3.1 Why Do We Need to Classify Elements?
- Over 100 elements are known, each with unique properties.
- Classification is important to:
- Identify patterns in chemical and physical properties.
- Predict properties of unknown elements.
- Make study systematic and easy.
- Early chemists observed recurring patterns (periodicity) in elements’ properties.
3.2 Genesis of Periodic Classification
- Dmitri Mendeleev (1869): Arranged elements by increasing atomic mass.
- Left gaps for undiscovered elements and predicted their properties.
- Example: Eka-Aluminium → Gallium, Eka-Silicon → Germanium.
- Limitations: Some elements did not fit by atomic mass (e.g., iodine and tellurium).
Other Early Classifications:
- Dobereiner’s Triads: Groups of three elements with average properties.
- Newlands’ Law of Octaves: Every eighth element has similar properties.
3.3 Modern Periodic Law and the Present Form of the Periodic Table
- Modern Periodic Law: Properties of elements are a periodic function of their atomic numbers.
- Moseley (1913): Proved atomic number, not mass, is fundamental.
Features of Modern Periodic Table:
- 18 groups (vertical) and 7 periods (horizontal).
- Elements in a group → Similar valence electron configuration → Similar chemical properties.
- Elements arranged in increasing atomic number.
- s-, p-, d-, f-block elements classified based on their valence orbitals.
3.4 Nomenclature of Elements with Atomic Numbers > 100
- Elements beyond atomic number 100 are synthetic (produced in labs).
- IUPAC Temporary Naming:
- Names based on digits of atomic number + “-ium”.
- Example: 101 → Unnilunium (Un=1, Nil=0, Un=1) → Later named Mendelevium (Md).
- Names based on digits of atomic number + “-ium”.
- Used until official names are confirmed.
3.5 Electronic Configurations of Elements and the Periodic Table
- Electronic configuration: Distribution of electrons in shells and subshells.
- Determines:
- Position in the periodic table
- Chemical properties and reactivity
- Periodic trends like ionization energy and atomic radius
Example Configurations:
- Sodium (Na, Z=11): 1s² 2s² 2p⁶ 3s¹
- Chlorine (Cl, Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵
3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f- Blocks
| Block | Groups | Valence Electrons | Examples |
|---|---|---|---|
| s-block | 1–2 + He | s-orbital | Li, Be, Na, Mg |
| p-block | 13–18 | p-orbital | B, C, N, O, F, Ne |
| d-block | 3–12 | d-orbital | Fe, Cu, Zn |
| f-block | Lanthanoids, Actinoids | f-orbital | La, U, Th |
Key Notes:
- s- and p-block → Main group elements
- d-block → Transition elements
- f-block → Inner transition elements
3.7 Periodic Trends in Properties of Elements
1. Atomic Radius
- Distance from nucleus to outermost electron.
- Across a period: Decreases (more protons → stronger attraction).
- Down a group: Increases (new shells added).
2. Ionization Energy
- Energy required to remove an electron.
- Across a period: Increases
- Down a group: Decreases
3. Electron Affinity
- Energy released when an atom gains an electron.
- Becomes more negative across a period, less negative down a group.
4. Electronegativity
- Ability to attract electrons in a bond.
- Across a period: Increases
- Down a group: Decreases
5. Metallic and Non-metallic Character
- Metals: Low ionization energy, left side of periodic table.
- Non-metals: High ionization energy, right side.
- Metallic character decreases across a period, increases down a group.
Quick Summary Table:
| Property | Across Period | Down Group |
|---|---|---|
| Atomic Radius | ↓ | ↑ |
| Ionization Energy | ↑ | ↓ |
| Electron Affinity | More negative | Less negative |
| Electronegativity | ↑ | ↓ |
| Metallic Character | ↓ | ↑ |