Ch 9

Chapter 9: Electron Configuration and the Periodic Table

Development of The Periodic Table
  • John Newlands: Proposed that elements arranged by increasing atomic mass show periodicity. Elements in the same group appear in horizontal rows.
  • Dmitri Mendeleev & Lothar Meyer: Independently grouped elements similarly, leading to the early development of the periodic table. Mendeleev's table emphasized valency.
Mendeleev's First Periodic Table (1869)
  • Example elements and their atomic masses:
    • Ti = 50, Zr = 90, V = 51, Nb = 94, Ta = 182, Cr = 62, Mo = 96, W = 186, etc.
Modern Periodic Table

displays organized groups:

  • Group classifications such as IA, IIA (Alkali Metals), IIIA (Aluminum group), IV-VI (Nonmetals), and VIIA (Halogens).
Electron Configuration of Ions
  • Group I:
    • Na: $10[Ne] 3s^1$
    • Cs: $54 [Xe] 6s^1$
  • Group II:
    • Mg: $10[Ne] 3s^2$
    • Ba: $54 [Xe] 6s^2$
  • Group III:
    • Al: $10[Ne] 3s^2 3p^1$
Electron Configuration of Anions
  • Group 5A: N: $2[He] 2s^2 2p^3$
  • Group 6A: S: $10[Ne] 3s^2 3p^4$
  • Group 7A: Cl: $10[Ne] 3s^2 3p^5$
Periodic Trends
  • Atomic Size and Ionization Energy:
    • Atomic size decreases across a period and increases down a group.
    • Ionization energy generally increases across a period and decreases down a group.
Effective Nuclear Charge
  • Defined by: Zeff=Z−sZ_{eff} = Z - s
    • Z: Atomic number
    • s: Screening constant (usually equal to the number of inner electrons); can be between 0 and Z.
Sizes of Atoms (Atomic Radius)
  • Bonding atomic radius: half the distance between covalently bonded nuclei.
  • Trends in atomic radius:
    • Decreases left to right due to increasing $Z_{eff}$.
    • Increases from top to bottom due to increasing principal quantum number (n).
Sizes of Ions
  • Cations: Smaller than parent atoms (outer electron removal reduces repulsion).
  • Anions: Larger than parent atoms (electron addition increases repulsion).
  • Isoelectronic series: Ions with the same number of electrons but increasing nuclear charge will have decreasing size.
Ionization Energy
  • Energy needed to remove an electron from an atom.
    • First Ionization Energy: Energy to remove the first electron; subsequent removals require exponentially more energy.
    • Measured in kJ/molkJ/mol.
    • Trends: more energy needed across a period, less down a group.
Electron Affinity
  • Energy change when adding an electron to an atom. Example for Chlorine:
    • Cl+e−→Cl−Cl + e^- → Cl^-
    • EA=−extΔH=349extkJ/molEA = - ext{ΔH} = 349 ext{ kJ/mol}
  • General trend: more exothermic left to right across a period.
Increasing Metallic Character
  • Properties of Metals, Nonmetals, and Metalloids:
    • Transition from metals (shiny, malleable, conductive) to nonmetals (brittle, poor conductors).
Alkali Metals
  • Soft, low-density metallic solids.
  • React severely with water, forming strong alkaline solutions.
  • Most reactive group of metals.
Alkaline Earth Metals
  • Higher densities and melting points than alkali metals; less reactive compared to them.
  • Reactive with water, with reactivity increasing down the group.
Group 6A: Oxygen Family
  • Oxygen: $O2$ and ozone ($O3$); tends to be strong oxidizing agent.
  • Sulfur: Most stable allotrope is S8 (ring structure).
Group VIIA: Halogens
  • Highly reactive nonmetals; known for distinct properties such as large negative electron affinities.
    • Form metal halides when reacting with metals; compounds have distinct structures (e.g. Cl, Br, I).
Group VIIIA: Noble Gases
  • Very low reactivity; monatomic gases (inert under standard conditions).
  • Has very high ionization energies and positive electron affinities.