Classification of Elements and Periodicity in Properties

• Historical Development

  • Dobereiner's Triads: Atomic mass of middle element = average of other two
  • Newlands' Law of Octaves: Every 8th element shows similar properties
  • Lothar Meyer's Curve: Plotted Atomic Volume vs Atomic Mass
  • Mendeleev's Periodic Law: Properties are periodic functions of Atomic Mass
    • Merits: Left gaps for undiscovered elements (Eka-Aluminium, Eka-Silicon)
    • Demerits: Position of Hydrogen, Isotopes, and anomalous pairs

• Modern Periodic Law & Table

  • Moseley's Experiment: X-ray spectra proved physical/chemical properties depend on Atomic Number (Z)
  • Modern Periodic Law: Properties are periodic functions of Atomic Number
  • Long Form Structure:
    • 7 Periods (Horizontal rows)
    • 18 Groups (Vertical columns)

• Nomenclature for Elements with Z > 100

  • Roots: 0=nil, 1=un, 2=bi, 3=tri, 4=quad, 5=pent, 6=hex, 7=sept, 8=oct, 9=enn
  • Suffix: -ium (e.g., Z = 101 $\rightarrow$ Unnilunium)

• Classification by Electronic Configuration

  • s-Block: Groups 1 & 2 ($ns^{1-2}$) | Reactive metals, strong reducing agents
  • p-Block: Groups 13 to 18 ($ns^2 np^{1-6}$) | Metals, non-metals, metalloids
  • Representative Elements: s-Block + p-Block (excluding Group 18)
  • d-Block: Groups 3 to 12 ($(n-1)d^{1-10} ns^{1-2}$) | Transition Elements, colored compounds, variable oxidation states
  • f-Block: Inner Transition Elements ($(n-2)f^{1-14} (n-1)d^{0-1} ns^2$)
    • Lanthanoids (4f series: $Ce$ to $Lu$)
    • Actinoids (5f series: $Th$ to $Lr$)

• Periodic Trends in Physical Properties

  • Atomic Radius:
    • Types: Covalent, Metallic, van der Waals radii
    • Across Period: Decreases (Increase in Effective Nuclear Charge, $Z_{eff}$)
    • Down Group: Increases (Addition of new energy shells)
  • Ionic Radius:
    • Cation < Parent Atom (Loss of electron, higher $Z_{eff}$)
    • Anion > Parent Atom (Gain of electron, higher repulsion)
    • Isoelectronic Species: Size decreases with increasing Nuclear Charge
  • Ionization Enthalpy ($\Delta_i H$):
    • Energy required to remove the most loosely bound electron from a isolated gaseous atom
    • Across Period: Increases
    • Down Group: Decreases
    • Factors: Atomic size, Shielding Effect, Penetration Effect, Half/Fully-filled subshells
  • Electron Gain Enthalpy ($\Delta_{eg} H$):
    • Energy change when an electron is added to a neutral gaseous atom
    • Across Period: Becomes more negative
    • Down Group: Becomes less negative
    • Exception: Chlorine has more negative $\Delta_{eg} H$ than Fluorine (due to inter-electronic repulsions in 2p orbital)
  • Electronegativity:
    • Tendency of an atom to attract shared pair of electrons
    • Across Period: Increases
    • Down Group: Decreases
    • Scale: Pauling Scale (Fluorine = 4.0, highest value)

• Periodic Trends in Chemical Properties

  • Valency / Oxidation State:
    • Valence electrons (1 to 4) or ($8 - \text{valence electrons}$)
    • Across Period: Increases from 1 to 4, then decreases to 0
    • Down Group: Remains same
  • Anomalous Properties of Second Period Elements:
    • Li, Be, B, C, N, O, F differ from their respective groups
    • Reasons: Small atomic size, high Electronegativity, high $\Delta_i H$, absence of d-orbitals
  • Diagonal Relationship:
    • Similarities between elements diagonally placed in Periods 2 and 3
    • Pairs: Li-Mg, Be-Al, B-Si
    • Reason: Similar charge/radius ratio
  • Chemical Reactivity & Oxides:
    • Basic Oxides: Formed by metallic elements on extreme left (e.g., $Na_2O$)
    • Acidic Oxides: Formed by non-metallic elements on extreme right (e.g., $Cl_2O_7$)
    • Amphoteric Oxides: Show both acidic and basic character (e.g., $Al_2O_3$, $BeO$)
    • Neutral Oxides: $CO$, $NO$, $N_2O$