Classification of Elements | IISER Smart Prep

1. Core Concept

The properties of elements are periodic functions of their atomic number (Z). The modern periodic table groups elements with similar valence shell electronic configurations together, leading to predictable trends in physical and chemical behavior.

2. Genesis of Periodic Classification

The quest for classification began with the discovery of elements. Key historical milestones include:

Dobereiner's Triads (1817): Grouped 3 elements with similar properties. Atomic mass of middle element ≈ average of other two. (e.g., Li, Na, K).
Newlands' Law of Octaves (1866): Elements arranged by mass; every 8th element had similar properties (like musical notes). Failed after Calcium.
Mendeleev's Periodic Law (1869): "Properties are periodic functions of atomic masses." Predicted unknown elements (Eka-Boron, Eka-Aluminium) with remarkable accuracy.
Lothar Meyer: Plotted atomic volume vs atomic mass; similar elements occupied similar positions on the curve.

3. Modern Periodic Law & Present Form

Modern Periodic Law (Moseley): "The physical and chemical properties of the elements are periodic functions of their atomic numbers (Z)."

Based on X-ray spectra: √ν = a(Z - b).

Structure of the Table:

Periods (Horizontal): 7 Periods. Period number = highest principal quantum number (n).
Groups (Vertical): 18 Groups. Elements in a group have the same valence electron configuration.

4. Electronic Configurations & Blocks

Types of Elements:

Block	Groups	Valence Configuration	Characteristics
s-block	1 & 2	ns^1-2	Reactive metals, low IE, form ionic compounds.
p-block	13 to 18	ns² np^1-6	Metals, non-metals, metalloids. Chalcogens, Halogens.
d-block	3 to 12	(n-1)d^1-10 ns^0-2	Transition elements, variable oxidation, colored ions.
f-block	Bottom row	(n-2)f^1-14 (n-1)d^0-1 ns²	Inner transition (Lanthanoids/Actinoids). Heavy metals.

Quick Identification:

Period: The max value of 'n' in configuration.
Block: The subshell (s/p/d/f) that receives the last electron.
Group:
- s-block: No. of valence electrons.
- p-block: 10 + No. of valence electrons.
- d-block: No. of (n-1)d + ns electrons.

5. Periodic Trends in Properties

Z_eff = Z - σ

Effective Nuclear Charge: Net positive charge experienced by an electron. Increases across a period.

1. Atomic & Ionic Radii:

Type: Covalent radius < Metallic radius < Van der Waals radius.
Trend: Increases down group (Shells ↑); Decreases across period (Z_eff ↑).
Ionic Radius: Cation < Parent Atom < Anion. (e.g., Fe³⁺ < Fe²⁺ < Fe).
Isoelectronic series: Radius ∝ 1/Z. (e.g., O^2- > F^- > Na⁺ > Mg²⁺).

2. Ionization Enthalpy (IE):

Energy required to remove an electron from a gaseous atom.

Trend: Increases across period; Decreases down group.
Exceptions:
- Be > B: Be has stable 2s² (penetration effect).
- N > O: N has stable half-filled 2p³ configuration.
- Mg > Al and P > S follow same logic.

3. Electron Gain Enthalpy (Δ_egH):

Across period: Becomes more negative.
Down group: Becomes less negative.
Exception: Cl > F and S > O. Small size of F/O causes inter-electronic repulsion, making electron entry harder than Cl/S.

4. Electronegativity:

Ability to attract shared pair. F (4.0) is the most electronegative. Cesium is have least.

Trend: Increases across period; Decreases down group.

6. Chemical Periodicity & Reactivity

Diagonal Relationship: Li/Mg, Be/Al, B/Si show similar properties due to target charge/radius ratio.
Oxidation State: Period 2 elements show max covalency of 4 (no d-orbitals). Lower elements can expand octet.
Nature of Oxides:
- Left (Metals): Basic (MgO, Na₂O).
- Middle: Amphoteric (Al₂O₃, As₂O₃).
- Right (Non-metals): Acidic (Cl₂O₇, SO₃).

7. IAT Exam Focus

Critical Catch-ups:

Isoelectronic Radius: Always sort by nuclear charge Z. Higher Z = smaller size.
IE Exceptions: Always check configurations for half-filled/full-filled stability.
Shielding Effect (σ): f < d < p < s. Poor shielding of d/f electrons causes Lanthanoid Contraction, making Zr ≈ Hf.
Metallic Character: Increases down; Decreases across.

8. Practice Mock Test

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Classification of Elements & Periodicity