Periodic Table FULL CHAPTER | Class 11th Inorganic Chemistry | Arjuna NEET

Key Concepts

• Periodic Table Development: Historical progression from Lavoisier to Moseley, including Döbereiner's Triads, Newlands' Octaves, Lothar Meyer's Curve, and Mendeleev's Periodic Table.
• Modern Periodic Law: Elements' physical and chemical properties are periodic functions of their atomic numbers.
• Block Classification: Elements categorized into s, p, d, and f blocks based on the subshell of the last electron.
• Group and Period Determination: Methods to identify an element's group, period, and block based on electronic configuration and atomic number.
• Periodic Trends: Variations in atomic size, ionization energy, electron gain enthalpy, and electronegativity across periods and down groups.
• Lanthanide Contraction: The gradual decrease in atomic and ionic radii of the lanthanide elements with increasing atomic number.
• Transitional Contraction: The unexpected size decrease from Aluminum to Gallium.
• Isoelectronic Species: Ions or atoms with the same number of electrons.
• Ionization Energy: Energy required to remove an electron from a gaseous atom.
• Electron Gain Enthalpy: Enthalpy change when an electron is added to a gaseous atom.
• Electronegativity: Ability of an atom to attract shared electrons in a chemical bond.

Periodic Table Development

• Lavoisier: Classified elements as metals and non-metals, which was insufficient due to the discovery of metalloids.
• Döbereiner's Triads:
  • Grouped elements in triads with similar chemical properties and increasing atomic masses.
  • The atomic mass of the middle element was approximately the average of the first and third elements.
  • Examples: Chlorine, Bromine, Iodine; Calcium, Strontium, Barium; Lithium, Sodium, Potassium.
  • Limitations: Could not classify all known elements, and did not apply to d and f block elements.
• Newlands' Octaves:
  • Arranged elements in increasing atomic masses, noting that every eighth element had similar properties to the first.
  • Compared to musical notes (sa, re, ga, ma, pa, dha, ni, sa).
  • Limitations: Only applicable up to Calcium; assumed only 56 elements existed; failed after the discovery of noble gases.
• A.E.B. de Chancourtois:
  • Proposed a cylindrical periodic table (telluric helix).
  • Elements were arranged in increasing atomic weight on a cylinder.
  • Not widely accepted due to its complexity.
• Lothar Meyer's Curve:
  • Plotted atomic volume versus atomic weight.
  • Observed that elements with similar properties occupied similar positions on the curve.
  • Alkali metals occupied the peaks, alkaline earth metals the descending positions, and halogens the ascending positions.
• Mendeleev's Periodic Table:
  • Based on the periodic law: "The physical and chemical properties of elements are periodic functions of their atomic weights."
  • Arranged elements in horizontal rows (periods) and vertical columns (groups).
  • 7 periods and 8 groups.
  • Groups 1-7 were divided into subgroups A and B.
  • Group 8 had three rows of transition metals.
  • Predicted the existence and properties of undiscovered elements (Eka-aluminum, Eka-silicon, Eka-boron, Eka-manganese).
  • Corrected atomic masses of some elements.
  • Noble gases were added without disturbing the table.
  • Limitations: Position of hydrogen was uncertain; isotopes were not accounted for; anomalous pairs existed (Ar-K, Co-Ni, Te-I); dissimilar elements were grouped together.

Modern Periodic Table

• Moseley's Experiment:
  • Bombarded metals with high-speed electrons, producing X-rays.
  • Plotted the square root of X-ray frequency versus atomic number.
  • Observed a straight-line relationship, indicating atomic number was a more fundamental property than atomic weight.
• Modern Periodic Law: "The physical and chemical properties of elements are periodic functions of their atomic numbers."
• Characteristics:
  • 18 groups.
  • Elements in the same group have similar outer electronic configurations and chemical properties.
  • Noble gases are placed in Group 18.
  • Two series of 14 elements (Lanthanides and Actinides) are placed at the bottom.
  • Periods: 1 (2 elements), 2 & 3 (8 elements), 4 & 5 (18 elements), 6 & 7 (32 elements).
  • Period 1: Very short period.
  • Periods 2 & 3: Short periods.
  • Periods 4 & 5: Long periods.
  • Periods 6 & 7: Very long periods.
  • Period 7: Complete period.

IUPAC Nomenclature for Elements with Atomic Number > 100

• Use numerical roots for each digit of the atomic number:
  • 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn.
• Add "ium" to the end of the root.
• Symbol: First letter of each root.
• Examples:
  • 101: Unnilunium (Unu)
  • 118: Ununoctium (Uuo)
• Example Question: IUPAC name of element with atomic number 119: Ununennium (Uue).

Block Classification

• s-block:
  • Last electron enters the s-subshell.
  • General electronic configuration: ns1-2.
  • Groups 1 and 2.
  • Group 1: Alkali metals (ns1).
  • Group 2: Alkaline earth metals (ns2).
  • Exception: Helium (1s2) is placed in the p-block (Group 18).
• p-block:
  • Last electron enters the p-subshell.
  • General electronic configuration: ns2 np1-6.
  • Groups 13 to 18.
  • Group 13: Boron family (ns2 np1).
  • Group 14: Carbon family (ns2 np2).
  • Group 15: Nitrogen family or Pnictogens (ns2 np3).
  • Group 16: Chalcogens (ns2 np4).
  • Group 17: Halogens (ns2 np5).
  • Group 18: Noble gases (ns2 np6, except He: 1s2).
• d-block:
  • Last electron enters the d-subshell.
  • General electronic configuration: (n-1)d1-10 ns0-2.
  • Groups 3 to 12.
  • Transition elements or transition metals.
  • Four series: 3d (Period 4), 4d (Period 5), 5d (Period 6), 6d (Period 7).
  • Exception: Zn, Cd, Hg, and Cn are d-block elements but not transition elements because they have a completely filled d-subshell.
• f-block:
  • Last electron enters the f-subshell.
  • Inner transition elements.
  • Belong to Group 3 (IIIB).
  • Two series: Lanthanides (4f, Period 6) and Actinides (5f, Period 7).
  • Lanthanides: Cerium (Ce) to Lutetium (Lu).
  • Actinides: Thorium (Th) to Lawrencium (Lr).
  • General electronic configuration: (n-2)f1-14 (n-1)d0-1 ns2.

Determining Group, Period, and Block

1. Write the electronic configuration of the element.
2. Period number = Principal quantum number (n) of the valence shell.
3. Block: Determined by the subshell where the last electron enters.
4. Group:
   • Atomic number 104-118: Group number = Last two digits of atomic number.
   • s-block: Group number = Number of valence electrons.
   • p-block: Group number = 10 + Number of valence electrons.
   • d-block: Group number = Number of (n-1)d electrons + Number of ns electrons.
   • f-block: Group 3 (IIIB).

Atomic Size

• Definition: Distance from the center of the nucleus to the outermost electron.
• Units: Picometers (pm) or Angstroms (Å).
• Types: Covalent radius, van der Waals radius, metallic radius, ionic radius.
• Factors Affecting Atomic Size:
  • Number of shells: Increases atomic size.
  • Nuclear charge (Z): Higher Z decreases atomic size.
  • Screening effect (σ): Higher σ increases atomic size.
  • Effective nuclear charge (Zeff): Zeff = Z - σ.
• Trends:
  • Down the group: Atomic size increases.
  • Across the period: Atomic size decreases (with exceptions).
• d-block Contraction:
  • Initial decrease in size due to increasing nuclear charge.
  • Middle elements: Size remains almost constant due to balancing of attractive and repulsive forces.
  • Later elements: Size increases slightly due to increased electron-electron repulsion.
• Lanthanide Contraction:
  • Gradual decrease in atomic and ionic radii of lanthanides due to poor shielding by 4f electrons.
  • Results in similar sizes for 4d and 5d transition metals.
• Transitional Contraction:
  • Unexpected size decrease from Aluminum to Gallium.
  • Due to the filling of the d-orbitals before Gallium, which increases the effective nuclear charge.

Ionic Radius

• Cations are smaller than their parent atoms.
• Anions are larger than their parent atoms.
• For isoelectronic species:
  • Higher positive charge = Smaller size.
  • Higher negative charge = Larger size.
• Down the group: Ionic radii increase.

Ionization Energy (IE)

• Definition: Energy required to remove an electron from a gaseous atom.
• Always endothermic (positive value).
• Successive ionization energies increase (IE1 < IE2 < IE3).
• Factors Affecting IE:
  • Atomic size: Larger size = Lower IE.
  • Nuclear charge: Higher nuclear charge = Higher IE.
  • Penetration effect: Higher penetration = Higher IE (s > p > d > f).
  • Shielding effect: Higher shielding = Lower IE.
  • Electronic configuration: Half-filled and fully-filled configurations are more stable and have higher IE.
• Trends:
  • Down the group: IE decreases.
  • Across the period: IE increases (with exceptions).
• Applications:
  • Metallic and non-metallic character.
  • Reactivity of metals.
  • Number of valence electrons.
  • Stability of oxidation states.

Electron Gain Enthalpy (ΔegH)

• Definition: Enthalpy change when an electron is added to a gaseous atom.
• Can be exothermic (negative ΔegH) or endothermic (positive ΔegH).
• Electron affinity: Energy released when an electron is added (always exothermic).
• Factors Affecting ΔegH:
  • Atomic size: Smaller size = Higher ΔegH (more negative).
  • Nuclear charge: Higher nuclear charge = Higher ΔegH (more negative).
  • Electronic configuration: Stable configurations have lower ΔegH.
• Trends:
  • Down the group: ΔegH generally decreases (less negative).
  • Across the period: ΔegH generally increases (more negative).
• Exceptions:
  • Chlorine has a more negative ΔegH than Fluorine due to smaller size and higher interelectronic repulsion in Fluorine.
  • Oxygen has a less negative ΔegH than Sulphur.
  • Noble gases have positive ΔegH.

Electronegativity (EN)

• Definition: Ability of an atom to attract shared electrons in a chemical bond.
• Factors Affecting EN:
  • Atomic size: Smaller size = Higher EN.
  • Nuclear charge: Higher nuclear charge = Higher EN.
  • Hybridization: sp > sp2 > sp3.
• Trends:
  • Down the group: EN decreases.
  • Across the period: EN increases.
• Pauling scale: Fluorine (4.0) > Oxygen (3.5) > Nitrogen = Chlorine (3.0).
• Mulliken scale: EN = (IE + EA) / 2.
• Allred-Rochow scale: EN is proportional to the force between the nucleus and valence electrons.

Miscellaneous Points

• Bromine: Liquid non-metal at room temperature.
• Mercury: Liquid metal at room temperature.
• Gallium and Cesium: Melt when held in hand.
• Promethium: Only radioactive lanthanide.
• Osmium and Iridium: High density.
• Pnictogens: Group 15 elements.
• Chalcogens: Group 16 elements.
• Transuranic elements: Elements after Uranium in the actinide series.
• Diagonal relationships: Li-Mg, Be-Al, B-Si.

Conclusion

This comprehensive summary covers the key concepts, historical development, periodic trends, and important exceptions related to the classification of elements and periodicity in properties. It provides a detailed understanding of the periodic table and its applications, focusing on actionable insights and specific details.