Chemistry 1 - Periodic Properties of Elements Study Notes

Lecture - Week 7

Chemistry 1 - 1411ESC: Periodic Properties of Elements

Learning Outcomes
  • Understand how the structure of the Periodic Table is linked to the electron configuration of elements.
  • Understand how position in the Periodic Table can be used to predict trends in:
    • Atomic size
    • Ionisation energy
    • Electron affinity
    • Formation and stability of ions
    • Size of ions
Electron Configurations and the Structure of the Periodic Table
Orbital Diagrams for Multi-Electron Atoms
  • Orbital arrangement:
    • 0 4s 0 3s 0 2s 0 1s 0 4p 0 3p 0 2p

    • |-1 -1 -1 +1 +1 +1
      | 0 4f -3 -2 -1 +1 +2 +3
      | 0 +1 +2 4d 0 +1 +2
      | 3d -2
  • Energy levels categorization:
    • $n = 1$
    • $n = 2$
    • $n = 3$
    • $n = 4$
    • Angular momentum quantum numbers:
      • $l = 0$ (s), $l = 1$ (p), $l = 2$ (d), $l = 3$ (f)
Main Group Elements (s-block & p-block)
  • Period 1:
    • $H ext { : Atomic Number } 1, ext{ Atomic Mass } 1.00794$
  • Period 2:
    • $He ext { : Atomic Number } 2, ext{ Atomic Mass } 4.0026$
    • $Li ext { : Atomic Number } 3, ext{ Atomic Mass } 6.941$
    • $Be ext { : Atomic Number } 4, ext{ Atomic Mass } 9.01218$
    • $B ext { : Atomic Number } 5, ext{ Atomic Mass } 10.811$
    • $C ext { : Atomic Number } 6, ext{ Atomic Mass } 12.0107$
    • $N ext { : Atomic Number } 7, ext{ Atomic Mass } 14.0067$
    • $O ext { : Atomic Number } 8, ext{ Atomic Mass } 15.9994$
    • $F ext { : Atomic Number } 9, ext{ Atomic Mass } 18.9984$
    • $Ne ext { : Atomic Number } 10, ext{ Atomic Mass } 20.1797$
  • Period 3 includes:
    • $Na, Mg, Al, Si, P, S, Cl, Ar$
Period Structure
  • The periodic structure accommodates increasing energy levels and orbitals:
    • Period 1 has 2 elements
    • Periods 2 and 3 have 8 elements
    • Periods 4 and 5 have 18 elements
    • Period 6 has 32 elements (including Lanthanides)
Electron Configurations
  • As shells fill with electrons, the arrangement can be observed using:
    • Development of orbital diagrams for electronic distribution.
    • Example for Lead (Pb, Z = 82):
    • Full configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p²
    • Condensed configuration: [Xe] 6s² 4f¹⁴ 5d¹⁰ 6p²
Core and Valence Electrons
Definitions
  • Core Electrons: Electrons in inner shells that experience stronger nuclear attraction as atomic number increases.
  • Valence Electrons: Electrons in the outermost shell (e.g., for H, He, Li, Be, B, Ne, Na, Mg, Al, etc.).
Periodicity of Electron Configuration
  • Observed patterns in electron configurations across groups lead to similarities in chemical properties. The corresponding trends include:
    • Group 1 (Li, Na, K) has configuration $[He] 2s^1$, $[Ne] 3s^1$, $[Ar] 4s^1$.
    • Group 2 (Be, Mg, Ca) has $[He] 2s^2$, $[Ne] 3s^2$, $[Ar] 4s^2$.
Shielding/Screening Effect
  • Definition: Refers to the reduction of effective nuclear charge experienced by outer electrons due to repulsion from inner-shell electrons.
  • In He, 1s electrons feel a nuclear charge of +2, while in Ar, they feel +18.
Effective Nuclear Charge (Z_eff)
Formula
  • $Z_{eff} = Z - S$
    • Where: Z = actual nuclear charge, S = shielding (screening) constant > 0 but < Z.
  • Outer shell electrons feel less total nuclear charge due to shielding by inner electrons.
Trends in the Periodic Table
Periodicity
  • Refers to the variation in physical and chemical properties based on position in the Periodic Table due to valence shell configurations. Periodic trends:
    • Atomic Size: Increases down a group and decreases across a period.
    • Ionization Energy: Increases across a period (left to right) and decreases down a group.
    • Electron Affinity: Defined as the energy change when a neutral atom gains an electron.
Ionic Radius
  • The radius of an ion changes based on whether it is a cation or anion:
    • Cations (positive): Smaller than their parent atom.
    • Anions (negative): Larger than their parent atom.
Summary of Trends
Ionization Energy Trends
  • Example: $Li ext { : I} = 520 ext{ kJ/mol}$.
    • Ionization energy decreases down a group due to increased atomic radius and less nuclear attraction.
    • Ionization energy increases across a period due to increased nuclear charge.
Electron Affinity Trends
  • Reflect stability of newly formed anions, influenced by negative values of energy changes during atom-electron interactions. Large negative values indicate stable anion formation.