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

 Over 100 Short Q&A for Periodic Classification of Elements (Class 11 NCERT)

Part 1: Genesis of Classification & Modern Periodic Table

Q. No.	Question	Answer
1	Who first classified elements into triads?	Johann Dobereiner.
2	State Dobereiner's Law of Triads.	The atomic mass of the middle element is approximately the arithmetic mean of the other two.
3	Who proposed the Law of Octaves?	John Newlands.
4	What was the basis of Newlands' classification?	Increasing atomic masses (properties repeated every eighth element).
5	What was the major criterion for Mendeleev's classification?	Atomic weight (mass) and similarity in chemical properties (oxide and hydride formulas).
6	State Mendeleev's Periodic Law.	The properties of elements are a periodic function of their atomic weights.
7	Name an element whose existence was predicted by Mendeleev as 'Eka-silicon'.	Germanium (Ge).
8	What was the main drawback of Mendeleev's Periodic Table regarding position?	The position of isotopes and anomalous pairs (e.g., Ar before K).
9	Who proposed the Modern Periodic Law?	Henry Moseley.
10	What is the basis of the Modern Periodic Table?	Atomic Number (Z).
11	State the Modern Periodic Law.	The properties of elements are a periodic function of their atomic numbers.
12	How many groups and periods are in the Modern Periodic Table?	18 Groups (vertical columns) and 7 Periods (horizontal rows).
13	What does the Period number correspond to?	The Principal Quantum Number (n) of the outermost shell.
14	What is the common feature for elements in the same group?	Same number of valence electrons (similar outermost electronic configuration).
15	Which period is the shortest?	Period 1 (contains only 2 elements, H and He).
16	How many elements are in the 6th period?	32 elements (6s, 4f, 5d, 6p).
17	Which element is the starting point of the 4th period?	Potassium (K,Z=19).
18	What is the group number for noble gases?	Group 18 (or 0 group).
19	What is the group number for alkali metals?	Group 1.
20	What is the group number for halogens?	Group 17.

Part 2: Classification by Blocks (s, p, d, f)

Q. No.	Question	Answer
21	What determines the block an element belongs to?	The orbital that receives the last differentiating electron.
22	What is the general electronic configuration of s-block elements?	ns1−2.
23	Name the groups that constitute the s-block.	Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals).
24	What is the general electronic configuration of p-block elements?	ns2np1−6.
25	Which group is the only non-metal in the s-block?	Hydrogen (H).
26	Which block contains maximum number of metals, non-metals, and metalloids?	p-block.
27	What is the general electronic configuration of d-block elements?	(n−1)d1−10ns0−2.
28	What is another name for d-block elements?	Transition Elements.
29	Name the groups that constitute the d-block.	Group 3 to Group 12.
30	What is the general electronic configuration of f-block elements?	(n−2)f1−14(n−1)d0−1ns2.
31	What is another name for f-block elements?	Inner-Transition Elements.
32	Name the two series of f-block elements.	Lanthanoids (Lanthanides) and Actinoids (Actinides).
33	What is the maximum number of electrons an f-subshell can accommodate?	14 electrons.
34	To which block does the element with Z=30 (Zn) belong?	d-block (Group 12).
35	To which block does the element with Z=54 (Xe) belong?	p-block (Group 18).
36	What is the group and period of the element with Z=17 (Cl)?	Group 17, Period 3.
37	What is the common oxidation state of all Lanthanoids?	+3.
38	What are s-block elements generally known for (electro-)?	They are highly Electropositive.
39	What are p-block elements generally known for (electro-)?	They are often highly Electronegative.
40	What are the elements after Z=92 called?	Transuranium Elements (all are radioactive).

Part 3: Atomic and Ionic Radii

Q. No.	Question	Answer
41	Define Atomic Radius.	The distance from the center of the nucleus to the outermost shell of electrons.
42	Define Covalent Radius.	Half the distance between the nuclei of two identical atoms covalently bonded.
43	Define Metallic Radius.	Half the internuclear distance between the nuclei of two adjacent metal atoms in a metallic crystal.
44	Define van der Waals Radius.	Half the distance between the nuclei of two non-bonded isolated atoms.
45	How does Atomic Radius vary across a Period (L → R)?	Decreases.
46	Why does atomic radius decrease across a period?	Increase in Effective Nuclear Charge (Zeff​) (valence electrons remain in the same shell).
47	How does Atomic Radius vary down a Group (T → B)?	Increases.
48	Why does atomic radius increase down a group?	Increase in the number of electron shells (Principle Quantum Number, n).
49	Which is larger: a parent atom (A) or its Cation (A+)?	Parent atom (A) (Cation is smaller).
50	Why is a cation smaller than its parent atom?	Higher Zeff​ and loss of the outermost shell (sometimes).
51	Which is larger: a parent atom (A) or its Anion (A−)?	Anion (A−) (Anion is larger).
52	Why is an anion larger than its parent atom?	Increased electron-electron repulsion (decreases Zeff​ per electron).
53	Define Isoelectronic Species.	Atoms or ions having the same number of electrons.
54	Arrange O2−,F−,Na+,Mg2+ in order of increasing ionic size.	Mg2+<Na+<F−<O2− (size decreases as Z increases for isoelectronics).
55	What is the main factor determining the size of isoelectronic species?	Nuclear Charge (Z) (Higher Z, smaller size).
56	Which radius is generally the largest among the four types?	van der Waals radius.
57	What is Shielding Effect (or Screening Effect)?	The reduction in the nuclear attraction felt by the valence electrons due to the presence of inner-shell electrons.
58	Which orbital type provides the least shielding effect?	f-orbitals.
59	How does Zeff​ generally change down a group?	It remains nearly constant (increase in Z is offset by increased shielding).
60	How does Zeff​ generally change across a period (L → R)?	It increases (charge increases, shielding is almost constant).

Part 4: Ionization Enthalpy (IE)

Q. No.	Question	Answer
61	Define Ionization Enthalpy (Δi​H).	The minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
62	Is the first ionization enthalpy generally endothermic or exothermic?	Endothermic (energy is required, Δi​H>0).
63	Why is IE2​ always greater than IE1​?	It is harder to remove an electron from an already stable cation (higher Zeff​).
64	How does IE generally vary across a Period (L → R)?	Increases.
65	Why does IE increase across a period?	Increase in Zeff​ and decrease in atomic size.
66	How does IE generally vary down a Group (T → B)?	Decreases.
67	Why does IE decrease down a group?	Increase in atomic size and increased shielding.
68	Which elements have the highest ionization enthalpy?	Noble Gases (due to stable, fully-filled configuration).
69	Which Group has the lowest ionization enthalpy?	Group 1 (Alkali Metals).
70	Which element has the highest IE in the entire table?	Helium (He).
71	Explain the jump in IE between IE1​ and IE2​ for an alkali metal (Na).	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	Why does N have a higher IE1​ than O?	N has a half-filled stable p-orbital (2p3) configuration.
73	Why does Mg have a higher IE1​ than Al?	Mg has a stable fully-filled s-orbital (3s2) configuration.
74	What unit is commonly used for ionization enthalpy?	kJ mol−1 or eV/atom.
75	Give one factor that influences IE besides size and Zeff​.	Stability of completely filled or half-filled subshells.

Part 5: Electron Gain Enthalpy (EGE) and Electronegativity

Q. No.	Question	Answer
76	Define Electron Gain Enthalpy (Δeg​H).	The enthalpy change when an electron is added to an isolated gaseous atom to form a gaseous anion.
77	If Δeg​H is negative, what does that mean?	The process is Exothermic (energy is released).
78	Which group has the largest negative Δeg​H (most electron affinity)?	Group 17 (Halogens, they need only one electron to complete the octet).
79	Which element has the most negative Δeg​H in the entire table?	Chlorine (Cl).
80	Why is the Δeg​H of F less negative than that of Cl?	Due to the small size of the F atom, the incoming electron experiences strong inter-electronic repulsion in the compact 2p subshell.
81	Why do noble gases have large positive Δeg​H values?	The incoming electron must enter the next higher energy shell, which is highly unfavorable.
82	How does Δeg​H generally vary across a Period (L → R)?	Becomes more negative (increases).
83	How does Δeg​H generally vary down a Group (T → B)?	Becomes less negative (decreases).
84	Define Electronegativity (EN).	The tendency of an atom to attract the shared pair of electrons towards itself in a covalent bond.
85	Is Electronegativity a measurable quantity like IE?	No, it is a relative value (a scale).
86	How does Electronegativity vary across a Period (L → R)?	Increases.
87	How does Electronegativity vary down a Group (T → B)?	Decreases.
88	Which element has the highest electronegativity?	Fluorine (F) (EN=4.0 on Pauling Scale).
89	Which elements have the lowest electronegativity?	Alkali Metals (like Cs,Fr).
90	What scale is most commonly used to measure EN?	The Pauling Scale.
91	How does EN relate to metallic character?	Lower EN implies higher metallic character.
92	How does EN relate to acidic character of oxides?	Higher EN (of the element) implies more acidic oxide.
93	Give the relationship between EN and bond polarity.	Greater the EN difference, the more polar the bond.
94	What type of oxide is formed by highly electronegative elements?	Acidic oxides (e.g., Cl2​O7​).
95	What type of oxide is formed by highly electropositive elements?	Basic oxides (e.g., Na2​O).

Part 6: Chemical Periodicity and Other Properties

Q. No.	Question	Answer
96	Define Valency (or Oxidation State for main group elements).	The combining capacity of an element (equal to the number of valence e− or 8−valence e−).
97	How does valency change across a period for main group elements?	Increases from 1 to 4 and then decreases to 0.
98	Why do elements in Group 1 have a valency of 1?	They have one valence electron (ns1) which they easily lose.
99	How does Metallic Character vary across a Period (L → R)?	Decreases.
100	How does Metallic Character vary down a Group (T → B)?	Increases.
101	What are elements that show properties of both metals and non-metals called?	Metalloids or Semimetals (e.g., Si,As,Ge).
102	What is the nature of Al2​O3​ (Aluminum Oxide)?	Amphoteric (reacts with both acids and bases).
103	What is the diagonal relationship?	Similarity in properties between an element and its diagonally opposite element in the next group and period (e.g., Li and Mg).
104	What is the cause of the diagonal relationship?	The elements have nearly the same ionic size and charge-to-radius ratio.
105	What causes the Lanthanoid Contraction?	Poor shielding of the 4f electrons.
106	What is a significant consequence of the Lanthanoid Contraction?	The atomic radii of elements in the 5d series are very similar to those in the 4d series (e.g., Zr and Hf).
107	What is the chemical nature of SO2​?	Acidic oxide (S is a non-metal).
108	What is the chemical nature of CaO?	Basic oxide (Ca is an alkaline earth metal).
109	What is the maximum valency of an element in the 3rd period?	7 (e.g., Cl in HClO4​).
110	Which element marks the end of the 2nd period?	Neon (Ne,Z=10).
111	Why are the elements of Group 17 called Halogens?	Because they are salt-forming elements (from Greek halo = salt, gen = born).
112	What is the term for the stability gained by N,P,As etc., with their p3 configuration?	Half-filled orbital stability.