📖 Chapter CN03 · NDA Class 11–12 Level🎯 NDA Level : High Priority
The Periodic Table is one of the most tested chapters in NDA Chemistry. Questions appear on periodic trends (atomic radius, ionisation energy, electronegativity), block identification, and properties of specific elements. An average student who memorises the four trend directions and can identify blocks from electronic configurations can score consistently from this chapter.
📌 What to expect in NDA (based on 2022–2025 pattern): (1) Modern periodic law and its statement vs Mendeleev's law; (2) Period and group structure — number of periods, groups, and elements in each; (3) s, p, d, f block identification — which elements belong where; (4) Trends across a period: atomic radius ↓, IE ↑, EA ↑, EN ↑; (5) Trends down a group: atomic radius ↑, IE ↓, EA ↓, EN ↓; (6) Specific exceptions (noble gases, electron affinity of N and O; IE₂ of Na vs Mg).
Topics at a Glance
① Periodic Table Structure
Modern law, periods, groups, element count
② s, p, d, f Blocks
Configuration basis, properties, examples
③ Atomic Radius & IE
Period/group trends, exceptions, Zeff
④ EA & Electronegativity
Trends, Pauling scale, fluorine vs chlorine
1. The Modern Periodic Table
1.1
Modern Periodic Law & Table Structure
Moseley's law replaced Mendeleev — atomic number is the true basis
📋 Mendeleev's Periodic Law (1869)
Properties of elements are a periodic function of their atomic masses
Arranged elements in increasing order of atomic mass
Left gaps for undiscovered elements (predicted Ga, Ge, Sc)
Flaw: Cobalt (58.9) placed before Nickel (58.7) — violated mass order
Could not explain position of isotopes or hydrogen
⚡ Modern Periodic Law (Moseley, 1913)
Properties of elements are a periodic function of their atomic numbers (Z)
Atomic number (not mass) determines element's properties
Periods (horizontal rows): 7 periods total
Period 1: 2 elements (H, He)
Period 2: 8 elements (Li → Ne)
Period 3: 8 elements (Na → Ar)
Period 4: 18 elements (K → Kr) ← first time d-block appears
Period 5: 18 elements (Rb → Xe)
Period 6: 32 elements (Cs → Rn) ← includes f-block (lanthanides)
Period 7: 32 elements (Fr → Og) ← includes f-block (actinides)
Groups (vertical columns): 18 groups total
Groups 1–2: s-block (alkali & alkaline earth metals)
Groups 3–12: d-block (transition metals)
Groups 13–18: p-block (includes non-metals, metalloids, noble gases)
Separate rows (below): f-block (lanthanides Z=57–71; actinides Z=89–103)
Total elements currently known: 118
Period number = number of electron shells. Group number (for s/p blocks) = number of valence electrons (e.g. Group 1 → 1 valence e⁻; Group 17 → 7 valence e⁻).
Fig. 1 — Schematic layout of the Modern Periodic Table showing s, p, d, f blocks by colour. Period numbers (P1–P7) on left; Group numbers across top.
📝 TOPIC-WISE PYQ
Periodic Table Structure — NDA Pattern Questions
Q1. The modern periodic law states that properties of elements are a periodic function of their:
(a) Atomic mass (b) Atomic number (c) Number of neutrons (d) Valence electrons
Answer: (b) Atomic number
Moseley (1913) showed that X-ray frequencies of elements increase regularly with atomic number, not atomic mass. The modern law — properties repeat periodically with increasing atomic number — resolved all anomalies (Co–Ni reversal, isotopes). The repetition occurs because electronic configuration repeats with increasing Z.
Q2. How many elements are there in Period 4 of the Modern Periodic Table?
(a) 8 (b) 18 (c) 32 (d) 2
Answer: (b) 18
Period 4 (K to Kr, Z=19 to 36) contains 18 elements — it is the first period to include the d-block (10 transition elements from Sc to Zn, Z=21–30) alongside 2 s-block (K, Ca) and 6 p-block (Ga–Kr) elements. Periods 1, 2, 3 have 2, 8, 8 elements respectively.
Q3. An element has electronic configuration 2, 8, 7. To which group and period does it belong?
(a) Group 7, Period 3 (b) Group 17, Period 3 (c) Group 7, Period 2 (d) Group 16, Period 3
Answer: (b) Group 17, Period 3
Total electrons = 2+8+7 = 17 → Z=17 = Chlorine. Number of shells = 3 → Period 3. Valence electrons = 7 → Group 17 (halogens). Chlorine is a classic NDA element — highly reactive non-metal, forms HCl, NaCl. Group 17 elements all have 7 valence electrons and high electronegativity.
2. s, p, d, f Block Elements
2.1
Block Classification Based on Electronic Configuration
The last orbital where valence electron enters determines the block
The four blocks arise from which type of subshell (s, p, d, or f) receives the last electron during Aufbau filling. Each block has a characteristic set of properties.
① s-Block — Groups 1 & 2 (Alkali and Alkaline Earth Metals)
Configuration: Outermost electron enters an s subshell. Group 1: ns¹; Group 2: ns². Elements: H, He, Li, Na, K, Rb, Cs, Fr (Gr.1) and Be, Mg, Ca, Sr, Ba, Ra (Gr.2). Properties: Highly reactive metals (except H, He); low ionisation energy; large atomic radius; form ionic compounds; strong reducing agents. Alkali metals (Gr.1) react vigorously with water. He (noble gas) is placed here due to ns² config but behaves like Group 18.
② p-Block — Groups 13–18 (Non-metals, Metalloids, Noble Gases)
Configuration: Last electron enters a p subshell. ns²np¹ to ns²np⁶. Elements: B, C, N, O, F, Ne and their heavier congeners. Includes all non-metals (except H), all metalloids (B, Si, Ge, As, Sb, Te), all halogens (F, Cl, Br, I, At), and all noble gases (He, Ne, Ar, Kr, Xe, Rn). Properties: Most diverse block; high electronegativity on right side; noble gases are chemically inert (complete octet). Halogens most reactive non-metals.
③ d-Block — Groups 3–12 (Transition Metals)
Configuration: Last electron enters a d subshell: (n−1)d¹⁻¹⁰ ns⁰⁻². Elements: Sc to Zn (Period 4), Y to Cd (Period 5), La/Hf to Hg (Period 6), Ac/Rf to Cn (Period 7). Properties: Hard, high melting points, good conductors; variable oxidation states (e.g. Fe: +2/+3; Mn: +2 to +7); form coloured compounds; act as catalysts (Fe in Haber process, Pt in catalytic converters); paramagnetic due to unpaired d-electrons. Zn, Cd, Hg sometimes excluded (d¹⁰ complete).
Configuration: Last electron enters an f subshell: (n−2)f¹⁻¹⁴. Lanthanides (Z=57–71): La to Lu — Period 6; 4f subshell fills. Actinides (Z=89–103): Ac to Lr — Period 7; 5f subshell fills. Properties: Lanthanides show very similar properties ("lanthanide contraction"); actinides are all radioactive. Used in magnets, nuclear reactors, defence technology (U, Pu in nuclear weapons/reactors).
Block
Groups
Filling subshell
Key examples
Characteristic property
s
1, 2
ns¹ or ns²
Na, K, Ca, Mg
Highly reactive metals, low IE
p
13–18
np¹ to np⁶
C, N, O, Cl, Ne
Non-metals, metalloids, noble gases
d
3–12
(n−1)d¹⁻¹⁰
Fe, Cu, Zn, Cr, Mn
Variable oxidation states, coloured ions
f
— (separate)
(n−2)f¹⁻¹⁴
La, Ce, U, Pu
Lanthanide contraction; radioactive actinides
📌 NDA Shortcut — Block Identification from Configuration:
If last filled subshell is s → s-block; p → p-block; d → d-block; f → f-block.
Exception: He (1s²) is in s-block but placed in Group 18 (noble gas behaviour). Zn ([Ar]3d¹⁰4s²) — last filled subshell is 4s, but Zn is conventionally placed in d-block.
📝 TOPIC-WISE PYQ
s, p, d, f Blocks — NDA Pattern Questions
Q1. An element with electronic configuration [Ar] 3d⁵ 4s¹ belongs to which block?
(a) s-block (b) p-block (c) d-block (d) f-block
Answer: (c) d-block
This is Chromium (Z=24) — the last subshell being filled is the 3d (an exception due to half-filled d stability). d-block elements have (n−1)d partially filled. Cr is in Group 6, Period 4. Its oxide Cr₂O₃ is used in green pigments and it shows variable oxidation states (+2, +3, +6).
Q2. Which of the following is a property of transition metals (d-block)?
(a) They form only colourless compounds (b) They have only one fixed oxidation state (c) They show variable oxidation states and form coloured compounds (d) They are all non-metals
Answer: (c) Variable oxidation states and coloured compounds
d-block (transition) metals have partially filled d-orbitals. Different numbers of d-electrons can be lost → variable oxidation states. Example: Mn shows +2, +3, +4, +6, +7 states. Coloured compounds arise because electrons in partially filled d-orbitals absorb visible light and transition between d-energy levels (d-d transition). CuSO₄ is blue, K₂Cr₂O₇ is orange, KMnO₄ is purple.
Q3. Lanthanides and actinides belong to which block?
(a) s-block (b) p-block (c) d-block (d) f-block
Answer: (d) f-block
Lanthanides (Z=57–71) fill the 4f subshell. Actinides (Z=89–103) fill the 5f subshell. Both are inner transition metals placed in the f-block. All actinides are radioactive — Uranium (Z=92) and Plutonium (Z=94) are used in nuclear reactors and weapons. NDA sometimes asks which actinides are used in nuclear technology.
🧠 TRICKY QUESTIONS
Block Classification — Conceptual Traps
Q. Helium (1s²) has a completely filled s-subshell — so is it in the s-block or the p-block?
Answer: s-block (by configuration) but placed in Group 18 (p-block side) by behaviour.
He has configuration 1s² — last filled orbital is 1s, so it is technically an s-block element. However, He is placed in Group 18 alongside the noble gases (which are p-block, having ns²np⁶) because it shares identical chemical inertness and a completely filled valence shell. NDA question wording matters: "Which block does He belong to?" → s-block. "Which group is He in?" → Group 18.
Q. Zinc (Z=30) has configuration [Ar] 3d¹⁰ 4s² — its 3d is completely filled. Is Zn a transition metal?
Answer: Debated — Zn is often excluded from "true" transition metals.
The IUPAC definition of a transition metal requires an atom with an incomplete d subshell (in ground state or as an ion). Zn in ground state has 3d¹⁰ (complete) and Zn²⁺ also has 3d¹⁰ (complete). So Zn does not qualify strictly. It also forms colourless compounds (ZnSO₄ is white) and has only +2 oxidation state. However, conventionally Zn is placed in the d-block in the periodic table. NDA accepts Zn as part of d-block but acknowledges it lacks "transition metal character."
3. Periodicity in Properties
📈 Trend Summary — All Four Properties at a Glance
Property
Across a Period (→)
Down a Group (↓)
Atomic Radius
↓ Decreases — Z increases, same shells, stronger pull
↓ Decreases — larger atom, more shielding, easier to remove e⁻
Electron Affinity (EA)
↑ Increases — more electronegative, more eager to gain e⁻
↓ Decreases — larger radius, added e⁻ farther from nucleus
Electronegativity (EN)
↑ Increases — highest at F (top-right)
↓ Decreases — lowest at Cs/Fr (bottom-left)
3.1
Atomic Radius
Measure of atom size — covalent, metallic, or van der Waals radius
Atomic radius is the distance from the nucleus to the outermost electrons. It cannot be measured directly — defined as half the internuclear distance between two identical bonded atoms (covalent radius for non-metals; metallic radius for metals).
Fig. 2 — Atomic radius decreases across Period 3 (Na→Ar) as nuclear charge increases but electrons are added to the same shell, increasing Zeff.
📋 Why radius decreases across a period
Moving left to right: Z (protons) increases
Electrons added to the same shell (shielding stays similar)
Effective nuclear charge Zeff = Z − shielding increases
Greater Zeff → stronger pull on electrons → shell contracts
Inner shells shield outer electrons more effectively
Zeff increases only slightly; shielding increases greatly
Result: Li (152 pm) < Na (186 pm) < K (227 pm) < Rb (248 pm)
📌 Exception: Noble Gases — Noble gas radii are van der Waals radii (non-bonding), which are larger than covalent/metallic radii. So Ar appears to have a larger radius than Cl if van der Waals radii are compared — but this is not a fair comparison. If the same type of radius is compared, the trend holds.
3.2
Ionisation Energy (IE)
Energy needed to remove an electron from a gaseous atom — most NDA-tested trend
Ionisation energy (IE₁) is the minimum energy required to remove the most loosely bound electron from a gaseous atom in its ground state: X(g) + Energy → X⁺(g) + e⁻. It is measured in kJ/mol or eV.
⚛ Ionisation Energy — Key Facts & Values
Successive IE: IE₁ < IE₂ < IE₃ … (always increases — each next e⁻ harder to remove)
Huge jump in IE signals core electron removal (used to find group):
Na: IE₁=496, IE₂=4562 kJ/mol → HUGE jump at 2nd → Group 1 (1 valence e⁻) ✓
Mg: IE₁=738, IE₂=1451, IE₃=7733 → HUGE jump at 3rd → Group 2 ✓
Trend across Period 3 (IE₁, kJ/mol):
Na(496) < Mg(738) < Al(577)* < Si(786) < P(1012) < S(999)* < Cl(1251) < Ar(1521)
* Exceptions: Al < Mg ; S < P (NDA favourite traps — explained below)
Trend down Group 1 (IE₁, kJ/mol):
Li(520) > Na(496) > K(419) > Rb(403) > Cs(376) — decreases down group ✓
Unit: kJ/mol. Larger IE₁ = atom holds electrons more tightly (e.g. noble gases have highest IE). Smallest IE₁ = Cs (most easily ionised) among natural elements.
🚨 Exception 1: IE(Al) < IE(Mg) — Why?
Mg: [Ne] 3s² — the 3s subshell is fully filled (stable)
Al: [Ne] 3s² 3p¹ — the extra e⁻ is in a higher energy 3p orbital
The 3p electron of Al is easier to remove than Mg's 3s²
Fully filled subshell = extra stability → higher IE for Mg
S: [Ne] 3s² 3p⁴ — 3p has one paired orbital (↑↓ | ↑ | ↑)
Paired electrons repel each other → easier to remove one
Half-filled subshell = extra stability → higher IE for P
So IE(S) < IE(P) despite S having higher Z
📝 TOPIC-WISE PYQ
Ionisation Energy — NDA Pattern Questions
Q1. The first ionisation energy of nitrogen (N) is greater than that of oxygen (O). This is because:
(a) N has smaller atomic radius (b) N has a half-filled 2p subshell which is extra stable (c) O has more electrons (d) N is a lighter element
Answer: (b) N has a half-filled 2p subshell which is extra stable
N: [He] 2s² 2p³ — half-filled 2p (↑|↑|↑), extra stable. O: [He] 2s² 2p⁴ — one pair in 2p, electron–electron repulsion makes it easier to remove one. Therefore IE₁(N) = 1402 kJ/mol > IE₁(O) = 1314 kJ/mol. This is the same principle as the S < P exception.
Q2. The successive ionisation energies of an element are 578, 1817, 2745, 11578 kJ/mol. To which group does this element belong?
(a) Group 1 (b) Group 2 (c) Group 3 (d) Group 13
Answer: (c) Group 3
The huge jump occurs between IE₃ and IE₄ (2745 → 11578 kJ/mol) — the 4th electron is a core electron. This means the element has 3 valence electrons → Group 3 (or Group 13 in modern notation). The jump indicates removing from a filled inner shell. This is Aluminium (Al) — IE₁=578, huge jump at IE₄.
Q3. Which of the following has the highest first ionisation energy?
(a) Na (b) Mg (c) Al (d) Si
Answer: (b) Mg
General trend: IE increases across period. But Al (Group 13) has a lower IE than Mg (Group 2) because Al loses a 3p electron while Mg has a stable filled 3s². Among Na, Mg, Al, Si: Na(496) < Al(577) < Mg(738) < Si(786). So Si has the highest here, but if the option "Si" appears, it's correct. In this specific set: Si has highest — but if the question lists only Na, Mg, Al then Mg wins due to the Al exception.
3.3
Electron Affinity (EA) & Electronegativity (EN)
EA = energy when atom gains e⁻; EN = power to attract shared e⁻ in a bond
⚡ Electron Affinity (EA)
Energy released when a gaseous atom gains one electron: X + e⁻ → X⁻
More negative EA = greater tendency to accept electron
Across period: EA generally increases (more electronegative)
Down group: EA generally decreases
Highest EA: Cl (not F!) — F's small size causes electron–electron repulsion when new e⁻ added in cramped 2p orbital
N and noble gases: nearly zero/positive EA (stable half-filled or full shells)
📈 Electronegativity (EN) — Pauling Scale
Ability of bonded atom to attract shared electrons toward itself
Pauling scale: F = 4.0 (highest); Cs = 0.7 (lowest)
Across period: EN increases → F most electronegative
Down group: EN decreases → Cs/Fr least electronegative
Noble gases: not assigned EN (no bonds)
EN determines bond polarity: large difference → ionic; small → covalent
Fig. 3 — Pauling electronegativity values for Period 2 and key Period 3 elements, showing steady increase across the period. Fluorine (4.0) is the absolute maximum.
📌 NDA Critical Exceptions — Electron Affinity: (1) F vs Cl: Cl has higher EA than F despite F being more electronegative. F's 2p orbital is too small and compact — adding a new electron causes severe electron–electron repulsion. Cl has larger 3p orbitals that accommodate the new electron more comfortably. (2) N (Group 15): N has nearly zero EA — the half-filled 2p³ (extra stable) strongly resists gaining another electron. Same for P in Group 15. (3) Noble gases: EA ≈ 0 or slightly positive — complete octet makes them unwilling to gain electrons.
📝 TOPIC-WISE PYQ
Electron Affinity & Electronegativity — NDA Pattern Questions
Q1. Which element has the highest electronegativity on the Pauling scale?
(a) Oxygen (b) Chlorine (c) Fluorine (d) Nitrogen
Answer: (c) Fluorine
Fluorine (EN = 4.0) is the most electronegative element on the Pauling scale. Being in Period 2 (small size) and Group 17 (7 valence electrons, needs only 1 more), it has the maximum pull on bonding electrons. Oxygen is second (EN = 3.5). EN increases towards the top-right corner of the periodic table (excluding noble gases).
Q2. The electron affinity of fluorine (F) is less than that of chlorine (Cl). What is the reason?
(a) F has higher electronegativity (b) F has a smaller 2p orbital causing electron–electron repulsion when gaining an e⁻ (c) Cl has higher nuclear charge (d) F is in Period 2 which always has low EA
Answer: (b) F has a smaller 2p orbital causing electron–electron repulsion
F's 2p orbital is very compact. When an incoming electron is added, the existing 7 electrons in the small shell cause significant repulsion — reducing the energy released (EA). Cl has a larger 3p orbital, so the new electron is accommodated more easily, releasing more energy. EA(Cl) ≈ −349 kJ/mol; EA(F) ≈ −328 kJ/mol. This is a classic NDA exception.
Q3. Consider the following statements about electronegativity:
I. It increases across a period from left to right.
II. It increases down a group.
III. Noble gases are assigned the highest electronegativity values.
(a) I only (b) I and II (c) II and III (d) All three
Answer: (a) I only
Statement I is correct — EN increases across a period as Zeff increases.
Statement II is wrong — EN decreases down a group (larger atoms have weaker pull).
Statement III is wrong — Noble gases are generally not assigned electronegativity values since they don't typically form bonds.
🧠 TRICKY QUESTIONS
Periodic Trends — The Classic NDA Traps
Q. Arrange Na, Mg, Al, Si in order of increasing first ionisation energy. Most students get this wrong.
Correct order: Na < Al < Mg < Si
Simple expectation (increasing Z): Na < Mg < Al < Si. But Al has lower IE than Mg (Al loses a 3p electron; Mg has stable filled 3s²).
Actual IE₁ (kJ/mol): Na(496) < Al(577) < Mg(738) < Si(786). Trap: Students assume it simply increases left to right, missing the Mg > Al exception. Similarly in Period 2: Li < B < Be < C < O < N < F < Ne — both Be > B and N > O are exceptions!
Q. Atomic radius of Ga (Z=31) is slightly less than Al (Z=13) even though Ga is in the next period below Al. Why?
Answer: d-block contraction (poor shielding by d-electrons)
Between Al and Ga, the 3d subshell (10 electrons) fills in Period 4. d-electrons are poor shielders of nuclear charge — they don't effectively shield the outer 4p electrons from the increasing nuclear charge. This causes a greater-than-expected increase in Zeff for Ga, pulling its outer electrons closer. This is analogous to the "lanthanide contraction" seen for elements after the f-block. NDA 2025 tested a similar concept — the general trend is "radius increases down a group" but the d-block contraction creates a notable exception for Ga, In.
Q. Noble gases have zero electron affinity (or slightly positive). Does this mean they are the most stable group in the table?
Answer: Yes — but for a specific reason (complete valence shell, not high EN).
Noble gases have ns²np⁶ (complete octet, except He with 1s²). They don't need to gain OR lose electrons. Their IE is the highest in each period (very hard to remove electrons) AND their EA is near zero (no tendency to gain). This double stability makes them chemically inert. The common trap: students confuse "high EN" (which is about bond polarity and pulling shared electrons) with "chemical stability" of noble gases. Noble gases are stable due to complete shells — not because of high electronegativity.
📄 CN03 Formula & Fact Sheet — Quick Reference
🔮 Periodic Table Structure
7 periods; 18 groups; 118 elements known
P1=2, P2=8, P3=8, P4=18, P5=18, P6=32, P7=32 elements
Period no. = number of electron shells
Group no. = valence electrons (s/p block)
Modern law: properties ∝ atomic number Z (Moseley)
⚛ Block Summary
s-block: Gr.1,2 — ns¹ or ns² (alkali/alkaline earth metals)
p-block: Gr.13–18 — np¹ to np⁶ (non-metals, metalloids, noble gases)
d-block: Gr.3–12 — (n−1)d¹⁻¹⁰ (transition metals)
f-block: lanthanides (4f) and actinides (5f)
He: s-block by config; placed in Group 18 by behaviour
📈 Atomic Radius Trend
Across period →: decreases (Zeff increases, same shell)
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