Classification of Elements

1. Introduction

• Matter around us exists as elements, compounds, and mixtures.
• There are 118 elements, each with unique properties.
• To study them systematically, scientists classified elements based on similarities in physical and chemical properties.

Definition of Classification of Elements:
Arrangement of elements in a systematic manner according to their similar properties.

Purpose of Classification:

• Elements have different physical and chemical properties.
• Some elements show similar properties.
• Helps in studying, predicting, and understanding the behavior of elements easily.

2. Historical Developments

(a) Döbereiner’s Triads (1829)

• Elements grouped in triads of three with similar properties.
• Atomic mass of the middle element ≈ average of the other two.
• Example: Li (7), Na (23), K (39) → 23 ≈ (7 + 39)/2
• Limitation: Worked only for a few triads.

(b) Newlands’ Law of Octaves (1864)

• Arranged elements by increasing atomic mass.
• Every 8th element had similar properties to the first.

Limitations:

1. Worked only up to calcium.
2. Failed for heavier elements.
3. Transition metals did not fit.

(c) Lothar Meyer’s Curve (1869)

• Plotted atomic volume vs atomic mass.
• Peaks corresponded to alkali metals.
• Showed periodicity of properties.

(d) Mendeleev’s Periodic Table (1869)

Arranged elements by increasing atomic mass.

Achievements:

• Left gaps for undiscovered elements.
• Predicted properties of Ga, Ge, Sc accurately.
• Grouped elements correctly.

Limitations:

• Atomic mass order sometimes failed (Co & Ni).
• Position of hydrogen unclear.
• Isotopes did not fit.

3. Modern Periodic Law

"Properties of elements are a periodic function of their atomic number (Z)."

Atomic number is the real basis of periodicity.

4. Modern Periodic Table Features

Periods (Horizontal Rows): 7

Period number = number of electron shells (n)

Groups (Vertical Columns): 18

Group number = number of valence electrons (except transition metals)

Blocks:

• s-block: Groups 1–2 → Highly reactive metals, form basic oxides
• p-block: Groups 13–18 → Metals, non-metals, metalloids, halogens, noble gases
• d-block: Groups 3–12 → Transition metals, colored compounds, variable oxidation states
• f-block: Lanthanides & Actinides → Inner transition elements, mostly radioactive (actinides)

5. Key Differences

Metals vs Non-metals

s-block vs p-block

6. Periodic Properties

(A) Atomic Radius

Across a period: Decreases (more nuclear charge pulls electrons closer)

Down a group: Increases (more electron shells added)

(B) Ionic Radius

Cations (+): Smaller (lose electrons)

Anions (−): Larger (gain electrons)

(C) Ionization Energy (IE)

1. Energy required to remove an electron.
2. Across period → increases
3. Down group → decreases
4. Low IE → Alkali metals; High IE → Noble gases

(D) Electron Affinity (EA)

1. Energy released when an atom gains an electron.
2. Across period → increases
3. Down group → decreases
4. Highest EA → Chlorine

(E) Electronegativity (EN)

1. Tendency to attract shared electrons.
2. Across period → increases
3. Down group → decreases
4. Most EN element → Fluorine

(F) Metallic Character

1. Decreases from left → right
2. Increases from top → bottom
3. Opposite trend to non-metallic character

7. Special Cases

1. Hydrogen: Behaves like alkali metals and halogens → special position
2. Helium: Noble gas, though s-block
3. Lanthanide contraction: Affects d-block sizes
4. Co & Ni anomaly: Fixed in modern table

8. Valency

1. Valency: Number of electrons an atom gains, loses, or shares.
2. s-block: Valency = group number
3. p-block: Varies (commonly 3, 4, 5, etc.)

9. Why Periodicity Occurs?

Because electronic configurations repeat at regular intervals.

Example:

Group 1: ns¹

Group 2: ns²

Group 17: ns² np⁵