📘 Chapter 3 – Classification of Elements and Periodicity in Properties
1. Why Classification of Elements is Needed
By 1800 → only 31 elements known.
Today → more than 118 elements.
To study them systematically → classification was essential.
Aim: group elements with similar properties together.
2. Early Attempts at Classification
(a) Dobereiner’s Triads (1829)
Groups of 3 elements with similar properties.
Middle element’s atomic mass ≈ average of other two.
Example:
Li (7), Na (23), K (39)
(7+39)/2 = 23 ✅
Limitation → not all elements fit into triads.
(b) Newlands’ Law of Octaves (1866)
Arranged elements in increasing atomic mass.
Every 8th element → similar properties (like musical notes).
Example: Li, Be, B, C, N, O, F → Na resembles Li.
Limitation → worked only up to Calcium; failed for heavier elements.
3. Mendeleev’s Periodic Table (1869)
Based on atomic mass.
Periodic law: “Properties of elements are a periodic function of their atomic masses.”
Features:
Rows = Periods, Columns = Groups.
Left gaps for undiscovered elements → predicted Sc, Ga, Ge correctly.
Distinguished metals & non-metals.
Limitations:
Position of isotopes not explained.
Increasing atomic mass order sometimes violated (e.g., Co & Ni).
No clear place for Hydrogen.
4. Modern Periodic Law (Moseley, 1913)
After discovery of protons.
Law: “Properties of elements are a periodic function of their atomic number (Z).”
Solved Mendeleev’s problems → isotopes have same Z, so same place.
5. Modern Periodic Table (Long Form)
Arranged by atomic number.
18 Groups, 7 Periods.
Based on electronic configuration.
Groups: Vertical columns → same valence electrons.
Periods: Horizontal rows → same number of shells.
6. Classification of Elements (Blocks)
Depending on last electron entry:
s-block: Groups 1 & 2 (alkali & alkaline earth metals).
p-block: Groups 13–18 (includes metals, non-metals, metalloids, noble gases).
d-block: Groups 3–12 (transition metals).
f-block: Lanthanides + Actinides (inner transition metals).
7. Periodic Trends in Properties
(a) Atomic Radius
Half distance between nuclei of two bonded atoms.
Across a period → decreases (↑ nuclear charge, same shell).
Down a group → increases (new shells added).
(b) Ionic Radius
Cations: smaller than parent atom (loss of e⁻ → stronger pull).
Anions: larger than parent atom (gain of e⁻ → more repulsion).
(c) Ionization Enthalpy (IE)
Energy required to remove the outermost electron.
Across period → increases (nuclear charge ↑, atomic size ↓).
Down group → decreases (shielding ↑, atomic size ↑).
(d) Electron Gain Enthalpy (EGE)
Energy change when an atom gains an electron.
Across period → more negative (nuclear charge ↑).
Down group → less negative (size ↑, shielding ↑).
Exception: Noble gases (positive EGE, no tendency to gain e⁻).
(e) Electronegativity
Tendency to attract shared pair of electrons in a bond.
Across period → increases.
Down group → decreases.
Most electronegative → Fluorine (4.0, Pauling scale).
8. Anomalous Properties of Second Period Elements
Elements: Li, Be, B, C, N, O, F, Ne.
Show differences compared to heavier group members due to:
Small size.
High electronegativity.
Absence of d-orbitals.
Example:
Li shows similarity with Mg (Diagonal Relationship).
Be shows similarity with Al.
9. Periodic Trends in Metallic/Non-metallic Character
Metallic character: ability to lose e⁻.
Non-metallic character: ability to gain e⁻.
Across a period → metallic ↓, non-metallic ↑.
Down a group → metallic ↑, non-metallic ↓.