Title: Periodic Classification of Elements (Class 11 CBSE)
Introduction
The periodic classification of elements is the systematic arrangement of elements based on their atomic number, properties, and electron configuration. This classification helps in understanding trends in chemical behavior and predicting the properties of unknown elements. The modern periodic table is an essential tool in chemistry and is based on the periodic law.
Early Attempts of Classification
- Dobereiner’s Triads (1817)
- Elements were grouped in sets of three, called triads, based on their atomic masses.
- The atomic mass of the middle element was approximately the average of the other two.
- Example: Li (6.9), Na (23), K (39) → Na’s mass ≈ (6.9 + 39)/2.
- Limitation: Could only classify a few elements.
- Newlands’ Law of Octaves (1864)
- Elements were arranged in increasing order of atomic mass.
- Every eighth element had similar properties (like musical notes).
- Limitation: Failed after calcium, as heavier elements did not follow the pattern.
- Mendeleev’s Periodic Table (1869)
- Elements were arranged in increasing atomic mass.
- Elements with similar properties were placed in groups.
- Left gaps for undiscovered elements (e.g., eka-aluminium → later discovered as gallium).
- Limitation: No fixed place for isotopes; some elements were placed in incorrect order.
Modern Periodic Table
- Proposed by Henry Moseley (1913), based on atomic number rather than atomic mass.
- Modern Periodic Law: The physical and chemical properties of elements are periodic functions of their atomic number.
Structure of the Modern Periodic Table
- Groups (Vertical Columns): 18 groups based on valence electrons.
- Periods (Horizontal Rows): 7 periods based on the number of electron shells.
- Block Classification:
- s-block (Groups 1 & 2): Alkali and alkaline earth metals.
- p-block (Groups 13-18): Includes halogens and noble gases.
- d-block (Transition elements): Middle of the periodic table.
- f-block (Lanthanides & Actinides): Placed separately at the bottom.
Periodic Trends in the Periodic Table
1. Atomic Radius
- Decreases across a period (due to increased nuclear charge pulling electrons closer).
- Increases down a group (due to addition of electron shells).
2. Ionization Energy (IE)
- Increases across a period (more energy needed to remove electrons from a more stable atom).
- Decreases down a group (outer electrons are farther from the nucleus and easier to remove).
3. Electron Affinity
- Increases across a period (non-metals have higher attraction for electrons).
- Decreases down a group (larger atoms have weaker attraction for additional electrons).
4. Electronegativity
- Increases across a period (elements tend to attract electrons more strongly).
- Decreases down a group (due to increase in atomic size and shielding effect).
5. Metallic and Non-Metallic Character
- Metallic character decreases across a period (left to right, elements become more non-metallic).
- Increases down a group (as atoms lose electrons more easily).
6. Reactivity Trends
- Metals: More reactive down a group (e.g., alkali metals like Cs are more reactive than Li).
- Non-metals: More reactive up a group (e.g., fluorine is the most reactive halogen).
Importance of the Periodic Table
- Helps in predicting element properties.
- Guides the discovery of new elements.
- Essential in industries, medicine, and environmental studies.
Conclusion
The periodic classification of elements is one of the most important achievements in chemistry. The modern periodic table, arranged by atomic number, helps in understanding element behavior and trends systematically. It continues to evolve with the discovery of new elements and advancements in quantum mechanics.