Periodic Properties of the Elements

Periodic Properties of the Elements

3.1 Aluminum: Low-Density Atom Result in Low Density Metal

  • Density and Atomic Properties
    • Low-density metals, such as aluminum, contrast with higher density elements like iron.
    • Density increases as you move down a group: the mass-to-volume ratio increases.
    • Density is a periodic property dependent on atomic structure and periodic table position.

3.2 The Periodic Law and the Periodic Table

  • Mendeleev's Discovery

    • Dmitri Mendeleev created the first periodic table in 1869, organizing elements by atomic mass.
    • Introduced gaps for predicted undiscovered elements, e.g., eka-silicon (predicted Germanium).

  • Structure of Modern Periodic Table

    • Arranged by increasing atomic number (not mass).
    • Rows = periods; Columns = groups/families.
    • Elements in the same group have similar properties due to similar electron configurations.

  • Classification of Elements

    • Metals, Nonmetals, Metalloids.
    • Main-group elements are predictable based on their position; transition metals are less predictable.

3.3 Electron Configuration: How Electrons Occupy Orbitals

  • Quantum-Mechanical Behavior of Electrons

    • Electrons occupy orbitals in a quantized manner around the nucleus.
    • Ground state: lowest energy configuration; Excited state: higher energy levels.

  • Principal Numbers and Subshells

    • Principal quantum number (n) indicates energy level; each level can contain different subshells (s, p, d, f).
    • Electron capacity for shells: n=1 (2), n=2 (8), n=3 (18), n=4 (32).

  • Filling Rules

    • Aufbau principle: electrons fill lower energy orbitals first.
    • Pauli exclusion principle: no two electrons can have the same set of four quantum numbers.
    • Hund's rule: electrons fill degenerate orbitals singly before pairing.

3.4 Electron Configurations, Valence Electrons, and the Periodic Table

  • Valence Electrons

    • Element groups dictate the number of valence electrons.
    • Main group electron configurations help predict reactivity and ionic charges.

  • Reactivity of Noble Gases

    • Noble gases are stable as their valence shells are fully filled (low energy).

3.5 Electron Configurations and Elemental Properties

  • Influence of Electron Configuration on Reactivity
    • Elements near noble gases tend to gain or lose electrons readily (high energy).
    • Specific configurations can predict ion charges, e.g., Alkali metals as +1, Halogens as -1.

3.6 Periodic Trends in the Size of Atoms and Effective Nuclear Charge

  • Atomic Size Trends

    • Atomic radius increases down a group (more electron shells), decreases across a period (effective nuclear charge increases).
    • Effective nuclear charge (Zeff) can be calculated: Zeff = Z (atomic number) - S (shielding constant).

  • Impact of Shielding

    • Inner electrons shield valence electrons from the full effect of nuclear charge, affecting their energy levels and ionization energy.

3.7 Ions: Electron Configurations, Magnetic Properties, Ionic Radii, and Ionization Energy

  • Cations and Anions

    • Cations are smaller than their neutral atoms due to electron loss; anions are larger due to electron gain.
    • Ionization energy is the energy needed to remove an electron; it generally decreases down a group and increases across a period.

  • Successive Ionization Energies

    • Provided for various elements, indicating difficulties in removing electrons after initial removal.

3.8 Electron Affinities and Metallic Character

  • Electron Affinity Trends

    • Energy released when adding an electron; generally, increases across a period, particularly for halogens.
    • Metals tend to lose electrons (oxidized) while nonmetals tend to gain electrons (reduced).

  • Metallic Character Trends

    • Increases down a group and decreases across a period. Metals are typically malleable, ductile, and good conductors of heat and electricity.

3.9 Examples of Periodic Chemical Behavior

  • Specific Groups
    • Alkali Metals (Group 1A): high reactivity due to one valence electron.
    • Halogens (Group 7A): high reactivity due to seven valence electrons (want to gain one).
    • Noble Gases (Group 8A): unreactive due to full outer electron shells.