Bonds

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).

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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 $(\Delta EN)$; 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: $(\Delta EN)$ is between 0.4 and 1.7; electrons favor the atom with the higher $(\Delta EN)$.
• 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.

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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.

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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