KEAM – Chemistry
Classification of Elements & Periodic Properties
Goal: Master the classification schemes and periodic trends that frequently appear on the KEAM exam.
Approach:
- Conceptual understanding of why trends exist.
- Derivations using simple quantum‑mechanical arguments.
- Key formulas & mnemonic tricks for quick recall.
- Tricky points that often trip candidates.
- A hard practice problem to test your grasp.
1. Classification of Elements
| Classification Type | Basis | Classic Examples | Notes for KEAM |
|---|---|---|---|
| Metals / Non‑Metals / Metalloids | Physical & electronic properties | Metals – K, Fe, Ag; Non‑Metals – C, O, N; Metalloids – Si, Ge, As | KEAM often asks “Which group contains most metalloids?” (Group 14 & 15). |
| Based on electronic configuration | Number of valence electrons | s‑block (groups 1–2), p‑block (groups 13–18), d‑block (groups 3–12), f‑block | Use to explain oxidation states & trends. |
| Based on the dominant valence shell | Principal quantum number (n) | Alkali metals (n=2–6), alkaline earths (n=2–6), transition metals (n=3–5) | Helps explain ionization energies. |
| Based on valence‑electron count | 1–8 valence electrons | Group 1: 1, Group 2: 2, …, Group 8: 8 | Useful for predicting typical oxidation states. |
| Based on physical state at 25 °C | Solid, liquid, gas | 1 – 2 – 3 – 4 | KEAM may ask “Which group does the liquid element belong to?” (Group 12). |
| Based on electronic affinity | Electron‑affinity sign & magnitude | 4 → 1 → 8 → 7 → 6 (most negative) | Helpful for periodic trend analysis. |
2. Periodic Properties & Their Trends
| Property | Trend Across a Period (left → right) | Trend Down a Group (top → bottom) | Key Equation / Concept |
|---|---|---|---|
| Atomic / Ionic Radius | ↓ (except for 3 d transition metals) | ↑ | ( r \approx \frac{a_0 (n+1)}{Z_{\text{eff}}} ) |
| First Ionization Energy (IE₁) | ↑ (except for 3 d transition metals) | ↓ | ( \text{IE}1 \propto \frac{Z{\text{eff}}^2}{n^2} ) |
| Electron Affinity | ↑ (most negative for group 8) | ↓ | No simple formula; trend follows ( \Delta H_f^\circ ). |
| Electronegativity (Pauling) | ↑ | ↓ | ( \chi = 0.359 \Delta H_f^\circ + 0.7 ) (Pauling) |
| Standard Electrode Potential | ↑ (more positive for non‑metals) | ↓ | Depends on half‑reaction, not a simple trend. |
| Oxidation States | More varied (especially for transition metals) | Tend to increase in magnitude | Governed by d‑electron count. |
| Bond Energy (C–C, C–H) | ↓ (except for 3 d transition metals) | ↑ | ( E_{\text{bond}} \propto \frac{Z_{\text{eff}}}{r} ) |
Tricky Point:
The 3 d transition metals (Sc–Zn) display anomalous trends: atomic radius increases and IE₁ decreases across the period because of poor shielding by 3d electrons. Remember this when answering “Which element has the largest radius among the 3 d metals?”
3. Derivations & Core Concepts
3.1 Effective Nuclear Charge (Z_eff)
-
Slater’s Rules (quick reference for Z_eff calculation)
- Write the electronic configuration.
- Group electrons: (1s), (2s,2p), (3s,3p), (3d), (4s,4p), …
- For an electron in group i:
- Contributions from electrons in same group: 0.35 (except 1s → 0.30).
- From groups inner than i: 1.00 each.
- From groups outer than i: 0.00.
- Z_eff = Z – Σ(shielding)
-
Application:
For Na⁺ (1s² 2s¹ → 2s² after ionisation),
( Z_{\text{eff}} = 11 – (2×