BS

Chapter 2

  • 2.1 Molecular Representations

    • Lewis structures: shows all atoms, bonds, and lone pairs explicitly which is impractical for larger molecules

    • partially condensed structures: carbons and hydrogens are condensed (CH3, CH2, etc)

    • condensed structures: only double and triple bonds are drawn. atoms are clustered in groups when possible

    • molecular formulas only tell the number and type of atoms, not how they’re arranged

  • 2.2 Drawing Bond-Line Structures

    • rules

      • 1. single bonded carbon atoms that are chained together should be drawn in a zigzag

      • 2. double bonds should be drawn to maximize bond angles

      • 3. the direction of single bonds is irrelevant

      • 4. all heteroatoms must be drawn with any attached hydrogens

      • 5. carbon atoms must never be drawn with more than four bonds.

  • 2.3 Identifying Functional Groups

    • functional groups are groups of atoms and bonds that have a predictable chemical behavior

    • Relevant Functional Groups

  • 2.5 Identifying lone pairs

    • Formal charges must always be drawn in bond-line structures as they help infer where lone pairs are without them having to be drawn

    • steps:

      • determine the appropriate number of valence electrons for the atom in question

      • determine if the atom exhibits that number

    • oxygen bonding patterns

      • negative charge indicates one bond and three lone pairs

      • no charge indicates two bonds and two lone pairs

      • positive charge indicates three bonds and one lone pair

    • Nitrogen bonding patterns

      • negative: two bonds, two lone pairs

      • no charge: three bonds, one lone pair

      • positive: four bonds and no lone pairs

  • 2.6 3D bond-line structures

    • geometry can be shown using dashes and wedges to signify when an atom is coming off the page or extending behind it

  • 2.7 Intro to Resonance

    • the inadequacy of bond-line structures

      • bond-line doesn’t take into account the fact that bonding electrons aren’t bound to a single spot

      • this is accurate for some simple double bonded carbons but for others it isn’t

      • allyl carbocations rotate electron density across double bonds and lone pairs

    • Resonance

      • the concept that lone pairs and pi bonds shift across a molecule

      • represented by resonance hybrids and structures

    • Resonance stabilization

      • molecules that exhibit resonance tend to be more stable then those that don’t

      • the spread of electrons is referred to as delocalization, while LP that don’t exhibit resonance are called localized electrons

  • 2.8 Curved Arrows

    • arrows are used to explain the movement of electron density as you construct resonance structures

    • rules

      • avoid breaking single bonds

      • obey the octet rule for second-row elements

    • resonance structures must have the same connectivity

  • 2.9 Formal charges in resonance structures

    • as electrons shift across resonance structures, so does formal charges

  • 2.10 Drawing resonance structures via pattern recognition

    • allylic lone pairs

      • referring to lone pairs on atoms adjacent to a double bond (meaning on an atom bonded to one of the carbons/atoms participating in the double bond)

      • if not C=C, then allylic-like

      • LP forms a new double bond between the allylic atom and vinylic atom, and the double bond between the vinylic atoms dissolves into a LP on the opposite vinylic to maintain their octets

    • allylic carbocation

      • a carbocation adjacent to a double bond

      • the carbocation and double bond switch sides

      • a carbocation next to conjugated pi bonds (separated by 1 sigma bond) will do the same

    • a lone pair adjacent to a carbocation

      • a double bond will form between the atom the LP was on and the carbocation

    • a pi bond between two atoms with different electronegativity

      • the pi bond moves onto the electronegative atom as a lone pair

    • conjugated pi bonds enclosed in a ring

      • pi bonds push over one position

  • 2.11 assessing the relative importance of resonance structures

    • some resonance structures will contribute more or less to the resonance hybrid then the others

    • major contributors have

      • the greatest number of filled octets

      • the fewest formal charges

      • then, charges on the most appropriate atom (positive on a carbon vs a sulfur, negative on an oxygen vs a nitrogen).

    • if all structures are equivalent, then they contribute equally to the resonance hybrid.

  • 2.12 the resonance hybrid

    • the resonance hybrid is the combination of all the resonance structures and most accurately represents what the actual molecule looks like

    • illustrates how electron pairs are delocalized across a molecule

    • formal charges in resonance structures are drawn as partial charges in the hybrid

  • 2.13 delocalized and localized lone pairs

    • delocalized lone pairs

      • changes the geometry of the atoms it’s located on when it is on that atom

      • not counted when determining hybridization by counting electron domains

    • localized lone pairs

      • doesn’t participate in resonance and is therefore not allylic to a pi bond

      • a lone pair on an atom that also has a pi bond cannot participate in resonance and is as such, localized