Ionic Bonds

  • Nature of Bond: Electrons are transferred, not shared.
  • Participants: Typically between one metal and one nonmetal, resulting in oppositely charged ions.
  • Bond Strength: Electrostatic interactions hold the ions together.

Ionic Solids

  • Solid at room temperature due to strong ionic bonds.
  • High boiling points.
  • Highly ordered crystal lattice structure:
    • Not malleable due to lack of give.
  • Solubility: At least slightly soluble in water.
  • Conductivity:
    • Cannot conduct electricity in solid form.
    • Conducts electricity as a liquid or in aqueous solution (ions are free to move).

Covalent Bonds

  • Nature of Bond: Electrons are shared.
  • Participants: Generally occurs between two atoms with high ionization energies (commonly two nonmetals).

Non-Polar Covalent Bonds

  • Definition: Atoms have less than a 0.4 ; electrons are shared equally.
  • Properties:
    • Use London dispersion forces (LDF).
    • Low melting and boiling points.
    • Often gases at room temperature due to weak intermolecular forces.
    • Larger molecules (e.g., wax) can be solid at room temperature.
    • Do not conduct electricity (no ions).

Polar Covalent Bonds

  • Definition: is between 0.4 and 1.7; electrons favor the atom with the higher .
  • Properties:
    • Stronger intermolecular forces than non-polar covalent (dipole-dipole forces; sometimes hydrogen bonding).
    • Dissolve in polar solvents if hydrogen bonding is present (e.g., sugar in water).

Coordinate Covalent Bonds

  • Definition: Both shared electrons come from the same atom.

Metallic Bonds

  • Nature of Bond: Forms between two metals.
  • Participants: Atoms with low ionization energies and low electronegativities.
  • Structure:
    • Positive nuclei in a lattice surrounded by delocalized electrons.
  • Electron Behavior: Electrons are delocalized, allowing for unique metallic properties.
  • Properties:
    • Malleable and ductile due to the unrestricted movement of delocalized electrons.
    • Good conductors of heat and electricity due to delocalized electrons.
    • Shiny: Light excites valence electrons, which absorb energy, oscillate, and reemit it.
    • Exhibit the photoelectric effect.

Network Solids

  • Definition: Macromolecules with giant covalently bonded structures in 1D, 2D, or 3D arrays.
  • Bond Strength: Covalent bonds are stronger than ionic bonds.
  • Properties: Depend on dimensional structure.

Allotropes

  • Definition: Different physical forms of the same element.

Dimensional Structures

  • 3D Network Solids:

    • Extremely strong bonds result in very high melting and boiling points.
    • Extremely hard and not soluble.
    • Examples: Diamond, quartz, silicon carbide.
  • 2D Network Solids:

    • Layers of 2D sheets with high melting and boiling points but softer structures.
    • Conduct electricity slightly due to delocalized double bonds.
    • Example: Graphite.
  • 1D Network Solids:

    • Form fibers with high melting and boiling points.
    • Examples: Asbestos.

Examples of Network Solids

  • Silicon compounds, mica, asbestos.

Atomic structures

  • bonds moving is only electricity conductive if the bonds move around OUTSIDE of the electrons
  • when calculating formal charges, count bonds as one, when calculating if it’s valance shell is full, count as two