Atomic Structure and Chemical Bonding Principles

Atomic Structure Fundamentals

Ionization Energy and Valence Electrons

• Ionization Energy (IE): Energy required to remove an electron from a gaseous atom in its ground state.
• Valence Electrons: The electrons located in the outermost energy level.

Quantum Numbers

These numbers describe the probable location and behavior of an electron within an atom:

1. Principal Quantum Number (n): Indicates the energy level where the electron is found.
2. Azimuthal Quantum Number (l): Determines the orbital shape. Values range from 0 to (n-1).
   • l = 0 (s orbital)
   • l = 1 (p orbital)
   • l = 2 (d orbital)
   • l = 3 (f orbital)
3. Magnetic Quantum Number (m_l): Determines the orientation of the orbital in space. Values range from -l to +l.
4. Spin Quantum Number (m_s): Determines the electron spin.

Periodic Table Structure

Elements are classified based on their valence electron configuration:

• Representative Elements (Group A): Elements filling s and p orbitals.
• Transition Elements (Group B): Elements filling d and f orbitals.

Transition Metal Electron Configurations (Group B)

Notable exceptions to standard filling order include Chromium (Cr, 24), Copper (Cu, 29), Molybdenum (Mo, 42), and Silver (Ag, 47).

Group definitions based on valence shell configuration:

• IB: ns¹ (n-1)d¹⁰
• IIB: ns² (n-1)d¹⁰
• IIIB: ns² (n-1)d¹
• IVB: ns² (n-1)d²
• VB: ns² (n-1)d³
• VIB: ns² (n-1)d⁴
• VIIB: ns² (n-1)d⁵
• VIII (Groups 8, 9, 10):
  • ns² (n-1)d⁶
  • ns² (n-1)d⁷
  • ns² (n-1)d⁸

Key Periodic Properties

Definitions of Core Properties

• Shielding Effect (Screening Effect): The decrease in the attractive force of the nucleus on the outermost electrons due due to inner electron shells.
• Effective Nuclear Charge (Z_eff): The net positive charge experienced by an electron in the outermost shell.
• Ionization Potential (IP): Energy required to remove an electron from a neutral gaseous atom, producing a cation.
• Electron Affinity (EA): Energy released when a neutral gaseous atom captures an electron, becoming an anion.
• Electronegativity (EN): The capability of an atom to attract electrons toward itself when forming a chemical bond.

Periodic Trends

• Shielding Effect: Increases downward (down a group).
• Effective Nuclear Charge (Z_eff):
  • Increases to the right (across a period).
  • Is relatively constant down a group.
• Electronegativity, Electron Affinity, Ionization Potential:
  • Increase to the right (across a period).
  • Decrease downward (down a group).
• Atomic Radius (Size): Increases downward and decreases to the right.

Metallic Character Trends

Metallic character decreases across a period (to the right) and increases down a group.

• The most metallic elements are found on the left and bottom of the table.
• The least metallic (most non-metallic) elements are found on the right and top.

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in molecules or ions in crystals.

Ionic Bonding

Involves the transfer of valence electrons, typically between a metal and a non-metal.

• Metal atoms lose electrons to form ions, resulting in a totally empty external energy level.
• Non-metallic atoms gain electrons to form ions, resulting in totally occupied external energy levels.

Note on Electron Tendency:

• Metals tend to lose electrons.
• Non-metals tend to gain electrons.

Covalent Bonding

Involves the sharing of valence electrons between atoms.

• Nonpolar Covalent Bond (Apolar): Equal sharing of electrons between two atoms of the same element.
• Polar Covalent Bond: Unequal sharing of electrons between atoms of different electronegativities.

Bond Type Determination

When the difference in electronegativity is greater than 1.7, the bond is considered primarily Ionic.

Summary of Electron Behavior:

• Ionic: Electrons are transferred.
• Covalent: Electrons are shared.