Historical Development of Periodic Table

The periodic table evolved through the contributions of several scientists who attempted to systematically arrange elements based on their properties.

Dobereiner's Triads (1829)

Johann Wolfgang Dobereiner grouped elements in sets of three with similar properties, showing that the atomic weight of the middle element was approximately the average of the other two.

Newlands' Law of Octaves (1864)

John Newlands observed that when elements were arranged by atomic weight, every eighth element had properties similar to the first, like musical octaves.

Mendeleev's Periodic Table (1869)

Dmitri Mendeleev created the first successful periodic table based on the periodic law:

Key Features of Mendeleev's Table

• Elements arranged in order of increasing atomic weights
• Vertical columns called groups with similar properties
• Horizontal rows called periods
• Left gaps for undiscovered elements (predicted properties accurately)
• Corrected atomic weights of some elements based on periodic trends

Merits of Mendeleev's Table

• Systematic classification of known elements
• Successful prediction of new elements (Ga, Ge, Sc)
• Correction of atomic weights (Be, U, etc.)
• Study of element properties became easier

Limitations of Mendeleev's Table

• No fixed position for hydrogen (resembles both alkali metals and halogens)
• Some elements with higher atomic weights placed before lower ones (Ar/K, Co/Ni, Te/I)
• No explanation for periodicity
• Lanthanides and actinides couldn't be accommodated properly
• No distinction between metals and non-metals

Modern Periodic Table

Henry Moseley's work on atomic numbers led to the modern periodic law:

Modern Periodic Table with Group and Period Classification

Structure of Modern Periodic Table

• 18 Groups (vertical columns) - Elements with same valence electrons
• 7 Periods (horizontal rows) - Elements with same number of shells
• 4 Blocks based on electron configuration:
  • s-block (Groups 1-2)
  • p-block (Groups 13-18)
  • d-block (Groups 3-12)
  • f-block (Lanthanides & Actinides)

Nomenclature of Groups

📊 Class 11: Essential Periodic Trends

Shielding Effect

The reduction in effective nuclear charge on outer electrons due to inner electrons is called shielding effect. It increases down a group as more inner shells are added.

Electron Gain Enthalpy

The energy change when an electron is added to neutral gaseous atom:

Trends: Becomes more negative across period (except Noble gases), Becomes less negative down group

Metallic Character

• Across period: Decreases (elements become less metallic and more non-metallic)
• Down group: Increases (easier electron loss due to larger atomic size)

Blocks in Periodic Table

S-Block Elements (Groups 1-2)

Elements where the last electron enters the s-orbital. Includes alkali metals (Group 1) and alkaline earth metals (Group 2).

• Outer configuration: ns^1-2
• Highly reactive metals (except H)
• Low ionization energies
• Form basic oxides and ionic compounds
• Strong reducing agents

P-Block Elements (Groups 13-18)

Elements where the last electron enters the p-orbital. Includes metals, metalloids, and non-metals.

• Outer configuration: ns²np^1-6
• Show variable oxidation states
• Form covalent compounds (smaller atoms) and ionic compounds (larger atoms)
• Include all non-metals and metalloids

D-Block Elements (Transition Metals, Groups 3-12)

Elements where the last electron enters the d-orbital of penultimate shell.

• Outer configuration: (n-1)d^1-10ns^0-2
• Show variable oxidation states
• Form colored compounds
• Act as catalysts
• Paramagnetic nature (unpaired d-electrons)
• Form complex compounds

F-Block Elements (Inner Transition Metals)

Elements where the last electron enters the f-orbital of antepenultimate shell.

• Outer configuration: (n-2)f^1-14(n-1)d^0-1ns²
• Lanthanides (4f series) and Actinides (5f series)
• Show +3 oxidation state predominantly
• Radioactive elements (especially actinides)
• Lanthanide contraction causes similar sizes of post-lanthanide elements

Anomalies and Special Cases

Position of Hydrogen

Hydrogen shows properties of both alkali metals (loses electron to form H⁺) and halogens (gains electron to form H^-), making its position ambiguous.

Anomalous Properties of Second Period Elements

Elements of second period (Li to F) show different behavior compared to their congeners due to:

• Small size
• High electronegativity
• Absence of d-orbitals
• Strong tendency to form pπ-pπ multiple bonds

Lanthanide Contraction

The steady decrease in atomic and ionic radii of lanthanides (La to Lu) due to poor shielding by 4f electrons causes:

• Similar sizes of 4d and 5d transition series elements
• Difficulty in separation of lanthanides
• Increased density and hardness of 5d series elements

Applications of Periodic Properties

In Biological Systems

• Na/K pump: Size difference crucial for nerve impulse transmission
• Mg in chlorophyll: Perfect size to coordinate with porphyrin ring
• Fe in hemoglobin: Variable oxidation states allow oxygen transport
• Ca in bones: Forms stable compounds with phosphates
• Trace elements: Many d-block elements are essential in small quantities

In Industry and Technology

• Aluminum production: Low reduction potential of Al makes electrolysis necessary
• Catalysts: Transition metals used due to variable oxidation states
• Semiconductors: p-block elements (Si, Ge) with intermediate conductivity
• Batteries: Based on redox reactions of alkali and transition metals
• Alloys: Combining metals with different properties to create superior materials

NCERT Important Questions & Mnemonics

Important NCERT Questions

Interactive Periodic Table

Naming of Elements with Atomic Number >100

For elements with atomic number >100, the digits are written in root words followed by 'ium'.