Short Q&A for Periodic Classification of Elements (Class 11 NCERT)

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Periodic Classification of Elements

112 Short Q&A • NCERT Class 11 Chemistry

Class 11 JEE / NEET

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112 Short Q&A for Periodic Classification

Part 1: Genesis of Classification & Modern Periodic Table
1. Q1: Who first classified elements into triads?
   A: Johann Dobereiner.
2. Q2: State Dobereiner's Law of Triads.
   A: The atomic mass of the middle element is approximately the arithmetic mean of the other two.
3. Q3: Who proposed the Law of Octaves?
   A: John Newlands.
4. Q4: What was the basis of Newlands' classification?
   A: Increasing atomic masses (properties repeated every eighth element).
5. Q5: What was the major criterion for Mendeleev's classification?
   A: Atomic weight (mass) and similarity in chemical properties (oxide and hydride formulas).
6. Q6: State Mendeleev's Periodic Law.
   A: The properties of elements are a periodic function of their atomic weights.
7. Q7: Name an element whose existence was predicted by Mendeleev as 'Eka-silicon'.
   A: Germanium (Ge).
8. Q8: What was the main drawback of Mendeleev's Periodic Table regarding position?
   A: The position of isotopes and anomalous pairs (e.g., Ar before K).
9. Q9: Who proposed the Modern Periodic Law?
   A: Henry Moseley.
10. Q10: What is the basis of the Modern Periodic Table?
    A: Atomic Number (Z).
11. Q11: State the Modern Periodic Law.
    A: The properties of elements are a periodic function of their atomic numbers.
12. Q12: How many groups and periods are in the Modern Periodic Table?
    A: 18 Groups (vertical columns) and 7 Periods (horizontal rows).
13. Q13: What does the Period number correspond to?
    A: The Principal Quantum Number (n) of the outermost shell.
14. Q14: What is the common feature for elements in the same group?
    A: Same number of valence electrons (similar outermost electronic configuration).
15. Q15: Which period is the shortest?
    A: Period 1 (contains only 2 elements, H and He).
16. Q16: How many elements are in the 6th period?
    A: 32 elements (6s, 4f, 5d, 6p).
17. Q17: Which element is the starting point of the 4th period?
    A: Potassium (K, Z=19).
18. Q18: What is the group number for noble gases?
    A: Group 18 (or 0 group).
19. Q19: What is the group number for alkali metals?
    A: Group 1.
20. Q20: What is the group number for halogens?
    A: Group 17.
Part 2: Classification by Blocks (s, p, d, f)
21. Q21: What determines the block an element belongs to?
    A: The orbital that receives the last differentiating electron.
22. Q22: What is the general electronic configuration of s-block elements?
    A: ns¹⁻².
23. Q23: Name the groups that constitute the s-block.
    A: Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals).
24. Q24: What is the general electronic configuration of p-block elements?
    A: ns²np¹⁻⁶.
25. Q25: Which group is the only non-metal in the s-block?
    A: Hydrogen (H).
26. Q26: Which block contains maximum number of metals, non-metals, and metalloids?
    A: p-block.
27. Q27: What is the general electronic configuration of d-block elements?
    A: (n−1)d¹⁻¹⁰ns⁰⁻².
28. Q28: What is another name for d-block elements?
    A: Transition Elements.
29. Q29: Name the groups that constitute the d-block.
    A: Group 3 to Group 12.
30. Q30: What is the general electronic configuration of f-block elements?
    A: (n−2)f¹⁻¹⁴(n−1)d⁰⁻¹ns².
31. Q31: What is another name for f-block elements?
    A: Inner-Transition Elements.
32. Q32: Name the two series of f-block elements.
    A: Lanthanoids (Lanthanides) and Actinoids (Actinides).
33. Q33: What is the maximum number of electrons an f-subshell can accommodate?
    A: 14 electrons.
34. Q34: To which block does the element with Z=30 (Zn) belong?
    A: d-block (Group 12).
35. Q35: To which block does the element with Z=54 (Xe) belong?
    A: p-block (Group 18).
36. Q36: What is the group and period of the element with Z=17 (Cl)?
    A: Group 17, Period 3.
37. Q37: What is the common oxidation state of all Lanthanoids?
    A: +3.
38. Q38: What are s-block elements generally known for (electro-)?
    A: They are highly Electropositive.
39. Q39: What are p-block elements generally known for (electro-)?
    A: They are often highly Electronegative.
40. Q40: What are the elements after Z=92 called?
    A: Transuranium Elements (all are radioactive).
Part 3: Atomic and Ionic Radii
41. Q41: Define Atomic Radius.
    A: The distance from the center of the nucleus to the outermost shell of electrons.
42. Q42: Define Covalent Radius.
    A: Half the distance between the nuclei of two identical atoms covalently bonded.
43. Q43: Define Metallic Radius.
    A: Half the internuclear distance between the nuclei of two adjacent metal atoms in a metallic crystal.
44. Q44: Define van der Waals Radius.
    A: Half the distance between the nuclei of two non-bonded isolated atoms.
45. Q45: How does Atomic Radius vary across a Period (L → R)?
    A: Decreases.
46. Q46: Why does atomic radius decrease across a period?
    A: Increase in Effective Nuclear Charge (Z_eff) (valence electrons remain in the same shell).
47. Q47: How does Atomic Radius vary down a Group (T → B)?
    A: Increases.
48. Q48: Why does atomic radius increase down a group?
    A: Increase in the number of electron shells (Principal Quantum Number, n).
49. Q49: Which is larger: a parent atom (A) or its Cation (A⁺)?
    A: Parent atom (A) (Cation is smaller).
50. Q50: Why is a cation smaller than its parent atom?
    A: Higher Z_eff and loss of the outermost shell (sometimes).
51. Q51: Which is larger: a parent atom (A) or its Anion (A⁻)?
    A: Anion (A⁻) (Anion is larger).
52. Q52: Why is an anion larger than its parent atom?
    A: Increased electron-electron repulsion (decreases Z_eff per electron).
53. Q53: Define Isoelectronic Species.
    A: Atoms or ions having the same number of electrons.
54. Q54: Arrange O²⁻, F⁻, Na⁺, Mg²⁺ in order of increasing ionic size.
    A: Mg²⁺ < Na⁺ < F⁻ < O²⁻ (size decreases as Z increases for isoelectronics).
55. Q55: What is the main factor determining the size of isoelectronic species?
    A: Nuclear Charge (Z) (Higher Z, smaller size).
56. Q56: Which radius is generally the largest among the four types?
    A: van der Waals radius.
57. Q57: What is Shielding Effect (or Screening Effect)?
    A: The reduction in the nuclear attraction felt by the valence electrons due to the presence of inner-shell electrons.
58. Q58: Which orbital type provides the least shielding effect?
    A: f-orbitals.
59. Q59: How does Z_eff generally change down a group?
    A: It remains nearly constant (increase in Z is offset by increased shielding).
60. Q60: How does Z_eff generally change across a period (L → R)?
    A: It increases (charge increases, shielding is almost constant).
Part 4: Ionization Enthalpy (IE)
61. Q61: Define Ionization Enthalpy (Δ_iH).
    A: The minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
62. Q62: Is the first ionization enthalpy generally endothermic or exothermic?
    A: Endothermic (energy is required, Δ_iH > 0).
63. Q63: Why is IE₂ always greater than IE₁?
    A: It is harder to remove an electron from an already stable cation (higher Z_eff).
64. Q64: How does IE generally vary across a Period (L → R)?
    A: Increases.
65. Q65: Why does IE increase across a period?
    A: Increase in Z_eff and decrease in atomic size.
66. Q66: How does IE generally vary down a Group (T → B)?
    A: Decreases.
67. Q67: Why does IE decrease down a group?
    A: Increase in atomic size and increased shielding.
68. Q68: Which elements have the highest ionization enthalpy?
    A: Noble Gases (due to stable, fully-filled configuration).
69. Q69: Which Group has the lowest ionization enthalpy?
    A: Group 1 (Alkali Metals).
70. Q70: Which element has the highest IE in the entire table?
    A: Helium (He).
71. Q71: Explain the jump in IE between IE₁ and IE₂ for an alkali metal (Na).
    A: Removing e⁻ from 3s¹ is easy (IE₁), but the second e⁻ must be removed from the stable noble gas core (2p⁶), causing a huge jump.
72. Q72: Why does N have a higher IE₁ than O?
    A: N has a half-filled stable p-orbital (2p³) configuration.
73. Q73: Why does Mg have a higher IE₁ than Al?
    A: Mg has a stable fully-filled s-orbital (3s²) configuration.
74. Q74: What unit is commonly used for ionization enthalpy?
    A: kJ mol⁻¹ or eV/atom.
75. Q75: Give one factor that influences IE besides size and Z_eff.
    A: Stability of completely filled or half-filled subshells.
Part 5: Electron Gain Enthalpy (EGE) and Electronegativity
76. Q76: Define Electron Gain Enthalpy (Δ_egH).
    A: The enthalpy change when an electron is added to an isolated gaseous atom to form a gaseous anion.
77. Q77: If Δ_egH is negative, what does that mean?
    A: The process is Exothermic (energy is released).
78. Q78: Which group has the largest negative Δ_egH (most electron affinity)?
    A: Group 17 (Halogens, they need only one electron to complete the octet).
79. Q79: Which element has the most negative Δ_egH in the entire table?
    A: Chlorine (Cl).
80. Q80: Why is the Δ_egH of F less negative than that of Cl?
    A: Due to the small size of the F atom, the incoming electron experiences strong inter-electronic repulsion in the compact 2p subshell.
81. Q81: Why do noble gases have large positive Δ_egH values?
    A: The incoming electron must enter the next higher energy shell, which is highly unfavorable.
82. Q82: How does Δ_egH generally vary across a Period (L → R)?
    A: Becomes more negative (increases).
83. Q83: How does Δ_egH generally vary down a Group (T → B)?
    A: Becomes less negative (decreases).
84. Q84: Define Electronegativity (EN).
    A: The tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond.
85. Q85: Is Electronegativity a measurable quantity like IE?
    A: No, it is a relative value (a scale).
86. Q86: How does Electronegativity vary across a Period (L → R)?
    A: Increases.
87. Q87: How does Electronegativity vary down a Group (T → B)?
    A: Decreases.
88. Q88: Which element has the highest electronegativity?
    A: Fluorine (F) (EN=4.0 on Pauling Scale).
89. Q89: Which elements have the lowest electronegativity?
    A: Alkali Metals (like Cs, Fr).
90. Q90: What scale is most commonly used to measure EN?
    A: The Pauling Scale.
91. Q91: How does EN relate to metallic character?
    A: Lower EN implies higher metallic character.
92. Q92: How does EN relate to acidic character of oxides?
    A: Higher EN (of the element) implies more acidic oxide.
93. Q93: Give the relationship between EN and bond polarity.
    A: Greater the EN difference, the more polar the bond.
94. Q94: What type of oxide is formed by highly electronegative elements?
    A: Acidic oxides (e.g., Cl₂O₇).
95. Q95: What type of oxide is formed by highly electropositive elements?
    A: Basic oxides (e.g., Na₂O).
Part 6: Chemical Periodicity and Other Properties
96. Q96: Define Valency (or Oxidation State for main group elements).
    A: The combining capacity of an element (equal to the number of valence e⁻ or 8−valence e⁻).
97. Q97: How does valency change across a period for main group elements?
    A: Increases from 1 to 4 and then decreases to 0.
98. Q98: Why do elements in Group 1 have a valency of 1?
    A: They have one valence electron (ns¹) which they easily lose.
99. Q99: How does Metallic Character vary across a Period (L → R)?
    A: Decreases.
100. Q100: How does Metallic Character vary down a Group (T → B)?
     A: Increases.
101. Q101: What are elements that show properties of both metals and non-metals called?
     A: Metalloids or Semimetals (e.g., Si, As, Ge).
102. Q102: What is the nature of Al₂O₃ (Aluminum Oxide)?
     A: Amphoteric (reacts with both acids and bases).
103. Q103: What is the diagonal relationship?
     A: Similarity in properties between an element and its diagonally opposite element in the next group and period (e.g., Li and Mg).
104. Q104: What is the cause of the diagonal relationship?
     A: The elements have nearly the same ionic size and charge-to-radius ratio.
105. Q105: What causes the Lanthanoid Contraction?
     A: Poor shielding of the 4f electrons.
106. Q106: What is a significant consequence of the Lanthanoid Contraction?
     A: The atomic radii of elements in the 5d series are very similar to those in the 4d series (e.g., Zr and Hf).
107. Q107: What is the chemical nature of SO₂?
     A: Acidic oxide (S is a non-metal).
108. Q108: What is the chemical nature of CaO?
     A: Basic oxide (Ca is an alkaline earth metal).
109. Q109: What is the maximum valency of an element in the 3rd period?
     A: 7 (e.g., Cl in HClO₄).
110. Q110: Which element marks the end of the 2nd period?
     A: Neon (Ne, Z=10).
111. Q111: Why are the elements of Group 17 called Halogens?
     A: Because they are salt-forming elements (from Greek halo = salt, gen = born).
112. Q112: What is the term for the stability gained by N, P, As etc., with their p³ configuration?
     A: Half-filled orbital stability.

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