Periodic Classification of Elements Revision Notes Class 10

Early Attempts at Classification

1. Dobereiner's Triads (1829)

• Law: Johann Wolfgang Döbereiner grouped elements with similar chemical properties into clusters of three (triads). He observed that when these elements were arranged in increasing order of atomic mass, the atomic mass of the middle element was roughly the arithmetic mean of the other two.
• Example: Lithium (Li, mass 6.9), Sodium (Na, mass 23), and Potassium (K, mass 39.1).

Mean Mass = (6.9 + 39.1) / 2 = 23 (Atomic Mass of Sodium)

• Limitation: He could identify only three triads from the elements known at that time (Total of 9 elements). It was not universally applicable to all known elements.

2. Newlands' Law of Octaves (1864)

• Law: John Newlands arranged the known elements in increasing order of their atomic masses. He found that every eighth element had properties similar to the first, comparing this to the seven notes of a musical scale (octaves).
• Limitations:
  • Applicable only up to Calcium (Ca). After Calcium, every eighth element did not possess similar properties.
  • He assumed only 56 elements existed in nature and no more would be discovered.
  • To fit elements into his table, he adjusted two elements into the same slot and placed completely dissimilar elements in the same note column (e.g., placing Cobalt and Nickel in the same column as Fluorine, Chlorine, and Bromine).

Mendeleev's Periodic Table (1869)

Mendeleev's Periodic Law: The chemical and physical properties of elements are a periodic function of their atomic masses.

Key Features:

1. Systematic Arrangement: Elements were organized into horizontal rows called Periods and vertical columns called Groups based on increasing atomic mass and chemical similarities (specifically formulas of their oxides and hydrides).
2. Gaps for Undiscovered Elements: Mendeleev boldly left gaps for undiscovered elements and predicted their properties by naming them using the prefix Eka (Sanskrit for 'one'). His predictions proved remarkably accurate when discovered later:
   • Eka-Boron → Scandium (Sc)
   • Eka-Aluminium → Gallium (Ga)
   • Eka-Silicon → Germanium (Ge)
3. Accommodation of Noble Gases: When noble gases (like Helium, Neon, Argon) were discovered later, they could be placed in a completely new group (Group 0) without disrupting the existing arrangement.

Limitations:

1. Anomalous Pairs: To maintain chemical similarity, elements with slightly higher atomic masses had to be placed before elements with lower atomic masses.
   Example: Cobalt (Co, atomic mass 58.9) was placed before Nickel (Ni, atomic mass 58.7).
2. Position of Isotopes: Isotopes of an element have different atomic masses but identical chemical properties. Placing them based strictly on atomic mass would require giving them separate slots, which was structurally impossible.
3. Position of Hydrogen: Hydrogen shows similarities to alkali metals (forms oxides, halides) as well as halogens (exists as a diatomic molecule, forms covalent compounds). Thus, its exact position was ambiguous.

The Modern Periodic Table (Moseley's Table)

In 1913, Henry Moseley demonstrated that the atomic number (the number of protons) of an element is a far more fundamental property for classification than its atomic mass.

Modern Periodic Law: The properties of elements are a periodic function of their atomic numbers.

Key Structural Features:

• Periods: There are 7 horizontal rows. The period number represents the total number of electron shells (n) present in the atoms of that period.
• Groups: There are 18 vertical columns. Elements in the same group possess the same number of valence electrons, giving them highly similar chemical properties.

Resolution of Mendeleev's Limitations:

• Isotopes: Because isotopes of the same element share the exact same atomic number, they naturally occupy the single same position in the table.
• Anomalous Pairs: When arranged by atomic number, anomalies like Cobalt (Z = 27) and Nickel (Z = 28) naturally fall into their correct logical order.

Summary Comparison of Classifications

Classification	Structural Basis	Principal Strength / Feature	Major Structural Limitation
Dobereiner's Triads	Atomic Mass	Identified a mathematical pattern linking mass to properties.	Only identified a total of 3 triads.
Newlands' Octaves	Atomic Mass	First to establish a repeating periodicity of properties.	Failed after Calcium; grouped dissimilar elements together.
Mendeleev's Table	Atomic Mass	Left gaps; accurately predicted undiscovered elements.	Couldn't assign fixed positions to isotopes or Hydrogen.
Modern Table	Atomic Number	Corrected ordering anomalies; highly systematic.	Unique, ambiguous position of Hydrogen remains.

Important Trends in the Modern Periodic Table

1. Valency

• Definition: The combining capacity of an atom, determined by its valence electrons.
• Across a Period (Left to Right): First increases from 1 to 4, and then decreases down to 0 (for representative groups).
• Down a Group (Top to Bottom): Remains constant because the number of valence electrons remains identical.

2. Atomic Size (Atomic Radius)

• Definition: The distance between the center of the nucleus and the outermost electron shell.
• Across a Period (Left to Right): Decreases. The nuclear charge increases (more protons), pulling the electron shells closer to the nucleus.
• Down a Group (Top to Bottom): Increases. New electron shells are added with each step down, expanding the physical size of the atom despite the increasing nuclear charge.

3. Metallic Character (Electropositivity)

• Definition: The tendency of an atom to lose electrons and form positive ions.
• Across a Period (Left to Right): Decreases. Increased nuclear charge binds valence electrons tightly, making them harder to lose.
• Down a Group (Top to Bottom): Increases. Atomic size grows and valence electrons move further from the nucleus, minimizing nuclear pull and making them easier to lose.

4. Non-Metallic Character (Electronegativity)

• Definition: The tendency of an atom to gain or attract shared electrons to form negative ions.
• Across a Period (Left to Right): Increases. The strong nuclear charge easily pulls in electrons.
• Down a Group (Top to Bottom): Decreases. Larger atomic size weakens the nucleus's ability to attract external electrons.

5. Nature of Oxides

• Across a Period (Left to Right): Oxides transition from Basic → Amphoteric → Acidic.
  • Basic Oxides: Formed by metals (e.g., Na₂O, MgO).
  • Amphoteric Oxides: Behave as both acidic and basic (e.g., Al₂O₃).
  • Acidic Oxides: Formed by non-metals (e.g., SO₂, CO₂).