Atomic Structure & Periodicity — Part 3: 35 MCQs with Explanations

Atomic number (Z) = number of protons = number of electrons in neutral atom.

Mass number (A) = number of protons (Z) + number of neutrons (N).

Isotopes have same atomic number (protons) but different mass numbers (different neutrons).

Number of neutrons = A - Z = 235 - 92 = 143.

d-subshell has l=2, with m_l = -2,-1,0,+1,+2 (5 orbitals). Each orbital holds 2 electrons, so 5 × 2 = 10 electrons.

Hund's rule: electrons occupy degenerate orbitals singly with parallel spins before pairing begins, minimizing electron-electron repulsion.

Cr has half-filled d-subshell stability: [Ar] 4s¹ 3d⁵ instead of expected [Ar] 4s² 3d⁴.

Total electrons = 10 (Ne) + 2 + 5 = 17 ⇒ chlorine (Z=17).

Energy order: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p... Due to penetration effect, 4s fills before 3d but 3d has higher energy once occupied.

Effective nuclear charge (Z_eff) increases across period, pulling electrons closer and making removal harder.

Ionization energy generally increases across period (Na < Mg > Al < Si). Mg has higher IE than Al due to s-p orbital stability, but Si has highest among these four.

Fluorine has highest electron affinity among elements (excluding chlorine which is slightly higher due to small size of F causing repulsion). Neon has zero/negative EA.

Increasing effective nuclear charge pulls electrons closer to nucleus, decreasing atomic radius.

All are isoelectronic (10 electrons). Ionic radius decreases with increasing nuclear charge: O²⁻ (8p) > Na⁺ (11p) > Mg²⁺ (12p) > Al³⁺ (13p).

Ionic bond results from electrostatic attraction between oppositely charged ions formed by complete electron transfer.

NH₃ has N-H bonds; N is highly electronegative with lone pair, enabling hydrogen bonding. HCl and H₂S have weak H-bonding; CH₄ has none.

Molecular orbital configuration of O₂: (σ1s²)(σ*1s²)(σ2s²)(σ*2s²)(σ2p_z²)(π2p_x², π2p_y²)(π*2p_x¹, π*2p_y¹). Bond order = ½(10-6) = 2. O₂ is paramagnetic.

N₂ has all electrons paired (bond order 3). O₂ and B₂ are paramagnetic (unpaired electrons); NO has odd electron.

ΔG° = -nFE°_cell, where n = moles electrons, F = Faraday constant (96485 C/mol).

Al³⁺ + 3e⁻ → Al. 3 moles electrons needed per mole Al ⇒ 3 faradays.

Fluorine has highest electronegativity (3.98 on Pauling scale).

Elements in same group have identical valence electron configuration, determining chemical behavior.

Total electrons = 18 + 1 + 5 = 24 ⇒ chromium (Z=24), which has half-filled d-subshell stability.

4th period contains 18 elements (K to Kr): 2 (s-block) + 10 (d-block) + 6 (p-block) = 18.

Poor shielding by 4f electrons causes greater effective nuclear charge, leading to gradual decrease in atomic/ionic radii across lanthanide series.

Inert pair effect (reluctance of s-electrons to participate in bonding) increases down groups 13-16, making lower oxidation states more stable.

Fluorine is more electronegative than oxygen, so O has +2 oxidation state in OF₂.

Iron (Z=26, [Ar] 4s² 3d⁶) is transition element. Zn, Cd, Hg have completely filled d-orbitals in elemental and common oxidation states.

Transition elements show variable oxidation states, form colored compounds due to d-d transitions, and often exhibit paramagnetism due to unpaired d-electrons.

Zn²⁺ has d¹⁰ configuration (all electrons paired) ⇒ diamagnetic. Others have unpaired electrons ⇒ paramagnetic.

Electron-withdrawing Cl groups stabilize conjugate base by inductive effect. More Cl atoms ⇒ stronger acid: Cl₃CCOOH (pKa ~0.7) > Cl₂CHCOOH > ClCH₂COOH > CH₃COOH (pKa 4.76).

-OH group donates electrons by resonance, activating ring and directing electrophiles to ortho/para positions.

-NO₂ withdraws electrons by resonance and induction, deactivating ring and directing electrophiles to meta position.

Oxidizing agent gets reduced. Cu²⁺ gains electrons to form Cu ⇒ oxidizing agent.

CH₃OH has O-H bond; oxygen is highly electronegative with lone pairs, enabling intermolecular hydrogen bonding.