TR

Lewis Structures

  • Key to bonding: electron sharing to achieve a full valence shell
  • Electrons are distributed in such a way that each element is surrounded by 8 electrons (an octet)
    • There are exceptions to the octet rule
    • Helium and hydrogen
    • Molecules with an odd number of valence electrons because the octet rule works by pairing electrons
      • Radical species: molecule with an unpaired electron
      • Usually very reactive
    • Octet deficient molecules
      • Some molecules are stable with an incomplete set
      • Group 13 elements Boron and Aluminum have this property
    • Valence shell expansion
      • Elements with n = or > 3 - have empty d orbitals, which means more than 8 electrons can fit around the central atom
      • Expanded shells are most common when the central atom is large and bonded to small, electronegative atoms
  • Each dot on a lewis structure represents a valence electron
  • Bonding electrons: electrons involved in bonding
  • Lone pair electrons: electrons not involved in bonding

Procedure for Drawing Lewis Structures

  • Draw a skeleton structure
    • H and F are always terminal atoms (on the end)
    • The element with the lowest ionization energy goes in the middle (with some exceptions)
  • Count the total number of valence electrons
    • If there’s a negative ion, add the absolute value of total charge to the count of valence electrons
    • If there’s a positive ion, subtract
  • Count the total number of electrons needed for each atom to have a full valence shell
  • Subtract the number of valence electrons (from step 2) from the total number of electrons for full shells (from step 3) to get the number of bonding electrons
  • Assign 2 bonding electrons to each bond
  • If bonding electrons remain, make double or triple bonds
    • Generally, double bonds only form between C, N, O, and S.
    • Triple bonds are usually restricted to C, N, O.
  • If valence electrons remain, assign as lone pairs, giving octets to all atoms except hydrogen
  • Determine formal charge

Formal Charge

  • Formal charge: a measure of the extent to which an atom has gained or lost an electron in the process of forming a covalent bond
  • FC=V-L-(½)B
    • V = number of valence electrons
    • L = number of lone pair electrons
    • B = bonding electrons
  • The sum formal charge of each atom should equal the total charge of the molecule
  • Formal charge ≠ oxidation number
  • Structures with lower absolute values of formal charges are more stable (because they have lower energy) - closest to zero
    • If 2 structures have the same absolute value of FC, the lower energy structure is the one where the negative charge is on the most electronegative atom

Resonance Structures

  • Resonance hybrid: a blend between two structures
  • Electrons in resonance structures are delocalized
  • Electron pairs are shared over several atoms, not just 2.
  • Resonance structures: 2 or more structures with the same arrangement of atoms, but a different arrangement of electrons
  • With resonance structures, there is no right or wrong structure - both are needed to describe the molecule

Resonance Stabilization

  • An electron energy level is more stable when it has more volume that it can occupy
  • A molecule with several resonance forms has more space for the electrons to be in
  • A molecule with several resonance forms has additional stability due to resonance -> resonance stability

Resonance Bonds

  • Resonance bonds in a molecule are somewhere between single bonds and double bonds
  • They’re harder to break than single bonds but easier to break than double bonds
  • They’re shorter than single bonds but longer than double bonds
  • Different resonance structures can be found by pushing electrons around (into/out of double bonds/atoms)
  • Use a curved arrow to show how electrons are moved
  • Never move electrons into expanded valence for C, N, O, or F
  • Conjugated bonds: alternating double bonds
    • Conjugated systems can highly increase the volume that the electrons can be in