Structure of Atom — Class 11 MCQs with Answers

Structure of Atom — important questions & MCQs with answers (Class 11 Chemistry)

6 solved questions from this chapter's difficulty levels, each with its answer and explanation. Tap an option to check yourself.

0 of 6 answered

1. Q1 of 6Easy

   Charge on a proton:

   A+1 CB−1.6×10⁻¹⁹ CC0D+1.6×10⁻¹⁹ C
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   Answer: +1.6×10⁻¹⁹ C

   Explanation: A proton carries the elementary charge, +1.6×10⁻¹⁹ C — the same size as the electron's charge but opposite sign. Picking −1.6×10⁻¹⁹ C swaps it with the electron; 0 C is the neutron's charge.

2. Q2 of 6Easy

   Maximum number of electrons in a shell with principal quantum number n:

   A2n²Bn²CnD2n
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   Answer: 2n²

   Explanation: A shell of principal quantum number n holds n² orbitals, and each orbital takes 2 electrons of opposite spin, so capacity is 2n². For n=2 that is 8 electrons; quoting the bare n² = 4 counts the orbitals and forgets to pair them.

3. Q3 of 6Medium

   Energy of nth Bohr orbit of H: En = −13.6/n² eV. E₂ =

   A−13.6 eVB−1.51 eVC−6.8 eVD−3.4 eV
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   Answer: −3.4 eV

   Explanation: Bohr energy scales as En = −13.6/n² eV, so at n = 2, E₂ = −13.6/4 = −3.4 eV. Forgetting to square n before dividing gives the wrong −6.8 eV instead.

4. Q4 of 6Medium

   Principal quantum number (n) describes:

   ASpinBEnergy level / shellCSubshell shapeDMagnetic orientation
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   Answer: Energy level / shell

   Explanation: The principal quantum number n specifies the shell or main energy level and roughly the electron's average distance from the nucleus, taking values 1, 2, 3, … Spin, subshell shape, and orientation are ms, l, m_l instead.

5. Q5 of 6Hard

   Energy required for electron transition n=2 to n=4 in H atom:

   A2.55 eVB10.2 eVC3.4 eVD1.51 eV
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   Answer: 2.55 eV

   Explanation: ΔE = 13.6(1/n₁² − 1/n₂²) = 13.6(1/4 − 1/16) = 13.6 × 3/16 = 2.55 eV. Using 1/16 − 1/4 instead (a negative result) or forgetting to subtract the fractions first are the usual arithmetic slips.

6. Q6 of 6Hard

   Pauli exclusion principle: no two electrons can have:

   AAll four quantum numbers sameBSame chargeCSame orbitDSame spin always
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   Answer: All four quantum numbers same

   Explanation: Pauli's exclusion principle forbids any two electrons in the same atom from sharing all four quantum numbers (n, l, m_l, ms) simultaneously — at minimum their spins must differ. Each orbital holds at most 2 electrons, of opposite spin.

Structure of Atom notes: key concepts (Class 11 Chemistry)

Building from Thomson, Rutherford and Bohr, this chapter replaces classical pictures with the quantum-mechanical atom. It links line spectra, quantised energy, dual nature of matter and Heisenberg's uncertainty to orbitals, four quantum numbers and the Aufbau, Pauli and Hund rules that fix electronic configuration.

Discovery of the electron

Cathode-ray experiments gave the electron's charge-to-mass ratio; Millikan's oil-drop fixed its charge as 1.6×10⁻¹⁹ C, so e/m yields the mass.

Rutherford nuclear model

α-scattering showed the atom is mostly empty with a tiny, dense, positively charged nucleus; electrons revolve around it like planets round the Sun.

Atomic and mass number

Atomic number Z is the proton count and fixes the element; mass number A = protons + neutrons. Isotopes share Z but differ in A (e.g. ¹H, ²H, ³H).

Bohr model of the atom

Electrons occupy fixed circular orbits of quantised angular momentum mvr = nh/2π; they neither absorb nor radiate energy while in a stationary state.

Quantisation of energy

Orbit energy is Eₙ = −13.6 Z²/n² eV; it is negative and rises toward zero as n grows, so ionisation frees the electron from n = 1 for hydrogen.

Hydrogen line spectrum

A photon is emitted or absorbed only when an electron jumps between levels, ΔE = hν; this explains the discrete Lyman, Balmer and Paschen line series.

Limitations of Bohr's model

Bohr fits only one-electron species; it fails for multi-electron atoms, ignores wave nature and uncertainty, and cannot explain fine structure.

Dual nature of matter

Every moving particle has a wavelength λ = h/mv; it is significant only for tiny masses, giving electrons a wave character that underlies the quantum model.

Heisenberg uncertainty

Position and momentum cannot both be known exactly: Δx·Δp ≥ h/4π. This forbids Bohr's precise orbits and replaces them with probability orbitals.

Four quantum numbers

Four numbers label an electron: n (shell, size/energy), l (subshell/shape, 0 to n−1), mₗ (orientation, −l to +l) and mₛ (spin, ±½).

Orbitals and their shapes

An orbital is a region of high probability, not a path: s is spherical, p is dumb-bell shaped, and d orbitals are double dumb-bells (except d_z²).

Nodes in orbitals

Nodes are zero-probability surfaces; an orbital has (n−1) total nodes, with l angular nodes and (n−l−1) radial nodes for any given orbital.

Aufbau principle

Electrons fill the lowest-energy orbital first, ordered by increasing (n + l); for equal (n + l) the lower n fills first, so 4s fills before 3d.

Pauli exclusion and Hund's rule

Pauli bars two electrons from sharing all four quantum numbers, capping an orbital at two paired spins; Hund fills degenerate orbitals singly before pairing.