Bonding

• Ionic bonding


  Charged ions held together by strong electrostatic attractions
• Ionic bonding


  • Oppositely charged ions form to get a full shell of electrons
  • Sodium gives up an electron to chlorine
• Ionic charges


  • Group 1 forms +1
  • Group 2 forms +2
  • Group 3 rarely form ionic bonds, they are covalent
  • Group 5 forms 3-
  • Group 6 forms 2-
  • Group 7 forms 1-
• Molecular ions


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

  1. Swap charges between ions

  2. Drop charges to get subscripts

  3. Simplify
• Ionic compounds


  • Have giant ionic structures
  • Dissolve well in polar water
  • Conduct electricity when molten or dissolved
  • Have high melting points
• Covalent bonding


  Sharing of electrons between atoms to achieve full shells
• Types of covalent bonds

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


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


  • Graphite - layers with delocalized electrons, can conduct electricity
  • Diamond - tightly packed, cannot conduct electricity
• Molecular shape


  Determined by number of bond pairs and lone pairs of electrons
• 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 planar (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
• Tetrahedral


  • Molecular shape with 4 bond pairs and 2 lone pairs
  • 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 molecules


  • Water
  • Hydrogen chloride
• Nonpolar molecules


  • Chlorine
  • Hydrocarbons
• Intermolecular forces


  Weak forces between molecules, not within covalent bonds
• Van der Waals forces

  • Weakest intermolecular force, induced dipole-dipole interactions
  • Larger molecules have stronger van der Waals forces
• Dipole-dipole forces


  • Stronger than van der Waals, exist between permanent dipoles
• Hydrogen bonding


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


  Giant lattice of positive metal ions with delocalized electrons, responsible for high melting points and conductivity
• Particle model states


  • Solid: tightly packed, regular arrangement, high density, vibrate on spot
  • Liquid: tightly packed, random arrangement, high density, move freely and slide over each other
  • Gas: very spaced out, random arrangement, low density, move freely
• Bond types


  • Giant covalent: solids, don't conduct, high melting points
  • Simple molecular: liquids/gases, may conduct if polar, low melting/boiling points
  • Giant ionic: solids, conduct if dissolved/molten, high melting points
  • Metallic: solids/liquids conduct, high melting points