Introduction
The classification of elements is a systematic arrangement of elements based on their properties. It helps in studying a large number of elements in an organized way. The modern classification is based on the atomic number, which determines the electronic configuration and properties of elements.
Periodicity refers to the repetition of similar properties after regular intervals when elements are arranged in increasing order of atomic number.
Early Attempts of Classification
1. Dobereiner’s Triads
Dobereiner grouped elements in sets of three called triads.
Condition:
Atomic mass of middle element ≈ (Atomic mass of first + third) / 2
Example:
Atomic mass relation:
Middle element ≈ (Element1 + Element3) / 2
2. Newlands’ Law of Octaves
Newlands arranged elements in increasing atomic mass.
Law: Every eighth element has properties similar to the first.
Representation:
Property(n) = Property(n + 7)
3. Mendeleev’s Periodic Table
Mendeleev arranged elements based on atomic mass and properties.
Mendeleev periodic law:
Properties ∝ Atomic mass
Limitations:
- Position of hydrogen uncertain
- Isotopes not explained
- Anomalies in atomic mass order
Modern Periodic Law
Modern periodic law states:
Properties ∝ Atomic number
This means properties of elements are periodic functions of their atomic numbers.
Modern Periodic Table
Features
- Elements arranged in increasing atomic number
- 7 periods (rows)
- 18 groups (columns)
Periods
Number of shells = period number
Example:
Number of shells = n
Groups
Elements in the same group have similar valence electrons.
Valence electrons determine chemical behavior.
Electronic Configuration and Periodicity
Electronic configuration follows the Aufbau principle.
Energy order rule:
Energy ∝ (n + l)
If (n + l) same → lower n has lower energy
Types of Elements
1. s-block Elements
Last electron enters s-orbital
General configuration:
ns¹ or ns²
2. p-block Elements
Last electron enters p-orbital
General configuration:
ns² np¹ to ns² np⁶
3. d-block Elements
Last electron enters d-orbital
General configuration:
(n−1)d¹ to (n−1)d¹⁰ ns⁰–²
4. f-block Elements
Last electron enters f-orbital
General configuration:
(n−2)f¹ to (n−2)f¹⁴
Periodic Properties
1. Atomic Radius
Distance from nucleus to outermost shell
Trend:
Across period → decreases
Down group → increases
Reason:
Effective nuclear charge increases across period
2. Ionic Radius
Cations: smaller than atom
Anions: larger than atom
3. Ionization Enthalpy
Energy required to remove electron
Equation:
X(g) → X⁺(g) + e⁻
Trend:
Across period → increases
Down group → decreases
4. Electron Gain Enthalpy
Energy released when electron is added
Equation:
X(g) + e⁻ → X⁻(g)
Trend:
Across period → more negative
Down group → less negative
5. Electronegativity
Ability to attract electrons
Trend:
Across period → increases
Down group → decreases
Effective Nuclear Charge
Effective nuclear charge formula:
Zeff = Z − S
Where:
Z = atomic number
S = shielding constant
Shielding Effect
Inner electrons reduce nuclear attraction on outer electrons
Shielding order:
s > p > d > f
Variation Summary
Across period:
Atomic radius ↓
Ionization energy ↑
Electronegativity ↑
Electron affinity ↑
Down group:
Atomic radius ↑
Ionization energy ↓
Electronegativity ↓
Anomalies in Periodic Trends
Some elements show irregular trends due to:
- Half-filled stability
- Fully-filled orbitals
Example concept:
Stability ∝ symmetry of electron distribution
Important Equations Summary
Energy relation:
E ∝ (n + l)
Effective nuclear charge:
Zeff = Z − S
Ionization process:
X(g) → X⁺(g) + e⁻
Electron gain:
X(g) + e⁻ → X⁻(g)
Conclusion
Classification of elements and periodicity help in understanding trends in properties of elements. These concepts are fundamental for predicting chemical behavior, bonding, and reactions.
Bonus: Exam Tips
- Learn trends with reasons, not just direction
- Focus on exceptions
- Practice electronic configuration questions
- Remember key equations