1. History of Classification

1. Dobereiner's Triads (1817):

Arranged elements in groups of three. The atomic weight of the middle element was approx the average of the other two.
Example: Li (7), Na (23), K (39). (7+39)/2 = 23.

2. Newlands' Law of Octaves (1865):

Every 8th element had properties similar to the 1st (like musical notes).
Limitation: Valid only up to Calcium (Atomic mass 40).

3. Mendeleev's Periodic Table:

Law: Properties are a periodic function of their Atomic Masses.

Merits: Predicted existence of undiscovered elements (Eka-Aluminium = Gallium, Eka-Silicon = Germanium). Left gaps.
Demerits: Position of Isotopes, Hydrogen's position, Anomalous pairs (Ar-K, Co-Ni).

2. Modern Periodic Table

Modern Periodic Law (Henry Moseley):

Physical and chemical properties of elements are a periodic function of their Atomic Numbers.

Structure (Long Form)

Periods (Horizontal Rows): 7 Periods. Represent principal quantum number (n).
Groups (Vertical Columns): 18 Groups. Elements have same valence shell configuration.

Block	General Config	Nature
s-block	ns^1-2	Reactive Metals
p-block	ns²np^1-6	Metals, Non-metals, Metalloids
d-block	(n-1)d^1-10ns^0-2	Transition Metals
f-block	(n-2)f^1-14...ns²	Inner Transition

[Image of Modern Periodic Table Structure]

3. Periodic Trends

1. Atomic Radius:

Across Period: Decreases (Effective nuclear charge Z_eff increases).
Down Group: Increases (Number of shells increases).
Order: Covalent < Metallic < Van der Waals radius.

2. Ionic Radius:

Cation: Smaller than parent atom (Cation < Atom).
Anion: Larger than parent atom (Anion > Atom).
Isoelectronic Species: Radius ∝ 1/Z (Atomic Number).
Example: Al³⁺ < Mg²⁺ < Na⁺ < F^- < O^2-.

3. Ionization Energy (IE):

Energy required to remove an electron from gaseous atom.
Trend: Increases across period, Decreases down group.

Exceptions:
Be > B (Penetration effect of 2s)
N > O (Half-filled stability of 2p³)

4. Electron Affinity (EGE):

Energy released when electron is added.
Highest EGE: Chlorine (Cl), NOT Fluorine. (F has small size repulsion).
Noble Gases: Positive EGE (Stable octet).

5. Electronegativity (EN):

Tendency to attract shared pair of electrons.
Most EN Element: Fluorine (4.0).
Trend: Increases across period, Decreases down group.

[Image of Periodic Trends Chart]

4. Chemical Behavior

Valency

Across Period: Increases from 1 to 4 with respect to Oxygen, then decreases. With respect to Hydrogen, increases 1 to 4 then decreases to 1.
Down Group: Remains constant.

Metallic Character

Across Period: Decreases (Non-metallic increases).
Down Group: Increases (Metallic increases).

Nature of Oxides

Basic	Na₂O, MgO (Metals)
Acidic	Cl₂O₇, SO₃ (Non-metals)
Amphoteric	Al₂O₃, ZnO, BeO
Neutral	CO, NO, N₂O

Numericals & HOTS

Q1. Element Position

An element has the electronic configuration [Kr] 4d¹⁰ 5s² 5p³. Identify its Period and Group number in the Modern Periodic Table.

Solution:
1. Period: Determined by the highest Principal Quantum Number (n).
Here, max n = 5. So, Period = 5.

2. Group: Last electron enters p-orbital.
For p-block, Group Number = 10 + Valence Electrons (ns + np).
Valence Electrons = 2 (from 5s) + 3 (from 5p) = 5.
Group = 10 + 5 = 15.

Q2. Radius Comparison

Arrange the following isoelectronic species in increasing order of ionic radius: O^2-, F^-, Na⁺, Mg²⁺. Explain why.

Solution:
All have 10 electrons.
Protons (Z): O(8), F(9), Na(11), Mg(12).

Concept: Radius is inversely proportional to Atomic Number (Z) for isoelectronic species.
Mg²⁺ has the highest nuclear pull (12 protons pulling 10 electrons), so it is smallest.
O^2- has the lowest nuclear pull (8 protons pulling 10 electrons), so it is largest.

Order: Mg²⁺ < Na⁺ < F^- < O^2-

Q3. Valence Electron Identification

The successive ionization energies (IE) of an element 'X' are: 740, 1450, 7730, 10500 kJ/mol. To which group does 'X' likely belong?

Solution:
Jump in IE:
IE₁ to IE₂: Small jump (740 → 1450).
IE₂ to IE₃: Massive jump (1450 → 7730).

Conclusion: The massive jump indicates that removing the 3rd electron breaks a stable noble gas core. Therefore, the element has 2 valence electrons.
Group: 2 (Alkaline Earth Metals)

Q4. Magnetic Moment

Calculate the spin-only magnetic moment of a divalent ion (M²⁺) if its atomic number is 25.

Solution:
Z = 25 (Manganese).
Config of Mn: [Ar] 3d⁵ 4s².
Config of Mn²⁺: Remove 2e^- from 4s → [Ar] 3d⁵.

Number of unpaired electrons (n) in 3d⁵ = 5.
μ = √[n(n+2)] = √[5(5+2)]
μ = √35
μ ≈ 5.92 B.M.

Q5. Amphoteric Identification

Identify the amphoteric oxides from the following list: Na₂O, CO, Al₂O₃, SO₃, ZnO, PbO.

Solution:
1. Na₂O: Strong Basic.
2. CO: Neutral.
3. SO₃: Acidic.
4. Al₂O₃, ZnO, PbO: These react with both acids (forming salts) and bases (forming complexes).

Ans: Al₂O₃, ZnO, PbO

Q6. Group 16/17 Trends (HOTS)

Arrange O, S, F, and Cl in decreasing order of Electron Gain Enthalpy (magnitude).

Solution:
Concept 1: Halogens (Group 17) have higher EGE than Chalcogens (Group 16). So, (F, Cl) > (O, S).
Concept 2: 2nd period elements (O, F) have lower EGE than 3rd period (S, Cl) due to small size repulsion.
So, Cl > F and S > O.

Combining: Cl has highest. O has lowest.
Order: Cl > F > S > O

Q7. IUPAC Name (Z > 100)

What is the IUPAC name and symbol for the element with atomic number 119?

Solution:
Digits: 1 (un), 1 (un), 9 (enn).
Suffix: -ium.

Name: Un + un + enn + ium = Ununennium.
Symbol: Uue.

Q8. Slater's Rule (HOTS)

Calculate the effective nuclear charge (Z_eff) experienced by a 4s electron in Potassium (Z=19).

Solution:
Config: (1s²) (2s² 2p⁶) (3s² 3p⁶) (4s¹).
We calculate shielding for the 4s electron.

1. (n-1) shell (3s, 3p): 8 electrons. Shielding = 8 × 0.85 = 6.80.
2. (n-2) and inner shells (1s, 2s, 2p): 10 electrons. Shielding = 10 × 1.00 = 10.00.

Total Screening (σ) = 16.80.
Z_eff = Z - σ = 19 - 16.80
Z_eff = 2.20

Q9. Bond Length

The bond length of Cl-Cl is 198 pm and C-C is 154 pm. Calculate the bond length of C-Cl.

Solution:
Covalent Radius of Cl (r_Cl) = 198 / 2 = 99 pm.
Covalent Radius of C (r_C) = 154 / 2 = 77 pm.

Bond Length C-Cl = r_C + r_Cl
= 77 + 99
Ans: 176 pm

Q10. Nitrogen vs Oxygen (HOTS)

IE₁ of Nitrogen is greater than Oxygen. However, IE₂ of Oxygen is greater than Nitrogen. Why?

Solution:
For IE₁: N (2p³) is half-filled stable. O (2p⁴) is unstable. So N > O.

For IE₂:
N⁺ config is 2p².
O⁺ config is 2p³ (Half-filled stable).
Removing 2nd electron from stable O⁺ requires more energy than removing from N⁺.
So, IE₂ (O) > IE₂ (N).

Important Relations

1. Effective Nuclear Charge (Z_eff)

Z_eff = Z - σ

(Z = Atomic Number, σ = Screening Constant)

Trend: Increases by 0.65 across a period (left to right).

2. Isoelectronic Species Radius

For species with same number of electrons:

Radius ∝ 1 / Atomic Number (Z)

Anion > Neutral > Cation

Ex: N^3- > O^2- > F^- > Na⁺ > Mg²⁺

3. Spin Only Magnetic Moment (μ)

μ = √[n(n + 2)] B.M.

(n = Number of unpaired electrons)
B.M. = Bohr Magneton

4. Electronegativity Relationships

Mulliken's Scale (X_M):

X_M = (IE + EA) / 2

Relation with Pauling (X_P):

X_P ≈ X_M / 2.8

(Where IE and EA are in eV/atom)

20 Golden Facts (NEET/JEE)

1. Highest Electron Affinity: Chlorine (Cl) has the highest electron affinity, NOT Fluorine. F has a small size, leading to inter-electronic repulsion, making it harder to add an electron.
2. IE₁ Exception (Period 2): Order is Li < B < Be < C < O < N < F < Ne.
Be > B (Full filled 2s orbital penetration).
N > O (Half filled 2p orbital stability).
3. Lanthanoid Contraction: The atomic radii of 4d and 5d series elements in the same group are almost identical (e.g., Zr ≈ Hf, Nb ≈ Ta) due to poor shielding by 4f electrons.
4. Noble Gas Radii: Noble gases have the largest atomic radii in their respective periods because they are measured as Van der Waals radii, which are larger than Covalent radii.
5. Amphoteric Oxides: Remember the list: Zn, Al, Be, Sn, Pb, Ga (Znabhi Aliyona Gayi Punjabi Song). Examples: ZnO, Al₂O₃, BeO, SnO₂, PbO.
6. Neutral Oxides: Only three common neutral non-metal oxides: CO, NO, N₂O (Laughing gas).
7. Highest Density: Osmium (Os) and Iridium (Ir) have the highest densities (~22.6 g/cm³).
8. Liquid Elements:
Metal: Mercury (Hg).
Non-metal: Bromine (Br).
Metals liquid just above room temp: Ga, Cs, Fr.
9. Most Electronegative: Fluorine (4.0).
Most Electropositive: Cesium (Cs).
10. Diagonal Relationship:
Li ↔ Mg
Be ↔ Al
B ↔ Si
Due to similar ionic potential (Charge/Size ratio).
11. Successive IE: IE₂ is always > IE₁ because it is harder to remove an electron from a positive ion.
Huge jump in IE occurs when breaking a noble gas configuration (e.g., Na IE₁ is low, IE₂ is extremely high).
12. Highest Melting Point: Tungsten (W) among metals. Carbon (Diamond) among non-metals.
13. Group 16 & 17 EA:
Group 16: S > Se > Te > Po > O (O is lowest).
Group 17: Cl > F > Br > I.
14. Inert Pair Effect: Heavier p-block elements (Tl, Pb, Bi) show stability in oxidation states that are 2 units lower than the group valence (e.g., Tl⁺¹, Pb⁺², Bi⁺³).
15. Acidic Character of Oxides: Increases across a period (Na₂O < MgO < ... < Cl₂O₇). Increases with oxidation state (MnO < Mn₂O₇).
16. Order of Screening Effect: s > p > d > f. Due to their shape, 's' orbitals shield nucleus most effectively, 'f' orbitals shield least (leading to lanthanoid contraction).
17. Coinage Metals: Group 11 (Cu, Ag, Au). Also known as noble metals due to low reactivity.
18. Zero Group: Mendeleev added the Zero group (Noble gases) later without disturbing the original table.
19. Representative Elements: s-block and p-block elements (except Noble gases) are collectively called Representative Elements.
20. Transuranic Elements: Elements with atomic number Z > 92 (Uranium). They are all synthetic and radioactive.

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