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1. Who first proposed the concept of atom as an indivisible particle?
ⓐ. John Dalton
ⓑ. J.J. Thomson
ⓒ. Niels Bohr
ⓓ. Rutherford
Correct Answer: John Dalton
Explanation: John Dalton in 1803 proposed the Atomic Theory which stated that atoms are indivisible, indestructible, and combine in simple whole-number ratios. Thomson later discovered the electron, Rutherford proposed the nuclear model, and Bohr refined the structure with quantized orbits.
2. Which discovery proved that the atom is not indivisible?
ⓐ. Discovery of electron
ⓑ. Discovery of photon
ⓒ. Discovery of positron
ⓓ. Discovery of neutron
Correct Answer: Discovery of electron
Explanation: The discovery of the electron by J.J. Thomson in 1897 showed that atoms are made of smaller subatomic particles, hence not indivisible. Neutrons were discovered later (Chadwick, 1932). Photons are quanta of light, not subatomic particles inside the atom.
3. What was the major drawback of Dalton’s Atomic Theory?
ⓐ. It explained isotopes correctly.
ⓑ. It explained radioactivity correctly.
ⓒ. It could not explain isotopes and subatomic particles.
ⓓ. It could explain atomic structure fully.
Correct Answer: It could not explain isotopes and subatomic particles.
Explanation: Dalton considered atoms as indivisible, identical particles for each element. But isotopes (same element with different masses) and the existence of subatomic particles (electrons, protons, neutrons) contradicted his postulates.
4. Which subatomic particle has the least mass?
ⓐ. Proton
ⓑ. Electron
ⓒ. Neutron
ⓓ. Positron
Correct Answer: Electron
Explanation: Electron mass = $9.1 \times 10^{-31}\,\text{kg}$, which is about 1/1836 the mass of a proton or neutron. Protons and neutrons are nearly equal in mass. A positron has the same mass as an electron but positive charge.
5. What is the charge-to-mass ratio (e/m) of an electron determined by J.J. Thomson?
ⓐ. $1.76 \times 10^{11}\,\text{C/kg}$
ⓑ. $9.1 \times 10^{-31}\,\text{C/kg}$
ⓒ. $1.67 \times 10^{-27}\,\text{C/kg}$
ⓓ. $3.0 \times 10^{8}\,\text{C/kg}$
Correct Answer: $1.76 \times 10^{11}\,\text{C/kg}$
Explanation: Thomson measured the deflection of cathode rays in electric and magnetic fields and calculated the e/m ratio as $1.76 \times 10^{11}\,\text{C/kg}$. This was a key step in discovering the electron’s properties.
6. Who discovered the neutron?
ⓐ. Goldstein
ⓑ. Rutherford
ⓒ. Chadwick
ⓓ. Bohr
Correct Answer: Chadwick
Explanation: James Chadwick discovered the neutron in 1932 through bombardment of beryllium with alpha particles. Neutrons are neutral particles that explain isotopes and provide nuclear stability.
7. What is the relative charge of a proton compared to an electron?
ⓐ. Equal and opposite
ⓑ. Double that of an electron
ⓒ. Half of electron’s charge
ⓓ. Zero
Correct Answer: Equal and opposite
Explanation: Proton has a charge of $+1.602 \times 10^{-19}\,C$, while electron has $-1.602 \times 10^{-19}\,C$. Magnitudes are equal but opposite in sign, ensuring charge balance in atoms.
8. Which experiment led to the discovery of the nucleus?
ⓐ. Cathode ray experiment
ⓑ. Gold foil experiment
ⓒ. Oil drop experiment
ⓓ. Photoelectric effect experiment
Correct Answer: Gold foil experiment
Explanation: Rutherford’s gold foil experiment (1909) involved alpha particles passing through thin gold foil. Most passed undeflected, but few deflected at large angles, proving a small, dense, positively charged nucleus at the atom’s center.
9. The nucleus of an atom is made up of:
ⓐ. Protons only
ⓑ. Neutrons only
ⓒ. Protons and neutrons
ⓓ. Electrons and neutrons
Correct Answer: Protons and neutrons
Explanation: The nucleus contains positively charged protons and neutral neutrons (collectively called nucleons). Electrons move around the nucleus in orbitals and do not reside inside it.
10. Which scientist’s model of the atom is often referred to as the “plum pudding model”?
ⓐ. Dalton
ⓑ. Rutherford
ⓒ. J.J. Thomson
ⓓ. Bohr
Correct Answer: J.J. Thomson
Explanation: Thomson proposed that electrons are embedded in a positively charged “soup” like plums in pudding. This model was later rejected after Rutherford’s gold foil experiment showed the existence of a dense nucleus.
11. Who discovered canal rays (positive rays) leading to the discovery of the proton?
ⓐ. J.J. Thomson
ⓑ. Eugen Goldstein
ⓒ. James Chadwick
ⓓ. Niels Bohr
Correct Answer: Eugen Goldstein
Explanation: In 1886, Eugen Goldstein discovered canal rays while working with gas discharge tubes. These rays were composed of positively charged particles, later identified as protons in hydrogen. This discovery showed that atoms contain positive particles, complementing Thomson’s discovery of electrons.
12. Which gas produced the smallest positive canal ray particle (proton) during Goldstein’s experiments?
ⓐ. Oxygen
ⓑ. Nitrogen
ⓒ. Hydrogen
ⓓ. Helium
Correct Answer: Hydrogen
Explanation: In canal ray experiments, the lightest positive particle was obtained from hydrogen gas, which was identified as the proton. Its charge is equal and opposite to that of the electron, but its mass is 1836 times greater.
13. What is the mass of a proton?
ⓐ. $9.1 \times 10^{-31}\,\text{kg}$
ⓑ. $1.67 \times 10^{-27}\,\text{kg}$
ⓒ. $1.76 \times 10^{11}\,\text{kg}$
ⓓ. $3.0 \times 10^{8}\,\text{kg}$
Correct Answer: $1.67 \times 10^{-27}\,\text{kg}$
Explanation: The proton mass is approximately $1.67 \times 10^{-27}\,\text{kg}$, much heavier than the electron ($9.1 \times 10^{-31}\,\text{kg}$). This high mass explains why protons contribute significantly to the atom’s mass.
14. What is the charge of a proton in coulombs?
ⓐ. $+1.602 \times 10^{-19}\,C$
ⓑ. $-1.602 \times 10^{-19}\,C$
ⓒ. $+9.1 \times 10^{-31}\,C$
ⓓ. $0\,C$
Correct Answer: $+1.602 \times 10^{-19}\,C$
Explanation: A proton carries a positive elementary charge equal in magnitude to that of the electron. This balance of charges ensures electrical neutrality in atoms.
15. Which experiment led to the discovery of electrons?
ⓐ. Rutherford’s scattering experiment
ⓑ. Goldstein’s canal ray experiment
ⓒ. J.J. Thomson’s cathode ray experiment
ⓓ. Millikan’s oil drop experiment
Correct Answer: J.J. Thomson’s cathode ray experiment
Explanation: J.J. Thomson in 1897 studied cathode rays in a discharge tube and concluded that they were composed of negatively charged particles called electrons. This discovery proved that atoms are divisible into smaller subatomic particles.
16. Which property of cathode rays was used by Thomson to measure the charge-to-mass ratio (e/m) of an electron?
ⓐ. Deflection in an electric and magnetic field
ⓑ. Their ability to produce fluorescence
ⓒ. Their ability to move a light paddle wheel
ⓓ. Their straight-line motion in vacuum
Correct Answer: Deflection in an electric and magnetic field
Explanation: By balancing the electric and magnetic deflections of cathode rays, Thomson calculated the e/m ratio for the electron, which was $1.76 \times 10^{11}\,\text{C/kg}$. Other listed properties demonstrated electrons’ particle nature but were not used for e/m calculation.
17. What conclusion did Thomson draw from cathode ray experiments?
ⓐ. Atoms are indivisible.
ⓑ. Atoms contain negatively charged particles.
ⓒ. Nucleus exists at the center of atom.
ⓓ. Mass of atom is concentrated in neutrons.
Correct Answer: Atoms contain negatively charged particles.
Explanation: Thomson showed that cathode rays consist of electrons, proving atoms are not indivisible. The nucleus was discovered later by Rutherford, and neutrons by Chadwick.
18. Which scientist measured the charge of an electron after Thomson’s work?
ⓐ. Rutherford
ⓑ. Millikan
ⓒ. Bohr
ⓓ. Goldstein
Correct Answer: Millikan
Explanation: Millikan’s oil drop experiment (1909) measured the charge of an electron ($-1.602 \times 10^{-19}\,C$). Combining this with Thomson’s e/m ratio, the mass of an electron was calculated.
19. Which of the following is not a property of cathode rays?
ⓐ. They produce fluorescence when striking a surface.
ⓑ. They travel in straight lines in a vacuum.
ⓒ. They are deflected towards the positive plate in an electric field.
ⓓ. They carry positive charge.
Correct Answer: They carry positive charge.
Explanation: Cathode rays are composed of electrons (negatively charged), hence are deflected towards the positive plate. Canal rays are positive, discovered by Goldstein.
20. In Thomson’s experiment, what was the value of e/m ratio obtained for the electron?
ⓐ. $1.67 \times 10^{-27}\,\text{C/kg}$
ⓑ. $9.1 \times 10^{-31}\,\text{C/kg}$
ⓒ. $1.76 \times 10^{11}\,\text{C/kg}$
ⓓ. $3.0 \times 10^{8}\,\text{C/kg}$
Correct Answer: $1.76 \times 10^{11}\,\text{C/kg}$
Explanation: Thomson calculated the charge-to-mass ratio of electrons as $1.76 \times 10^{11}\,\text{C/kg}$. This was the first quantitative evidence of a subatomic particle’s property, revolutionizing atomic theory.
21. Who discovered the neutron?
ⓐ. J.J. Thomson
ⓑ. Ernest Rutherford
ⓒ. James Chadwick
ⓓ. Niels Bohr
Correct Answer: James Chadwick
Explanation: In 1932, James Chadwick discovered the neutron by bombarding beryllium with alpha particles, producing a neutral radiation later identified as neutrons. This completed the discovery of all major subatomic particles: electrons, protons, and neutrons.
22. In Chadwick’s experiment, which element was bombarded with alpha particles to discover the neutron?
ⓐ. Carbon
ⓑ. Beryllium
ⓒ. Hydrogen
ⓓ. Oxygen
Correct Answer: Beryllium
Explanation: When beryllium was bombarded with alpha particles, a highly penetrating radiation was produced. It was not deflected by electric or magnetic fields, showing it was neutral. These neutral particles were identified as neutrons.
23. What type of radiation was observed in Chadwick’s experiment that proved the existence of neutrons?
ⓐ. Radiation deflected by electric fields
ⓑ. Neutral radiation unaffected by electric and magnetic fields
ⓒ. Positively charged radiation
ⓓ. Radiation with very low penetration power
Correct Answer: Neutral radiation unaffected by electric and magnetic fields
Explanation: The radiation produced from beryllium bombardment was highly penetrating and not influenced by electric or magnetic fields. This confirmed the existence of a neutral subatomic particle, the neutron.
24. What is the mass of a neutron compared to a proton?
ⓐ. Equal
ⓑ. Half of a proton
ⓒ. Slightly lesser
ⓓ. Slightly greater
Correct Answer: Slightly greater
Explanation: A neutron’s mass ($1.675 \times 10^{-27}\,kg$) is slightly greater than that of a proton ($1.673 \times 10^{-27}\,kg$). Despite being close in value, this small difference affects nuclear stability.
25. Why was the discovery of the neutron important?
ⓐ. It explained why atoms emit light.
ⓑ. It explained the presence of isotopes.
ⓒ. It explained cathode rays.
ⓓ. It explained the plum pudding model.
Correct Answer: It explained the presence of isotopes.
Explanation: Isotopes are atoms of the same element with equal protons but different masses. The existence of neutrons in the nucleus explained why isotopes differ in mass but not in chemical behavior.
26. Which reaction represents the production of neutrons in Chadwick’s experiment?
ⓐ. $^{9}_{4}Be + ^{4}_{2}\alpha \rightarrow ^{12}_{6}C + ^{1}_{0}n$
ⓑ. $^{14}_{7}N + ^{1}_{1}p \rightarrow ^{15}_{8}O$
ⓒ. $H_2 \rightarrow 2H^+ + 2e^-$
ⓓ. $^{40}_{19}K \rightarrow ^{40}_{20}Ca + e^-$
Correct Answer: $^{9}_{4}Be + ^{4}_{2}\alpha \rightarrow ^{12}_{6}C + ^{1}_{0}n$
Explanation: When beryllium was bombarded with alpha particles, it produced carbon and a neutral particle, identified as the neutron. This reaction directly represents Chadwick’s discovery.
27. Which property of neutrons makes them highly useful in nuclear reactions?
ⓐ. Positive charge that attracts electrons
ⓑ. High mass and neutral charge allowing deep penetration
ⓒ. Low mass and strong deflection in fields
ⓓ. Ability to emit photons
Correct Answer: High mass and neutral charge allowing deep penetration
Explanation: Neutrons, being neutral, can penetrate nuclei without being repelled by positive charges. Their high mass makes them effective projectiles in nuclear fission and other nuclear reactions.
28. What is the charge on a neutron?
ⓐ. $+1.602 \times 10^{-19}\,C$
ⓑ. $-1.602 \times 10^{-19}\,C$
ⓒ. Zero
ⓓ. Twice that of an electron
Correct Answer: Zero
Explanation: Neutrons are electrically neutral, carrying no charge. This is why the radiation in Chadwick’s experiment was not deflected by electric or magnetic fields.
29. Before the discovery of neutrons, which problem remained unresolved in atomic models?
ⓐ. Distribution of electrons in shells
ⓑ. Stability of the nucleus and isotopes
ⓒ. Quantization of energy levels
ⓓ. Nature of cathode rays
Correct Answer: Stability of the nucleus and isotopes
Explanation: Earlier models could not explain isotopes and how protons alone could stay together in the nucleus despite repulsive forces. Neutrons explained nuclear stability by providing strong nuclear force without repulsion.
30. Which of the following statements is true about neutrons?
ⓐ. Neutrons determine the element’s atomic number.
ⓑ. Neutrons determine the element’s mass number and isotopes.
ⓒ. Neutrons carry a positive charge.
ⓓ. Neutrons revolve around the nucleus.
Correct Answer: Neutrons determine the element’s mass number and isotopes.
Explanation: Neutrons contribute to the mass number ($A = Z + N$) and explain isotopes. Atomic number (Z) is determined by protons, not neutrons. Neutrons are inside the nucleus, not revolving outside.
31. Which subatomic particle has the smallest mass?
ⓐ. Proton
ⓑ. Neutron
ⓒ. Electron
ⓓ. Positron
Correct Answer: Electron
Explanation: The electron has a mass of $9.1 \times 10^{-31}\,kg$, which is about 1/1836 the mass of a proton or neutron. Both protons ($1.67 \times 10^{-27}\,kg$) and neutrons ($1.675 \times 10^{-27}\,kg$) are much heavier. A positron has the same mass as an electron but opposite charge.
32. Which particle determines the atomic number of an element?
ⓐ. Proton
ⓑ. Neutron
ⓒ. Electron
ⓓ. Nucleus
Correct Answer: Proton
Explanation: The atomic number $Z$ is defined as the number of protons in the nucleus. Electrons balance the charge but do not define $Z$. Neutrons affect mass number and isotopes, not the atomic number.
33. Which subatomic particle is responsible for chemical bonding and reactivity?
ⓐ. Proton
ⓑ. Neutron
ⓒ. Electron
ⓓ. Photon
Correct Answer: Electron
Explanation: Electrons, particularly the valence electrons in the outermost shell, determine how atoms interact chemically. Protons and neutrons remain inside the nucleus and do not participate directly in bonding.
34. Which particle has no charge but nearly the same mass as a proton?
ⓐ. Electron
ⓑ. Photon
ⓒ. Alpha particle
ⓓ. Neutron
Correct Answer: Neutron
Explanation: Neutrons have zero charge and mass $1.675 \times 10^{-27}\,kg$, slightly greater than protons. Their neutrality makes them important in nuclear stability and reactions like fission.
35. Which subatomic particle defines isotopes of an element?
ⓐ. Electron
ⓑ. Proton
ⓒ. Neutron
ⓓ. Photon
Correct Answer: Neutron
Explanation: Isotopes are atoms of the same element with the same number of protons (same atomic number) but different numbers of neutrons, leading to different mass numbers.
36. Which subatomic particle is positively charged?
ⓐ. Electron
ⓑ. Photon
ⓒ. Neutron
ⓓ. Proton
Correct Answer: Proton
Explanation: The proton carries a charge of $+1.602 \times 10^{-19}\,C$. Electrons are negative, neutrons are neutral, and photons are chargeless quanta of light.
37. What is the relative mass of a proton compared to an electron?
ⓐ. 1/1836 times
ⓑ. 1836 times
ⓒ. Equal
ⓓ. Slightly more
Correct Answer: 1836 times
Explanation: A proton is approximately 1836 times heavier than an electron. This is why the nucleus contributes nearly all of the atom’s mass while electrons contribute negligible mass.
38. Which subatomic particle contributes to both the mass and stability of the nucleus but not to the atomic number?
ⓐ. Proton
ⓑ. Positron
ⓒ. Electron
ⓓ. Neutron
Correct Answer: Neutron
Explanation: Neutrons add to the mass number and provide nuclear stability via the strong nuclear force. However, they do not change the atomic number, which depends only on protons.
39. Which of the following sets correctly matches the charge of subatomic particles?
ⓐ. Proton = negative, Electron = positive, Neutron = neutral
ⓑ. Proton = positive, Electron = negative, Neutron = neutral
ⓒ. Proton = neutral, Electron = positive, Neutron = negative
ⓓ. Proton = positive, Electron = neutral, Neutron = negative
Correct Answer: Proton = positive, Electron = negative, Neutron = neutral
Explanation: Protons carry $+1$ charge, electrons carry $-1$ charge, and neutrons are electrically neutral. This balance ensures atomic neutrality.
40. Which subatomic particle orbits outside the nucleus?
ⓐ. Proton
ⓑ. Neutron
ⓒ. Electron
ⓓ. Nucleon
Correct Answer: Electron
Explanation: Electrons move in shells/orbitals outside the nucleus, whereas protons and neutrons (collectively nucleons) remain confined inside the nucleus. The arrangement of electrons governs atomic properties and bonding.
41. Which statement best describes Thomson’s “plum pudding” model of the atom?
ⓐ. A uniformly positive sphere with electrons embedded like “plums”
ⓑ. A tiny dense nucleus with electrons orbiting in shells
ⓒ. A nucleus containing protons and neutrons with an electron cloud
ⓓ. Electrons revolving in fixed quantized orbits around the nucleus
Correct Answer: A uniformly positive sphere with electrons embedded like “plums”
Explanation: In 1904, J.J. Thomson proposed that the atom is a diffuse, uniformly positive sphere. Negatively charged electrons are embedded within it to balance charge. There is no central nucleus in this picture. Options B and C describe later nuclear/quantum models, and D belongs to Bohr’s model.
42. What did the plum pudding model successfully account for?
ⓐ. Large-angle α-particle backscattering
ⓑ. Nuclear binding energy values
ⓒ. Electrical neutrality of atoms
ⓓ. Discrete line spectra of hydrogen
Correct Answer: Electrical neutrality of atoms
Explanation: By embedding electrons ($-e$) in a diffuse positive “pudding” of total charge $+Ze$, the net charge can be zero. The model cannot explain large-angle scattering (A), nuclear energetics (B), or sharp spectral lines (D).
43. Which experimental observation most directly contradicted Thomson’s model?
ⓐ. Photoelectric effect thresholds
ⓑ. Large-angle deflection of α-particles by thin metal foils
ⓒ. Brownian motion of particles
ⓓ. Continuous radiation spectrum from hot bodies
Correct Answer: Large-angle deflection of α-particles by thin metal foils
Explanation: Rutherford’s gold-foil experiment observed rare but dramatic backscattering events. A diffuse positive charge (Thomson) would only cause small deflections. Large-angle scattering implies a compact, massive, positively charged nucleus.
44. In the plum pudding model, how is the positive charge distributed?
ⓐ. Concentrated entirely at the center
ⓑ. Localized in discrete shells
ⓒ. Confined to a thin surface layer
ⓓ. Spread uniformly throughout the atom
Correct Answer: Spread uniformly throughout the atom
Explanation: Thomson assumed a continuous “pudding” of positive charge filling the atom’s volume. A central concentration (A) is Rutherford’s nucleus; shells (B) and surface confinement (C) are not features of his model.
45. Which key limitation made the plum pudding model incompatible with atomic emission spectra?
ⓐ. It predicted negative nuclear charge
ⓑ. It lacked any electrons
ⓒ. It could not provide quantized energy levels
ⓓ. It required neutrons for stability
Correct Answer: It could not provide quantized energy levels
Explanation: Discrete line spectra (e.g., hydrogen) require quantized electronic energies. Thomson’s picture has no mechanism for quantization. The nucleus’s charge is positive (not A), electrons are present (not B), and neutrons (D) are not part of the explanation for spectra.
46. What deflection pattern for α-particles would the plum pudding model predict?
ⓐ. Many particles stopped dead at the foil
ⓑ. Frequent backscattering beyond $90^\circ$
ⓒ. Mostly small, gentle deflections with virtually no large-angle events
ⓓ. Spiral trajectories around point nuclei
Correct Answer: Mostly small, gentle deflections with virtually no large-angle events
Explanation: A diffuse positive charge spreads the Coulomb force, producing only slight path deviations. The observed rare but large deflections require a compact, highly charged center, contradicting Thomson.
47. According to Thomson’s model, most of an atom’s mass is:
ⓐ. In orbiting electrons
ⓑ. Uniformly spread with the diffuse positive charge
ⓒ. In a central neutron core
ⓓ. Confined to quantized shells
Correct Answer: Uniformly spread with the diffuse positive charge
Explanation: Without a nucleus, the positive “pudding” carried the atom’s mass in Thomson’s view. Electrons are too light (A). Options C and D reflect later nuclear/quantum ideas absent in the model.
48. Which comparison is correct regarding Thomson vs. Rutherford models?
ⓐ. Both have nuclei; Rutherford removes electrons
ⓑ. Thomson has quantized orbits; Rutherford has diffuse charge
ⓒ. Thomson has electrons only; Rutherford has no positive charge
ⓓ. Thomson: no nucleus/diffuse $+ $ charge; Rutherford: tiny dense $+$ nucleus with external electrons
Correct Answer: Thomson: no nucleus/diffuse $+ $ charge; Rutherford: tiny dense $+$ nucleus with external electrons
Explanation: Rutherford replaced the diffuse positive background with a compact nucleus. Thomson’s model lacks a nucleus and places electrons in a uniform positive matrix, not in orbits around a center.
49. For an atom with $Z$ embedded electrons in Thomson’s model, which expresses charge neutrality?
ⓐ. $+Ze + Z(-e) = 0$
ⓑ. $+Ze + (-e) = 0$
ⓒ. $+e + Z(-Ze) = 0$
ⓓ. $Z(+e) – e^2 = 0$
Correct Answer: $+Ze + Z(-e) = 0$
Explanation: The diffuse sphere carries total positive charge $+Ze$; $Z$ electrons contribute $-Ze$. Their sum is zero. Options B–D miscount charges or use incorrect forms.
50. Which direct conceptual weakness led to the rejection of the plum pudding model after scattering experiments?
ⓐ. It predicted that electrons are heavier than protons
ⓑ. It required photons to hold electrons in place
ⓒ. It provided no compact positive center to cause observed backscattering
ⓓ. It denied the existence of electric fields inside atoms
Correct Answer: It provided no compact positive center to cause observed backscattering
Explanation: Backscattering requires a small, massive, highly charged nucleus. Thomson’s diffuse positive charge cannot produce such strong localized Coulomb forces. Options A, B, and D are not features or necessary assumptions of the model.
51. In Rutherford’s gold foil experiment, what was the source of alpha particles?
ⓐ. Uranium salts
ⓑ. Radium compound
ⓒ. Polonium source
ⓓ. Thorium salts
Correct Answer: Polonium source
Explanation: Rutherford used alpha particles emitted from a polonium source. These positively charged, heavy particles were directed at a thin gold foil to probe atomic structure. Other radioactive elements also emit alpha particles but polonium was used in this experiment.
52. What was the key observation from Rutherford’s gold foil experiment?
ⓐ. All alpha particles passed straight through the foil without deflection.
ⓑ. A few alpha particles were deflected at large angles, some nearly back.
ⓒ. Alpha particles were absorbed completely by gold foil.
ⓓ. All alpha particles were strongly repelled at the surface.
Correct Answer: A few alpha particles were deflected at large angles, some nearly back.
Explanation: Most alpha particles passed undeflected, but about 1 in 8000 deflected by large angles, with a few rebounding. This could only be explained by a small, dense, positively charged nucleus at the atom’s center.
53. What conclusion did Rutherford draw from the gold foil experiment?
ⓐ. Atom is indivisible and solid.
ⓑ. Positive charge and most of mass are concentrated in a small nucleus.
ⓒ. Electrons are embedded in a positive sphere.
ⓓ. Atom has no internal structure.
Correct Answer: Positive charge and most of mass are concentrated in a small nucleus.
Explanation: The rare large deflections suggested a very small region of dense positive charge (nucleus) where most mass is concentrated, while electrons move around it in space.
54. Which observation supported that most of the atom is empty space?
ⓐ. Almost all alpha particles passed through the foil without deflection.
ⓑ. A few alpha particles rebounded.
ⓒ. Electrons are lighter than protons.
ⓓ. Isotopes exist in nature.
Correct Answer: Almost all alpha particles passed through the foil without deflection.
Explanation: Since most alpha particles passed straight through, it indicated that atoms are mostly empty space, contradicting Thomson’s “plum pudding” model.
55. According to Rutherford’s nuclear model, how are electrons arranged?
ⓐ. Embedded in a positive sphere.
ⓑ. Orbiting around the dense nucleus.
ⓒ. Fixed inside the nucleus.
ⓓ. Randomly moving in empty space.
Correct Answer: Orbiting around the dense nucleus.
Explanation: Rutherford proposed that electrons revolve around the nucleus, similar to planets orbiting the sun. This planetary model explained atomic structure but had its own limitations.
56. Which drawback of Rutherford’s model was highlighted by classical electromagnetic theory?
ⓐ. Electrons would spiral into the nucleus losing energy as radiation.
ⓑ. Electrons cannot be neutral.
ⓒ. Nucleus should contain only electrons.
ⓓ. Neutrons were not discovered yet.
Correct Answer: Electrons would spiral into the nucleus losing energy as radiation.
Explanation: According to Maxwell’s theory, accelerating charges emit radiation. Thus, revolving electrons would lose energy, spiral inward, and collapse into the nucleus — a major flaw in Rutherford’s model.
57. Why could Rutherford’s model not explain atomic stability?
ⓐ. Nucleus was too small.
ⓑ. No quantization of electron energy levels.
ⓒ. Electrons were fixed inside the nucleus.
ⓓ. No discovery of neutrons.
Correct Answer: No quantization of electron energy levels.
Explanation: The model assumed electrons revolve in orbits but did not explain why they remain in fixed paths without losing energy. Quantization, later introduced by Bohr, explained stability.
58. What fraction of alpha particles were deflected back at angles greater than $90^\circ$ in Rutherford’s experiment?
ⓐ. About 1 in 10
ⓑ. About 1 in 100
ⓒ. About 1 in 8000
ⓓ. About 1 in 2
Correct Answer: About 1 in 8000
Explanation: Very few alpha particles were deflected backward, showing that the nucleus occupies a very tiny volume compared to the entire atom. This rare event gave evidence of the dense nucleus.
59. Which feature of the atom was directly revealed by Rutherford’s experiment?
ⓐ. Existence of isotopes
ⓑ. Existence of nucleus
ⓒ. Quantization of energy levels
ⓓ. Discovery of neutron
Correct Answer: Existence of nucleus
Explanation: The experiment showed a small, dense, positively charged core (nucleus) at the center of the atom. Isotopes and neutrons were discovered later, and quantization was explained by Bohr.
60. Which statement is a limitation of Rutherford’s nuclear model?
ⓐ. It explained the nucleus well but failed to explain electron stability and atomic spectra.
ⓑ. It proved that electrons are embedded in positive charge.
ⓒ. It explained line spectra of hydrogen completely.
ⓓ. It introduced the concept of neutrons.
Correct Answer: It explained the nucleus well but failed to explain electron stability and atomic spectra.
Explanation: Rutherford’s model described the nucleus accurately but could not explain why revolving electrons remain stable or why atoms emit discrete line spectra. This limitation was resolved by Bohr’s model using quantized orbits.
61. According to Bohr’s model, electrons revolve around the nucleus in:
ⓐ. Random elliptical orbits
ⓑ. Fixed circular orbits with quantized energy
ⓒ. Spiral paths gradually falling into the nucleus
ⓓ. Uniform positive spheres
Correct Answer: Fixed circular orbits with quantized energy
Explanation: Bohr proposed that electrons revolve in discrete circular orbits called stationary states, without radiating energy. Each orbit corresponds to a quantized energy level, which solved the stability problem of Rutherford’s model.
62. What is the condition for allowed orbits in Bohr’s theory?
ⓐ. Electron energy must be maximum
ⓑ. Electron angular momentum must be quantized
ⓒ. Nucleus must be electrically neutral
ⓓ. Electron must travel at speed of light
Correct Answer: Electron angular momentum must be quantized
Explanation: Bohr stated that angular momentum of electrons is quantized: $$mvr = n\frac{h}{2\pi}, \quad n=1,2,3,…$$ This condition allows only specific orbits, explaining stability and quantized spectra.
63. In Bohr’s model, the energy of the $n^{th}$ orbit of hydrogen is given by:
ⓐ. $E_n = -\dfrac{13.6}{n^2}\,\text{eV}$
ⓑ. $E_n = \dfrac{13.6}{n}\,\text{eV}$
ⓒ. $E_n = -13.6n^2\,\text{eV}$
ⓓ. $E_n = \dfrac{13.6}{n^3}\,\text{eV}$
Correct Answer: $E_n = -\dfrac{13.6}{n^2}\,\text{eV}$
Explanation: The energy of the hydrogen atom in the $n^{th}$ orbit is inversely proportional to $n^2$. Negative sign indicates the electron is bound to the nucleus. Ground state energy is $-13.6$ eV.
64. Which transition in the hydrogen atom gives the Lyman series?
ⓐ. Transitions to $n=1$
ⓑ. Transitions to $n=2$
ⓒ. Transitions to $n=3$
ⓓ. Transitions to $n=4$
Correct Answer: Transitions to $n=1$
Explanation: Lyman series corresponds to transitions from higher levels ($n=2,3,…$) down to the first orbit ($n=1$), producing spectral lines in the ultraviolet region.
65. What is the radius of the first Bohr orbit ($n=1$) for hydrogen atom?
ⓐ. $0.529 \,\text{Å}$
ⓑ. $1.06 \,\text{Å}$
ⓒ. $0.265 \,\text{Å}$
ⓓ. $5.29 \,\text{Å}$
Correct Answer: $0.529 \,\text{Å}$
Explanation: The Bohr radius for hydrogen is given by: $$r_n = n^2 \times 0.529\,\text{Å}$$ For $n=1$, radius is $0.529\,\text{Å}$.
66. Which postulate of Bohr explains the discrete line spectrum of hydrogen?
ⓐ. Electrons revolve in random paths.
ⓑ. Energy is emitted or absorbed when electrons jump between stationary states.
ⓒ. Electrons radiate energy continuously.
ⓓ. Electrons have infinite possible energy levels.
Correct Answer: Energy is emitted or absorbed when electrons jump between stationary states.
Explanation: Bohr’s postulate states: $\Delta E = h\nu$. This explains why hydrogen shows discrete lines instead of a continuous spectrum.
67. What is the ionization energy of hydrogen atom in ground state?
ⓐ. $0\,\text{eV}$
ⓑ. $6.8\,\text{eV}$
ⓒ. $27.2\,\text{eV}$
ⓓ. $13.6\,\text{eV}$
Correct Answer: $13.6\,\text{eV}$
Explanation: Ionization energy is the energy required to remove the electron from ground state ($n=1$) to infinity. This value is $13.6\,\text{eV}$.
68. Which spectral series of hydrogen falls in the visible region?
ⓐ. Lyman series
ⓑ. Balmer series
ⓒ. Paschen series
ⓓ. Brackett series
Correct Answer: Balmer series
Explanation: Balmer series corresponds to transitions to $n=2$. These transitions produce lines in the visible region, explaining hydrogen’s visible spectral lines.
69. What is the frequency of radiation emitted when an electron jumps from $n=3$ to $n=2$ in hydrogen?
ⓐ. $\nu = \dfrac{R_Hc}{9}$
ⓑ. $\nu = \dfrac{R_Hc}{36}$
ⓒ. $\nu = R_Hc\left(\dfrac{1}{2^2}-\dfrac{1}{3^2}\right)$
ⓓ. $\nu = R_Hc\left(\dfrac{1}{3^2}-\dfrac{1}{2^2}\right)$
Correct Answer: $\nu = R_Hc\left(\dfrac{1}{2^2}-\dfrac{1}{3^2}\right)$
Explanation: Using Bohr’s frequency condition: $$\nu = R_Hc\left(\dfrac{1}{n_1^2}-\dfrac{1}{n_2^2}\right), \quad n_2 > n_1$$ or $n_1=2, n_2=3$, the frequency is $R_Hc\left(\dfrac{1}{4}-\dfrac{1}{9}\right)$.
70. Which limitation of Bohr’s model is correct?
ⓐ. It fails for hydrogen but works for complex atoms.
ⓑ. It explains Zeeman effect and fine structure.
ⓒ. It cannot explain spectra of multi-electron atoms.
ⓓ. It considers electron as a standing wave.
Correct Answer: It cannot explain spectra of multi-electron atoms.
Explanation: Bohr’s model works well only for hydrogen and hydrogen-like ions. It fails to explain finer details (fine structure, Zeeman effect, He spectrum) and was later replaced by quantum mechanical models.
71. Which of the following is a correct postulate of Bohr’s atomic model?
ⓐ. Electrons revolve randomly around the nucleus.
ⓑ. Electrons radiate energy continuously while in orbit.
ⓒ. Electrons revolve in certain permitted circular orbits without radiating energy.
ⓓ. Electrons exist only inside the nucleus.
Correct Answer: Electrons revolve in certain permitted circular orbits without radiating energy.
Explanation: Bohr postulated the existence of stationary orbits where electrons do not radiate energy, ensuring stability. Random orbits (A) and continuous radiation (B) contradict experimental spectra, and (D) is false.
72. According to Bohr’s quantization rule, the angular momentum of an electron is:
ⓐ. $mvr = nh$
ⓑ. $mvr = n\dfrac{h}{2\pi}$
ⓒ. $mvr = \dfrac{h}{2\pi n}$
ⓓ. $mvr = n^2h$
Correct Answer: $mvr = n\dfrac{h}{2\pi}$
Explanation: Bohr proposed quantized angular momentum for stable orbits: $$mvr = n\frac{h}{2\pi}, \quad n=1,2,3…$$ This condition restricts electrons to specific orbits, explaining discrete spectra.
73. Which postulate explains the stability of atoms?
ⓐ. Electrons radiate energy only in ground state.
ⓑ. Electrons emit radiation while revolving in any orbit.
ⓒ. Electrons in stationary states do not radiate energy.
ⓓ. Electrons are held by neutrons.
Correct Answer: Electrons in stationary states do not radiate energy.
Explanation: Classical physics predicted collapse of electrons into the nucleus. Bohr avoided this by postulating that electrons in allowed orbits do not emit energy, explaining atomic stability.
74. According to Bohr, when is radiation emitted or absorbed by an atom?
ⓐ. Continuously while electrons move in orbit.
ⓑ. Only when an electron jumps between two stationary states.
ⓒ. When the nucleus vibrates.
ⓓ. When protons change into neutrons.
Correct Answer: Only when an electron jumps between two stationary states.
Explanation: Bohr’s second postulate states that radiation occurs only when electrons transition between energy levels, with frequency: $$\Delta E = h\nu$$
75. What does the quantum number $n$ in Bohr’s model represent?
ⓐ. Shape of orbit
ⓑ. Size and energy of orbit
ⓒ. Spin of electron
ⓓ. Type of radiation emitted
Correct Answer: Size and energy of orbit
Explanation: The principal quantum number $n$ determines the radius and energy of an orbit. Higher $n$ means larger orbits and higher energy. Shape is linked to azimuthal quantum number, not Bohr’s model.
76. Which postulate accounts for discrete line spectra of hydrogen?
ⓐ. Angular momentum is quantized.
ⓑ. Radiation occurs during electronic transitions between energy states.
ⓒ. Electrons collapse into the nucleus after emitting photons.
ⓓ. Electrons exist only in ground state.
Correct Answer: Radiation occurs during electronic transitions between energy states.
Explanation: Bohr explained that line spectra arise because only specific energy differences are allowed, giving discrete photon frequencies when electrons jump between orbits.
77. In Bohr’s model, the energy difference between two orbits is given by:
ⓐ. $\Delta E = h\nu$
ⓑ. $\Delta E = mc^2$
ⓒ. $\Delta E = \dfrac{1}{2}mv^2$
ⓓ. $\Delta E = nh$
Correct Answer: $\Delta E = h\nu$
Explanation: The photon emitted or absorbed has energy equal to the difference between two stationary states: $$E_2 – E_1 = h\nu$$ This explains why spectral lines have fixed frequencies.
78. What prevents electrons from spiraling into the nucleus in Bohr’s model?
ⓐ. Continuous radiation of energy
ⓑ. Quantization of angular momentum and energy
ⓒ. Presence of neutrons in nucleus
ⓓ. Random orbital motion
Correct Answer: Quantization of angular momentum and energy
Explanation: Electrons remain in stable orbits due to quantized angular momentum. Without quantization, classical theory predicted collapse into the nucleus.
79. In Bohr’s model, the lowest energy state of hydrogen corresponds to:
ⓐ. $n=0$
ⓑ. $n=1$
ⓒ. $n=2$
ⓓ. $n=\infty$
Correct Answer: $n=1$
Explanation: The ground state of hydrogen is the first orbit ($n=1$), with energy $-13.6$ eV. Higher values of $n$ represent excited states, and $n=\infty$ corresponds to ionization.
80. Which of the following is NOT a postulate of Bohr’s model?
ⓐ. Electrons revolve in certain stable orbits without radiating energy.
ⓑ. Angular momentum of electron orbits is quantized.
ⓒ. Radiation is emitted/absorbed only during transitions between orbits.
ⓓ. Electrons move in elliptical orbits with varying energies.
Correct Answer: Electrons move in elliptical orbits with varying energies.
Explanation: Bohr proposed only circular orbits with quantized angular momentum. Elliptical orbits with additional quantization were later suggested by Sommerfeld.
81. Which spectral series of hydrogen lies in the ultraviolet region?
ⓐ. Balmer series
ⓑ. Lyman series
ⓒ. Paschen series
ⓓ. Brackett series
Correct Answer: Lyman series
Explanation: The Lyman series corresponds to electron transitions from higher energy levels ($n \geq 2$) down to the ground state ($n=1$). These transitions release photons in the ultraviolet region.
82. Which formula did Bohr use to explain the hydrogen line spectrum?
ⓐ. Planck’s radiation law
ⓑ. Rydberg formula
ⓒ. Einstein’s photoelectric equation
ⓓ. Rutherford scattering formula
Correct Answer: Rydberg formula
Explanation: The frequencies of hydrogen spectral lines are given by Rydberg’s formula: $$\frac{1}{\lambda} = R_H\left(\frac{1}{n_1^2} – \frac{1}{n_2^2}\right), \quad n_2>n_1$$ where $R_H$ is the Rydberg constant.
83. The Balmer series of hydrogen corresponds to transitions ending at:
ⓐ. $n=1$
ⓑ. $n=2$
ⓒ. $n=3$
ⓓ. $n=4$
Correct Answer: $n=2$
Explanation: The Balmer series consists of electron transitions from higher levels ($n \geq 3$) down to $n=2$. These lines appear in the visible spectrum.
84. Which transition produces the first line of the Balmer series?
ⓐ. $n=3 \to n=2$
ⓑ. $n=4 \to n=2$
ⓒ. $n=5 \to n=2$
ⓓ. $n=6 \to n=2$
Correct Answer: $n=3 \to n=2$
Explanation: The lowest-energy Balmer transition occurs when an electron falls from the third orbit to the second. This line is observed in the red region of the visible spectrum.
85. Which hydrogen series falls in the infrared region?
ⓐ. Lyman series
ⓑ. Balmer series
ⓒ. Paschen series
ⓓ. All of the above
Correct Answer: Paschen series
Explanation: The Paschen series corresponds to transitions from higher levels ($n \geq 4$) down to $n=3$. These lines fall in the infrared region.
86. The frequency of radiation emitted when an electron falls from $n=4$ to $n=2$ is given by:
ⓐ. $\nu = R_Hc\left(\dfrac{1}{4} – \dfrac{1}{2}\right)$
ⓑ. $\nu = R_Hc\left(\dfrac{1}{2^2} – \dfrac{1}{4^2}\right)$
ⓒ. $\nu = R_Hc\left(\dfrac{1}{4^2} – \dfrac{1}{2^2}\right)$
ⓓ. $\nu = R_Hc\left(\dfrac{1}{3^2} – \dfrac{1}{4^2}\right)$
Correct Answer: $\nu = R_Hc\left(\dfrac{1}{2^2} – \dfrac{1}{4^2}\right)$
Explanation: Using the Rydberg formula, the frequency for a transition $n_2 \to n_1$ is proportional to $\dfrac{1}{n_1^2} – \dfrac{1}{n_2^2}$. For $n_1=2, n_2=4$, we get $\nu = R_Hc\left(\dfrac{1}{4}-\dfrac{1}{16}\right)$.
87. What is the wavelength of the first line of the Lyman series ($n=2 \to n=1$)?
ⓐ. 1216 Å
ⓑ. 4861 Å
ⓒ. 6563 Å
ⓓ. 1026 Å
Correct Answer: 1216 Å
Explanation: The Lyman-$\alpha$ line arises from $n=2 \to n=1$ transition. Its wavelength is 1216 Å (in the ultraviolet region). The other values correspond to Balmer series lines.
88. Why does hydrogen emit line spectra instead of a continuous spectrum?
ⓐ. Electrons emit energy at random frequencies.
ⓑ. Only certain electron transitions with quantized energy differences are allowed.
ⓒ. The nucleus vibrates and emits lines.
ⓓ. Hydrogen atoms absorb all frequencies equally.
Correct Answer: Only certain electron transitions with quantized energy differences are allowed.
Explanation: Bohr’s postulates state that energy levels are quantized. When electrons jump between these levels, only specific energy differences ($\Delta E$) are released, producing line spectra.
89. Which transition in hydrogen corresponds to the Brackett series?
ⓐ. To $n=1$
ⓑ. To $n=2$
ⓒ. To $n=3$
ⓓ. To $n=4$
Correct Answer: To $n=4$
Explanation: The Brackett series is observed when electrons fall from higher levels ($n \geq 5$) down to $n=4$. These spectral lines lie in the infrared region.
90. The energy of emitted photon in hydrogen spectrum is related to:
ⓐ. Energy difference between nucleus and proton
ⓑ. Energy difference between two electron orbits
ⓒ. Energy of ground state only
ⓓ. Energy absorbed by nucleus
Correct Answer: Energy difference between two electron orbits
Explanation: According to Bohr, $$\Delta E = h\nu = E_{n_2} – E_{n_1}$$
91. Which of the following is a major limitation of Bohr’s model?
ⓐ. It explains hydrogen spectrum completely.
ⓑ. It includes wave nature of electrons.
ⓒ. It accounts for Zeeman effect in full detail.
ⓓ. It cannot explain spectra of multi-electron atoms.
Correct Answer: It cannot explain spectra of multi-electron atoms.
Explanation: Bohr’s model works well for hydrogen and hydrogen-like ions, but fails for atoms containing more than one electron. It could not explain the fine details of their spectral lines.
92. Why could Bohr’s model not explain the Zeeman effect?
ⓐ. It ignored nuclear charge.
ⓑ. It did not include electron spin and magnetic interactions.
ⓒ. It assumed elliptical orbits.
ⓓ. It used incorrect value of Planck’s constant.
Correct Answer: It did not include electron spin and magnetic interactions.
Explanation: Zeeman effect (splitting of spectral lines in a magnetic field) requires consideration of spin and orbital magnetic moments. Bohr’s model only treated electrons as orbiting particles without intrinsic spin.
93. Which assumption of Bohr’s model contradicts the uncertainty principle?
ⓐ. Electrons have wave nature.
ⓑ. Electrons occupy quantized energy levels.
ⓒ. Electrons emit radiation when moving in orbits.
ⓓ. Electrons revolve in well-defined circular orbits with known radius and velocity.
Correct Answer: Electrons revolve in well-defined circular orbits with known radius and velocity.
Explanation: According to Heisenberg’s uncertainty principle, position and momentum cannot be simultaneously known with certainty. Bohr’s model violates this by assigning exact orbits and velocities.
94. Bohr’s model could not explain the fine structure of hydrogen spectrum because:
ⓐ. It did not consider relativistic corrections and electron spin.
ⓑ. It assumed elliptical orbits.
ⓒ. It ignored neutrons.
ⓓ. It used wrong values for electron mass.
Correct Answer: It did not consider relativistic corrections and electron spin.
Explanation: Precise measurements of hydrogen lines show fine splitting due to spin–orbit coupling and relativistic effects. Bohr’s simple orbital model cannot account for these details.
95. Which limitation of Bohr’s model was later resolved by quantum mechanics?
ⓐ. Failure to explain isotopes.
ⓑ. Failure to explain discrete energy levels.
ⓒ. Failure to explain stability of nucleus.
ⓓ. Failure to explain wave–particle duality of electron.
Correct Answer: Failure to explain wave–particle duality of electron.
Explanation: Bohr’s model treated electrons as classical particles. Quantum mechanics (Schrödinger, de Broglie) introduced wave nature, standing wave patterns, and probability distributions, resolving this limitation.
96. Which experimental fact could not be explained by Bohr’s model?
ⓐ. Line spectra of hydrogen atom
ⓑ. Photoelectric effect
ⓒ. Stark effect (splitting of spectral lines in electric field)
ⓓ. Energy quantization in atoms
Correct Answer: Stark effect (splitting of spectral lines in electric field)
Explanation: Bohr’s model explained gross features of hydrogen spectrum but not Stark or Zeeman effects, which require advanced quantum mechanical treatment.
97. Why is Bohr’s model called semi-classical?
ⓐ. It uses only classical mechanics.
ⓑ. It uses only quantum mechanics.
ⓒ. It combines classical electron orbits with quantum postulates.
ⓓ. It ignores Planck’s theory.
Correct Answer: It combines classical electron orbits with quantum postulates.
Explanation: Bohr’s model assumes classical circular orbits but imposes quantization of angular momentum using Planck’s constant, making it semi-classical.
98. Which of the following could not be justified by Bohr’s model?
ⓐ. Hydrogen emission spectrum
ⓑ. Ionization energy of hydrogen
ⓒ. Energy of stationary states
ⓓ. Relative intensities of spectral lines
Correct Answer: Relative intensities of spectral lines
Explanation: Bohr calculated energies and frequencies correctly but could not predict why some lines are brighter than others. Quantum mechanics with transition probabilities solved this.
99. Bohr’s model fails for heavier atoms because:
ⓐ. Nucleus is not positively charged.
ⓑ. Electron–electron repulsions and shielding are not considered.
ⓒ. Planck’s constant does not apply to heavy atoms.
ⓓ. Electrons in heavy atoms have no angular momentum.
Correct Answer: Electron–electron repulsions and shielding are not considered.
Explanation: In multi-electron atoms, inner electrons shield outer ones from full nuclear charge. Bohr’s model ignores these interactions, so it fails for heavy atoms.
100. Which statement best summarizes the limitation of Bohr’s theory?
ⓐ. It is accurate for all atoms in explaining spectral lines.
ⓑ. It cannot account for fine details like splitting of lines or multi-electron spectra.
ⓒ. It denies the existence of quantization.
ⓓ. It considers electrons as probability clouds.
Correct Answer: It cannot account for fine details like splitting of lines or multi-electron spectra.
Explanation: While Bohr’s model revolutionized atomic theory, it was limited to hydrogen-like species. Modern quantum mechanics explains finer details like spin, relativistic effects, and multi-electron interactions.
Welcome to Class 11 Chemistry MCQs – Chapter 2: Structure of Atom (Part 1). This chapter from the NCERT/CBSE Class 11 Chemistry syllabus is a turning point in learning, as it bridges classical models of the atom with the modern quantum mechanical framework. It explains how scientists discovered the electron, proton, and neutron, developed atomic models, and eventually arrived at the modern picture of the atom. These concepts are essential for board examinations, JEE Main, JEE Advanced, NEET, and state-level competitive tests. A solid understanding of this chapter also supports advanced topics in Chemical Bonding, Periodicity, Organic Chemistry, and Quantum Chemistry.
These MCQs are universally relevant since atomic structure is a fundamental concept in Chemistry across the world.
Navigation & pages: The chapter contains 475 MCQs carefully arranged into 5 parts (100 + 100 + 100 + 100 + 75). Part 1 includes the first 100 MCQs, further distributed into 10 pages with 10 questions per page. Use the page numbers above to view questions and the Part buttons above to continue with the next sets.
What you will learn & practice
- Discovery of subatomic particles: electron, proton, neutron
- Thomson’s atomic model and its limitations
- Rutherford’s α-particle scattering experiment and nuclear model
- Bohr’s atomic model — postulates, energy levels, and limitations
- Electromagnetic radiation, photoelectric effect, and hydrogen spectrum
- de Broglie’s concept of matter waves and Heisenberg’s Uncertainty Principle
- Quantum mechanical model of the atom
- Quantum numbers (n, l, m, s) and their significance
- Shapes of orbitals (s, p, d, f)
- Rules of electronic configuration: Aufbau principle, Pauli’s Exclusion Principle, Hund’s Rule
- Stability of half-filled and fully filled orbitals
- Numerical problems based on wavelength, frequency, energy of photons, ionization enthalpy, etc.
How this practice works
- Click an option to check instantly: green dot = correct, red icon = incorrect. The Correct Answer with explanation appears immediately.
- Use the 👁️ Eye icon to reveal the solution without attempting.
- Use the 📝 Notebook icon to jot quick notes (not saved permanently).
- Use the ⚠️ Alert icon to report an error directly to us.
- Use the 💬 Message icon to ask doubts, share ideas, or start a discussion.
Real value: These MCQs are strictly aligned with the NCERT/CBSE syllabus, crafted from previous-year exam questions, and supported by clear, exam-focused explanations. They are highly effective for concept mastery, one-mark MCQ practice, and fast revision before exams.
The complete chapter contains 475 solved MCQs divided into 5 structured parts. This page (Part 1) presents the first 100 MCQs with answers and explanations. Explore more with: Class 11 Chemistry – Chapter Index, Class 11 – Subject Index, All Classes – MCQ Library.
- 👉 Total MCQs in this chapter: 475 (100 + 100 + 100 + 100 + 75)
- 👉 This page: First 100 MCQs with answers & explanations (in 10 pages)
- 👉 Best for: Boards • JEE/NEET • atomic theory practice • quick revision • Chemistry quizzes
- 👉 Next: Use the Part buttons and page numbers above to continue
FAQs on Structure of Atom ▼
▸ What are Structure of Atom MCQs in Class 11 Chemistry?
These are multiple-choice questions from Chapter 2 of NCERT Class 11 Chemistry – Structure of Atom. They test key concepts like subatomic particles, atomic models, quantum numbers, orbitals, and electronic configurations.
▸ How many MCQs are available in this chapter?
There are a total of 475 MCQs from Structure of Atom. They are divided into 5 sets – four sets of 100 questions each and one set of 75 questions.
▸ Are these Chemistry MCQs useful for NCERT and CBSE board exams?
Yes, these MCQs are based on NCERT/CBSE syllabus and are very useful for Class 11 students preparing for school and state board exams.
▸ Are Structure of Atom MCQs important for JEE and NEET?
Yes, this chapter is very important for JEE, NEET, and other entrance exams as it forms the foundation of modern atomic theory and quantum mechanics.
▸ Do these MCQs include answers and explanations?
Yes, every MCQ is provided with the correct answer and supported with explanations wherever necessary, ensuring clear understanding of the concepts.
▸ Who should practice Structure of Atom MCQs?
These MCQs are useful for Class 11 students, CBSE/state board candidates, and aspirants preparing for JEE, NEET, NDA, UPSC, and other competitive exams.
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Yes, solving these MCQs regularly helps in quick revision, improves memory retention, and boosts exam performance by strengthening accuracy and speed.
▸ Do these MCQs cover both basics and advanced topics?
Yes, the MCQs cover a wide range of topics from the discovery of subatomic particles to advanced concepts like the quantum mechanical model, orbitals, and electron configurations.
▸ Do these MCQs include questions on Bohr’s atomic model?
Yes, several MCQs are based on Bohr’s model of the atom, testing concepts like quantized energy levels, hydrogen spectrum, and limitations of the model.
▸ Are there MCQs on quantum numbers and orbitals?
Yes, many MCQs focus on quantum numbers, orbital shapes (s, p, d, f), and electron distribution, which are essential for understanding modern atomic theory.
▸ Are electronic configuration and anomalous configurations included?
Yes, MCQs include standard and anomalous electronic configurations of elements such as Cr and Cu, which are frequently asked in competitive exams.
▸ Why are the 475 MCQs divided into 5 parts?
The MCQs are divided into 5 sets to make practice structured and manageable. This helps students revise step by step without getting overloaded.
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Yes, teachers and coaching centers can use these MCQs as ready-made quizzes, assignments, or test material for students.
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