Periodic Classification of Elements

Correct Option: (c) The atoms lose their electrons more easily.

Left to Right Trends (Period 2: Li → Ne):

• (a) Correct: Metallic character decreases (Li, Be metals → B metalloid → C,N,O,F,Ne non-metals)
• (b) Correct: Valence electrons increase (Li:1, Be:2, B:3, C:4, N:5, O:6, F:7, Ne:8)
• (c) Incorrect: Ionization energy increases, so atoms lose electrons less easily moving right
• (d) Correct: Oxides change from basic (Li₂O) → amphoteric (BeO) → acidic (B₂O₃, CO₂, N₂O₅, etc.)

Key concept: Nuclear charge increases → stronger attraction on electrons → harder to remove electrons.

Correct Option: (b) Mg

Analysis:
• Formula XCl₂ indicates valency 2 (X²⁺ ion)
• Solid with high melting point suggests ionic compound
• Group 2 elements form MCl₂ (MgCl₂, CaCl₂, etc.)

Group comparison:
• Group 1 (Na): Forms NaCl (valency 1)
• Group 2 (Mg): Forms MgCl₂ (valency 2) ✓
• Group 13 (Al): Forms AlCl₃ (valency 3)
• Group 14 (Si): Forms SiCl₄ (covalent, low MP)

Likely elements: Mg, Ca, Sr, Ba (alkaline earth metals, Group 2)

(a) Neon (Ne) - Atomic number 10: K=2, L=8 (both full)
Also: Helium (He) - 1 shell with 2 electrons (full)

(b) Magnesium (Mg) - Atomic number 12: K=2, L=8, M=2
Group 2, Period 3, Valency 2

(c) Silicon (Si) - Atomic number 14: K=2, L=8, M=4
Group 14, Period 3, Valency 4

(d) Boron (B) - Atomic number 5: K=2, L=3
Group 13, Period 2, Valency 3

(e) Carbon (C) - Atomic number 6: K=2, L=4
2nd shell (4 electrons) = 2 × 1st shell (2 electrons)
Also: Any element with 6 electrons total

(a) Boron Group (Group 13): B, Al, Ga, In, Tl
• Valence electrons: 3 electrons in outermost shell (ns²np¹)
• Valency: 3 (form M³⁺ ions or share 3 electrons)
• Nature: Metalloids to poor metals
• Oxides: Amphoteric (B₂O₃, Al₂O₃)
• Example compounds: BCl₃, AlCl₃, GaAs

(b) Fluorine Group (Group 17 - Halogens): F, Cl, Br, I, At
• Valence electrons: 7 electrons in outermost shell (ns²np⁵)
• Valency: 1 (gain 1 electron to form X⁻ ions)
• Nature: Highly reactive non-metals
• Physical state: F₂, Cl₂ (gases), Br₂ (liquid), I₂ (solid)
• Oxidizing agents: Strong tendency to gain electrons
• Example compounds: NaCl, HCl, CHCl₃

(a) Atomic number: 17
• Calculation: 2 + 8 + 7 = 17 electrons
• Neutral atom: Number of electrons = Atomic number
• Element: Chlorine (Cl)

(b) Chemically similar to: Fluorine (F), Atomic number 9

Reason:
• Electronic configuration 2,8,7 indicates 7 valence electrons
• Elements with same valence electrons belong to same group (Group 17 - Halogens)
• F (2,7) and Cl (2,8,7) both have 7 valence electrons
• Similar chemical properties: Form -1 ions, diatomic molecules, strong oxidizing agents

Others:
• N (2,5): 5 valence electrons (Group 15)
• P (2,8,5): 5 valence electrons (Group 15)
• Ar (2,8,8): 8 valence electrons (Group 18 - Noble gas)

Analysis of positions:
• A: Group 16, 2nd period → Oxygen (O)
• B: Group 16, 3rd period → Sulphur (S)
• C: Group 17, 3rd period → Chlorine (Cl)

(a) A (Oxygen): Non-metal
Group 16 contains oxygen family (O, S, Se, Te, Po) - all non-metals except Po

(b) Reactivity comparison: C (Cl) is more reactive than A (O)
• In non-metals (Groups 16-17), reactivity increases upward in group
• Cl (Group 17) is more reactive than O (Group 16) in same period
• Halogens are most reactive non-metals

(c) Size comparison: C (Cl) is smaller than B (S)
• Both in Period 3
• Atomic size decreases left to right in period
• Order of size: Na > Mg > Al > Si > P > S > Cl
• S (atomic radius 104 pm) > Cl (atomic radius 99 pm)

(d) Ion formed by A (Oxygen): Anion (O²⁻)
• Group 16 elements gain 2 electrons to achieve octet
• Form divalent anions: O²⁻, S²⁻, etc.
• Oxide ion (O²⁻) is common in ionic compounds (MgO, CaO, etc.)

Electronic configurations:

Nitrogen (N, Z=7): 2, 5
• K shell: 2 electrons
• L shell: 5 electrons (valence shell)
• Period 2, Group 15

Phosphorus (P, Z=15): 2, 8, 5
• K shell: 2 electrons
• L shell: 8 electrons
• M shell: 5 electrons (valence shell)
• Period 3, Group 15

More electronegative: Nitrogen (N)

Reason:
• Electronegativity decreases down a group
• Nitrogen (Period 2) is above phosphorus (Period 3) in Group 15
• Nitrogen has smaller atomic size → stronger nuclear attraction on bonding electrons
• Electronegativity values: N (3.04) > P (2.19) on Pauling scale
• General trend: Electronegativity ∝ 1/Atomic size

Electronic configuration determines two key aspects of an element's position:

1. Period Number:
• Equal to the number of electron shells
• Example: Na (2,8,1) → 3 shells → Period 3
• Example: Cl (2,8,7) → 3 shells → Period 3

2. Group Number:
• For main group elements (s and p block):
  Group number = Number of valence electrons
• Example: Mg (2,8,2) → 2 valence electrons → Group 2
• Example: Cl (2,8,7) → 7 valence electrons → Group 17
• For noble gases: 8 valence electrons (except He:2) → Group 18

3. Block Identification:
• Last electron enters s orbital → s-block (Groups 1-2)
• Last electron enters p orbital → p-block (Groups 13-18)
• Last electron enters d orbital → d-block (Transition metals)

Example correlation:
• Element with config 2,8,4 → 3 shells (Period 3), 4 valence electrons (Group 14) → Silicon (Si)

Elements and their positions:

• Calcium (Ca, Z=20): Group 2, Period 4
• Z=12: Magnesium (Mg) - Group 2, Period 3
• Z=19: Potassium (K) - Group 1, Period 4
• Z=21: Scandium (Sc) - Group 3, Period 4 (Transition metal)
• Z=38: Strontium (Sr) - Group 2, Period 5

Elements resembling calcium: Z=12 (Mg) and Z=38 (Sr)

Reason:
• Similar properties occur within same group (vertical column)
• Calcium (Group 2): Alkaline earth metals
• Magnesium (Z=12): Same group (Group 2), different period
• Strontium (Z=38): Same group (Group 2), different period
• All Group 2 elements have: 2 valence electrons, form M²⁺ ions, similar chemical reactions

Others:
• Potassium (Z=19): Group 1 (alkali metal) - different properties
• Scandium (Z=21): Transition metal - different properties

Aspect	Mendeléev's Periodic Table	Modern Periodic Table
Basis of classification	Atomic mass	Atomic number
Periodic Law	Properties are periodic function of atomic mass	Properties are periodic function of atomic number
Position of isotopes	Problem (same element, different masses)	Correctly placed (same atomic number)
Position of hydrogen	No fixed position (confusion)	Group 1 (and sometimes 17) - still ambiguous
Anomalous pairs	Ar-K, Co-Ni, Te-I (mass order violated)	Correct order based on atomic number
Groups	Groups I to VIII (and subgroups)	Groups 1 to 18
Prediction of new elements	Successfully predicted (eka-aluminium, etc.)	Based on electronic configuration
Explanation of periodicity	No explanation given	Based on electronic configuration

Yes, some Döbereiner's triads existed in Newlands' Octaves.

Döbereiner's Triads (1817):
• Lithium (Li), Sodium (Na), Potassium (K)
• Calcium (Ca), Strontium (Sr), Barium (Ba)
• Chlorine (Cl), Bromine (Br), Iodine (I)

Newlands' Octaves (1866) - Part of his table:

Comparison:
1. Li, Na, K triad: Li (row1), Na (row2), K (row3) - all in "re" column ✓
2. Ca, Sr, Ba triad: Only Ca appears (row3), Sr/Ba not in early rows
3. Cl, Br, I triad: Cl (row3), Br appears later, I not in first 3 rows

Conclusion: Newlands extended Döbereiner's idea to octaves (every 8th element), but his table was limited to known elements and had inconsistencies.

Limitations of Döbereiner's Triads (1817):

1. Limited scope: Only identified 3 triads from 30 known elements:
  • Li, Na, K
  • Ca, Sr, Ba
  • Cl, Br, I

2. Could not classify all elements: Many elements didn't fit into triads

3. Not a comprehensive system: No overall arrangement of all elements

4. Atomic mass measurements: Some atomic masses were inaccurate at that time

5. New element discovery: As more elements were discovered, triad pattern wasn't universal

6. No prediction power: Didn't lead to predictions of new elements

Historical significance: First attempt to relate properties to atomic mass, inspired later classifications.

Limitations of Newlands' Law of Octaves (1866):

1. Applicable only to light elements: Worked only up to calcium (atomic mass 40)

2. Assumed only 56 elements: Believed no more elements would be discovered

3. Forced fitting: Placed unlike elements together to maintain octave pattern
  • Example: Co and Ni in same slot as F, Cl, Br
  • Fe placed far from Co and Ni (its similar elements)

4. No place for noble gases: Not discovered then, couldn't be accommodated

5. No explanation: Didn't explain why properties repeated every 8th element

6. Ridiculed by contemporaries: Compared to musical analogy wasn't taken seriously

7. Anomalies with heavier elements: Properties didn't match after certain point

Historical significance: First attempt at periodicity based on atomic mass, precursor to modern periodic table.

Using Mendeléev's Table (Table 5.4 in NCERT):

Group trends in Mendeléev's Table:
• Group I: R₂O
• Group II: RO
• Group III: R₂O₃
• Group IV: RO₂

Predictions:
1. Potassium (K): Group I → Formula: K₂O
  (Actually correct: Potassium oxide)

2. Carbon (C): Group IV → Formula: CO₂
  (Actually correct: Carbon dioxide; also forms CO)

3. Aluminium (Al): Group III → Formula: Al₂O₃
  (Actually correct: Aluminium oxide)

4. Silicon (Si): Group IV → Formula: SiO₂
  (Actually correct: Silicon dioxide)

5. Barium (Ba): Group II → Formula: BaO
  (Actually correct: Barium oxide; also forms BaO₂)

Mendeléev's success: His table allowed such predictions before elements were thoroughly studied.

Mendeléev left gaps for undiscovered elements and named them with prefix "eka" (Sanskrit for "one"):

1. Eka-boron → Scandium (Sc)
• Predicted: 1871
• Discovered: 1879 by Lars Fredrik Nilson
• Properties matched predictions

2. Eka-aluminium → Gallium (Ga)
• Predicted: 1871
• Discovered: 1875 by Paul Émile Lecoq de Boisbaudran
• Properties matched remarkably well

3. Eka-silicon → Germanium (Ge)
• Predicted: 1871
• Discovered: 1886 by Clemens Winkler
• Properties matched closely

Other predicted elements:
• Eka-manganese → Technetium (Tc, discovered 1937)
• Dvi-manganese → Rhenium (Re, discovered 1925)

Significance: Successful predictions validated Mendeléev's periodic table.

Mendeléev's criteria (1869):

1. Atomic mass as primary basis: Arranged elements in increasing order of atomic mass

2. Chemical properties similarity: Grouped elements with similar properties together

3. Compounds formed: Focused on hydrides and oxides formulae
  • Example: Group I forms R₂O and RH
  • Group IV forms RH₄ and RO₂

4. Periodic recurrence: Noted properties repeated at regular intervals

5. Left gaps for undiscovered elements: Predicted properties of missing elements

6. Sometimes inverted order: Placed heavier element before lighter when properties demanded
  • Example: Co (58.9) before Ni (58.7)
  • Example: Te (127.6) before I (126.9)

7. Used cards for each element: Wrote properties on cards, arranged and rearranged

Result: Published first periodic table in 1872 with 63 known elements.

Noble gases (He, Ne, Ar, Kr, Xe, Rn) are placed in Group 18 because:

1. Unique electronic configuration: Have complete valence shell (octet, except He: duplet)
  • He: 2
  • Ne: 2,8
  • Ar: 2,8,8
  • Kr: 2,8,18,8

2. Extremely low reactivity: Chemically inert under normal conditions
  • Previously called "inert gases"
  • Form compounds only under extreme conditions (Xe compounds known)

3. Zero valency: Don't gain, lose, or share electrons easily

4. Monatomic gases: Exist as single atoms, not diatomic molecules

5. Similar physical properties: Colorless, odorless, low boiling points

6. Historical discovery: Discovered late (1890s onwards) because of low abundance and reactivity

7. Placement in periodic table: Between highly electronegative halogens (Group 17) and highly electropositive alkali metals (Group 1)

Modern understanding: Their stability is due to completely filled orbitals, making them perfect bridge between periods.

Anomalies removed by Modern Periodic Table (Moseley, 1913):

Anomaly in Mendeléev's Table	Solution in Modern Table
Position of isotopes Same element, different atomic masses	Atomic number same → same position
Anomalous pairs Ar (39.9) before K (39.1) Co (58.9) before Ni (58.7) Te (127.6) before I (126.9)	Atomic number order correct: Ar (18) before K (19) Co (27) before Ni (28) Te (52) before I (53)
No explanation for periodicity	Periodicity explained by electronic configuration
Position of hydrogen ambiguous	Placed in Group 1 (and sometimes 17) based on config
Lanthanides & Actinides problem	Placed separately at bottom
Transition elements placement	d-block clearly identified

Key change: Basis changed from atomic mass to atomic number (proton number).

Elements similar to magnesium (Mg):

1. Calcium (Ca)
2. Strontium (Sr)
Also: Barium (Ba), Beryllium (Be) - all Group 2 elements

Basis for choice - Same Group (Group 2):

Electronic configuration similarity:
• Mg: [Ne] 3s² (2,8,2)
• Ca: [Ar] 4s² (2,8,8,2)
• Sr: [Kr] 5s² (2,8,18,8,2)
• All have 2 valence electrons in s-orbital

Similar chemical properties:
1. Reaction with oxygen: Form oxides MO
  2Mg + O₂ → 2MgO
  2Ca + O₂ → 2CaO

2. Reaction with water: Form hydroxides + H₂
  Mg + 2H₂O → Mg(OH)₂ + H₂ (slow with hot water)
  Ca + 2H₂O → Ca(OH)₂ + H₂ (faster)

3. Reaction with acids: Form salts + H₂
  Mg + 2HCl → MgCl₂ + H₂
  Ca + 2HCl → CaCl₂ + H₂

4. Ion formation: Lose 2 electrons to form M²⁺ ions

Group 2 = Alkaline earth metals

First ten elements (Atomic numbers 1-10):

1. Hydrogen (H) - Non-metal (sometimes behaves like metal)
2. Helium (He) - Noble gas (non-metal)
3. Lithium (Li) - Metal (alkali metal)
4. Beryllium (Be) - Metal (alkaline earth metal)
5. Boron (B) - Metalloid/Semi-metal
6. Carbon (C) - Non-metal
7. Nitrogen (N) - Non-metal
8. Oxygen (O) - Non-metal
9. Fluorine (F) - Non-metal (halogen)
10. Neon (Ne) - Noble gas (non-metal)

Metals among first ten: Lithium (Li) and Beryllium (Be)

Periodic trends in Period 2:
• Left side: Metals (Li, Be)
• Middle: Metalloid (B)
• Right side: Non-metals (C, N, O, F, Ne)
• Extreme right: Noble gas (Ne)

Properties of these metals:
• Li: Soft, reactive, stored in oil, forms Li⁺ ions
• Be: Hard, lightweight, forms Be²⁺ ions, amphoteric oxide