Bonding

• Ionic bonding


  Charged ions held together by strong electrostatic attractions
• Formation of ionic bond

  1. Atom gives up electron to another atom

  2. Oppositely charged ions are attracted to each other
• Molecular ions


  Hydroxide (OH-), Nitrate (NO3-), Ammonium (NH4+), Sulfate (SO4 2-), Carbonate (CO3 2-)
• Determining formula of ionic compound

  1. Write the two ions

  2. Swap the charges over

  3. Drop the charges to get the subscripts

  4. Simplify the formula
• Giant ionic structures


  • Cubic arrangement of ions
  • High melting points due to strong electrostatic forces
• Covalent bonding


  Sharing of electrons between atoms to achieve stable electron configurations
• Types of covalent bonds

  • Single
  • Double
  • Triple
• Dative covalent (coordinate) bond


  One atom donates a pair of electrons to another atom
• Giant covalent structures


  • Graphite: Layers of hexagons with delocalized electrons, can conduct electricity
  • Diamond: Tetrahedral structure with strong covalent bonds, does not conduct electricity
• Determining molecular shape


  1. Count the number of bond pairs and lone pairs on the central atom

  2. Use the VSEPR theory to determine the shape based on minimizing electron pair repulsion
• Molecular shapes with no lone pairs

  • Linear (2 bond pairs)
  • Trigonal planar (3 bond pairs)
  • Tetrahedral (4 bond pairs)
  • Trigonal bipyramidal (5 bond pairs)
  • Octahedral (6 bond pairs)
• Molecular shapes with lone pairs

  • Pyramidal (3 bond pairs, 1 lone pair)
  • Bent (2 bond pairs, 2 lone pairs)
  • Trigonal pyramidal (3 bond pairs, 2 lone pairs)
• Octahedral


  Molecular shape with 6 bond pairs or lone pairs arranged in an octahedral geometry
• Pyramidal


  • Molecular shape with 3 bond pairs and 1 lone pair
  • Example: ammonia
• Bent/Nonlinear


  • Molecular shape with 2 bond pairs and 2 lone pairs
  • Bond angle shrinks from 107 to 104.5 degrees
• Trigonal planar


  • Molecular shape with 3 bond pairs and 2 lone pairs
  • Bond angle remains at 120 degrees
• Octahedral with 4 bond pairs and 2 lone pairs

  • Molecular shape
  • Bond angle remains at 90 degrees
• Electronegativity


  Ability of an atom to attract electrons towards itself in a covalent bond
• The further up and right in the periodic table, the more electronegative the element (excluding noble gases)
• Polar bond


  Covalent bond where atoms have a difference in electronegativity, resulting in an uneven distribution of electrons
• Polar bonds


  • H-Cl
  • H2O
• Symmetrical molecules with polar bonds can still be nonpolar overall
• Van der Waals forces


  Weak intermolecular forces caused by temporary induced dipoles
• Van der Waals forces

  • Hold iodine molecules together in a crystal structure
• Dipole-dipole forces


  Stronger intermolecular forces between permanent dipoles
• Hydrogen bonding


  Strongest intermolecular force, occurs between hydrogen and highly electronegative elements like N, O, F
• Ice expands when cooled due to hydrogen bonding pushing molecules apart
• Metallic bonding


  Bonding in metals where positive metal ions are attracted to a sea of delocalized electrons
• Metallic bonding


  • Allows high thermal and electrical conductivity
  • Gives high melting points
• Particle model states


  • Solid
  • Liquid
  • Gas
• Bond types


  • Giant covalent
  • Simple molecular
  • Giant ionic
  • Metallic