General Chemistry - Periodic Table and Properties

Introduction to the Periodic Table

  • The periodic table is divided into parts that represent orbitals related to energy levels.
  • The periods are aligned horizontally and correspond to defined energy levels.
  • Groups, running vertically, represent an increase in valence electrons.
  • Chemical and physical characteristics of an element are significantly determined by its valence electrons.

Types of Elements

  • Metals: Found below a staircase on the periodic table.

    • Characteristics:
    • Luster (shine)
    • Malleability (ability to be shaped)
    • Durability
    • Electro positivity (tendency to lose electrons)
    • Valence electrons in metals are loosely held, making it easier to ionize them (process of losing electrons).

  • Nonmetals: Located above the staircase.

    • Characteristics:
    • Generally brittle
    • Lack luster
    • High ionization energies, meaning valence electrons are tightly held.
    • Capable of very high electronegativities and electron affinity.
    • Smaller atomic radii compared to metals.

  • Metalloids: Along the staircase, exhibiting properties of both metals and nonmetals.

    • Common metalloids include: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium, and Astatine.

Chemical Trends in the Periodic Table

  • As one moves across a period from left to right:
    • The number of protons increases.
    • The number of electrons increases, while the energy level remains the same.
  • Example: Comparison of Fluorine (atomic number 9) and Chlorine (atomic number 17).
    • Fluorine has 2 energy levels, Chlorine has 3. Therefore, Chlorine is larger, as it has an additional energy level.

Effective Nuclear Charge (Z_eff)

  • Definition: The net positive charge experienced by valence electrons due to the shielding effect from inner core electrons.
  • Calculation example for Fluorine:
    • 9 protons (charge) - 2 core electrons = 7 effective nuclear charge.
  • The effective nuclear charge increases as you move across a period due to the increased number of protons while the shielding effect remains constant.

Atomic and Ionic Radius

Atomic Radius:

  • The size of an atom is determined by the number of energy levels and the effective nuclear charge.
  • Atomic radius increases down a group due to the addition of energy levels.
  • Atomic radius decreases across a period from left to right due to increased nuclear charge pulling electrons closer.
    • Example: Francium is the largest, while Helium (noble gas) is much smaller despite having fewer protons.

Ionic Radius:

  • The ionic radius is not the same as the atomic radius. When an atom ionizes:

    • Cations (positive ions): Formed by losing electrons; ionic radius is typically smaller than atomic radius—example Lithium.
    • Anions (negative ions): Formed by gaining electrons; ionic radius is typically larger than atomic radius.

  • Comparison of Sodium (Na) and Magnesium (Mg):

    • Na loses 1 electron (becomes Na+ with configuration Ne) vs. Mg loses 2 electrons (becomes Mg2+ with configuration Ne). Despite being isoelectronic (same electron configuration), Mg2+ is smaller due to a higher nuclear charge pulling the electrons in more closely.

Ionization Energy

  • Definition: The energy required to remove an electron from an atom.
  • Trends:
    • Ionization energy increases up a group and across a period due to:
    • Increased nuclear charge (more protons) makes it harder to remove electrons.
    • Less energy levels means more effective shielding.
  • Example: Fluorine has the highest ionization energy due to high electronegativity compared to other elements.

Electron Affinity

  • Definition: The amount of energy released when an atom gains an electron. Generally, the process is exothermic.
  • Higher electron affinity is found in elements that want to gain electrons (e.g., nonmetals like halogens).
  • Trends: Electron affinity increases up a group and from left to right across a period.

Electronegativity

  • Definition: The ability of an atom to attract electrons within a chemical bond.
  • Trends: Electronegativity increases up a period and across a group.
  • Reference bond for electronegativity scaled from 0 (non-polar) to 4 (very polar).
  • Example values: Carbon-2.5, Oxygen-3.5, and Fluorine-4.0. The difference determines bond character (covalent vs ionic).

Summary of Groups

  • Alkali Metals (Group 1):

    • Highly reactive with water, creating hydroxides and hydrogen gas.

  • Alkaline Earth Metals (Group 2):

    • Also reactive but less than alkali metals; react with water and acids.

  • Chalcogens (Group 16):

    • Important for biological processes; reactive that can create oxidative stress.

  • Halogens (Group 17):

    • Very reactive; exist in diatomic forms; form salts when they react with metals.

  • Noble Gases (Group 18):

    • Inert under most conditions; low electronegativity; do not form stable compounds easily.

  • Transition Metals:

    • Have a wide range of oxidation states, exhibit color in compounds due to d-orbital electron transitions. They can form complex ions and various hydrated states, often causing vivid colors due to jump and release energy in visible light ranges.

Conclusion

  • The information about periodic trends is critical for understanding elemental behavior, bonding characteristics, and only emphasizes the trends observed across various groups and periods.
  • Important principles regarding ionization energy, ionic and atomic radius, effective nuclear charge, and electronegativity establish the foundational knowledge of chemistry.