁻¹, what mass contains 1 mole of sodium atoms?
1 mole of Na atoms has mass ≈ 23 g (molar mass equals atomic mass in g mol⁻¹).
Q45
How do you calculate the number of neutrons in
^{27}_{13}Al?Neutrons = A − Z = 27 − 13 = 14 neutrons.
Q46
Explain in one line why atomic masses in periodic table are not whole numbers.
Because listed atomic masses are weighted averages of isotopic masses based on natural abundances, giving non-integer values.
Quick Conceptual & Application (Q47–Q50)
Q47
Why do we say "most of the atom is empty space"?
Because electrons occupy a large volume around a tiny dense nucleus; Rutherford's scattering showed most α-particles pass through.
Q48
State one difference between atoms and molecules.
An atom is a single particle of an element; a molecule is two or more atoms chemically bonded together.
Q49
How are electronic configurations useful in predicting bonding?
By showing valence electrons and how many electrons an atom needs to gain/lose/share to achieve a stable configuration, thus predicting bond types.
Q50
Give two short tips to remember electron shell filling for first 20 elements.
Memorize distributions: 2,8,8,2 up to Ca and remember that 1–20 elements usually fill K,L and M shells with 2,8,8 pattern for quick recall.