Take timed practice tests on Periodic Table and Periodicity for JEE Main and JEE Advanced with session-wise drills, score review, and explanation-led revision.
Six 20-question timed sessions plus a 60-question chapter module. Each item is original and reframed for copyright safety.
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1. The modern periodic law states that properties of elements are periodic functions of their:
Explanation: Modern Periodic Law (Moseley): properties of elements are periodic functions of their atomic number (Z), not atomic mass. This resolved anomalies in Mendeleev's table (e.g., Ar/K, Co/Ni, Te/I).
2. Elements in the same group of the periodic table have:
Explanation: Elements in the same group have the same number of valence electrons and similar chemical properties. For example, all Group 1 elements have 1 valence electron.
3. Atomic radius across a period (left to right) generally:
Explanation: Across a period: nuclear charge increases, electrons are added to the same shell, increased Zeff pulls electrons closer → atomic radius decreases from left to right.
4. The first ionisation energy (IE₁) generally increases:
Explanation: IE₁ increases: (1) across a period — higher Zeff, smaller radius; (2) up a group — smaller atoms, valence electrons closer to nucleus. Exceptions: Group 13 (IE
5. The most electronegative element in the periodic table is:
Explanation: Fluorine has the highest electronegativity (Pauling scale: 4.0) due to its very small size and high nuclear charge relative to its radius. Electronegativity: F > O > N > Cl.
6. Among halogens, Cl has higher electron affinity than F because:
Explanation: Although F is more electronegative, adding an electron to the tiny 2p orbital of F causes significant electron-electron repulsion. Cl's larger 3p orbital accommodates the extra electron with less repulsion → Cl has higher EA. EA: Cl > F > Br > I.
7. The diagonal relationship exists between:
Explanation: Diagonal relationship: elements diagonally adjacent in the periodic table have similar properties due to similar size/charge density. Pairs: Li-Mg, Be-Al, B-Si. Example: LiCl and MgCl₂ both form covalent-character halides.
8. The correct order of ionic radii is:
Explanation: Na⁺, Mg²⁺, Al³⁺ are isoelectronic (10 e⁻ each). For isoelectronic species: higher nuclear charge → smaller ionic radius. Z: Na=11, Mg=12, Al=13. Radius order: Al³⁺
9. Metallic character increases:
Explanation: Metallic character: ability to lose electrons. Increases down a group (larger atoms, lower IE) and decreases across a period (higher IE, smaller atoms). Most metallic element: Cs or Fr (bottom-left). Least metallic: F or He.
10. The second ionisation energy of Na is very large compared to its first because:
Explanation: Na: 1s²2s²2p⁶3s¹. IE₁ removes the lone 3s electron (relatively easy). IE₂ removes an electron from the filled 2p⁶ noble-gas-core configuration — extremely stable → requires enormous energy. IE₂/IE₁ ratio ≈ 9 for Na.
11. Non-metals are generally found in which part of the periodic table?
Explanation: Non-metals are concentrated in the top-right of the periodic table (high electronegativity, high IE). Metals are bottom-left. Metalloids form a staircase between them.
12. Anomalous behaviour of the first element of a group is due to:
Explanation: First members of groups (Li, Be, B, C, N, O, F) show anomalous behaviour due to: (1) exceptionally small atomic/ionic size, (2) high charge density (polarising power), (3) maximum covalency = 4 (no d-orbitals in n=2 shell). Example: BeF₂ vs MgF₂, BF₃ vs BCl₃.
13. The effective nuclear charge (Zeff) experienced by valence electrons in Na is approximately:
Explanation: Slater's rules for Na (3s electron): screening from (2s,2p) × 0.85 = 8 × 0.85 = 6.8; from (1s,2s) already counted in 2s group; from core 2 × 1.0 = 2.0. Total screening = 8.8. Zeff = 11 − 8.8 = 2.2. The approximate value using simple rules is ~2.2–2.5.