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Chapter 1-7 Notes: Lewis Structures and Formal Charge (Vocabulary)

Lewis Structures, Octet Rule, and Formal Charge — Study Notes

  • Core idea: Lewis structures are a way to depict how atoms in a molecule are bonded and where lone pairs reside. The octet rule is foundational for many main-group elements, with occasional exceptions discussed later (e.g., boron, aluminum, sulfur, phosphorus).
  • Structural varieties:
    • Acyclic (open-chain) vs cyclic (ring) structures.
    • Insaturation concept (rings and multiple bonds) will be covered later; rings often imply the possibility of additional double bonds to satisfy valence and electron count.
  • When exploring a molecular formula, multiple structures may be possible (constitutional isomers).
    • Constitutional isomers: same molecular formula but different connectivity.
    • If two drawn structures are just different placements of atoms without changing connectivity, they are the same structure, not constitutional isomers.
  • Practical workflow emphasized in the talk:
    • Start with the skeleton, then fill in hydrogens to satisfy octets (except when constraints require otherwise).
    • Use electron-counting and octet checks as a “sanity check” to avoid violating fundamental rules.
    • If you exceed octets for non-hydrogen atoms or misplace electrons, revise by shifting bonds or hydrogens or introducing double bonds.
  • The relationship between formal charge and molecular formula:
    • Formal charge helps locate electron-rich vs electron-poor regions, which is useful for predicting reactivity and ion stability in solution.
    • The speaker emphasizes checking both neutral and charged species to understand where charges reside and how they affect connectivity.
  • Value of practice: The lecturer stresses repeated practice with problems (e.g., from Canvas handouts) to master electron counting, octet rules, and formal charges.

Key concepts and definitions

  • Octet rule: Most second-row elements (C, N, O, F) prefer to have eight electrons around them (a full octet) in stable structures.
  • Formal charge (FC): a bookkeeping method to determine the distribution of electrons in a molecule.
  • Anion vs cation:
    • Anion: a species with a negative formal charge (an extra electron).
    • Cation: a species with a positive formal charge (one electron fewer).

Formal charge: definition and calculation

  • General rule used in the lecture: for a non-hydrogen atom, formal charge is calculated as
    • ext{FC} = V - (B + L)
    • where:
    • V = valence electrons of the atom (from the periodic table)
    • B = number of bonds to the atom (each bond counts as 1; a double bond counts as 2 bonds, etc.)
    • L = number of nonbonded electrons (lone-pair electrons) on the atom
  • Important interpretation:
    • If FC = 0 for an atom, that atom is neutral in the context of that bonding situation.
    • If FC = -1 for an atom, that atom bears a negative formal charge (electron-rich site).
    • If FC = +1 for an atom, that atom bears a positive formal charge (electron-poor site).
  • Worked examples (using the same counting approach as the lecture):
    • Oxygen with one bond to something and three lone pairs:
    • V(O) = 6, B = 1, L = 6
    • ext{FC}(O) = 6 - (1 + 6) = -1
    • So the oxygen bears a negative formal charge when it has one bond and three lone pairs.
    • Carbon with four bonds and no lone pairs:
    • V(C) = 4, B = 4, L = 0
    • ext{FC}(C) = 4 - (4 + 0) = 0
    • Nitrogen with four bonds and no lone pairs:
    • V(N) = 5, B = 4, L = 0
    • ext{FC}(N) = 5 - (4 + 0) = +1
    • A nitrogen with four bonds is positively charged (a common carbocation-like nitrogation in resonance forms).
    • Carbon with three bonds and one lone pair:
    • V(C) = 4, B = 3, L = 2
    • ext{FC}(C) = 4 - (3 + 2) = -1
    • A carbon with three bonds and one lone pair bears a negative charge in this counting framework.
  • Key trends to recognize (with practice, these become intuitive):
    • Oxygen with one bond and three lone pairs → FC = −1 (electron-rich O−).
    • Nitrogen with four bonds → FC = +1 (positively charged N).
    • Carbon with four bonds and zero lone pairs → FC = 0 (neutral carbon).
    • Carbon with four bonds is typically neutral in these countings unless other constraints push a charge elsewhere in the molecule.
  • Important caveats about octets and common exceptions:
    • Most atoms depicted in this course (C, N, O, etc.) prefer eight electrons around them.
    • Boron (B) and aluminum (Al) are mentioned as cases that may not follow a full octet in some structures (electron-deficient species).
    • Sulfur (S) and phosphorus (P) can sometimes expand the octet in some contexts (beyond eight electrons).

Electron counting for a given formula and charged species

  • General approach:
    • Start with the total number of valence electrons from all atoms in the formula: sum of each atom's valence electrons.
    • If the species has a negative charge, add one electron for each negative charge to the total count.
    • If the species has a positive charge, subtract one electron for each positive charge from the total count.
    • Use the total electron count to guide the construction of the skeleton (non-hydrogen atoms first), then fill in hydrogens to satisfy octets, and finally adjust by forming multiple bonds if necessary to accommodate all electrons.
  • Example framework (for a neutral C4H8O-like formula):
    • Count valence electrons: 4\times 4\,+\,8\times 1\,+\,6=(16+8+6)=30 electrons.
    • Build the non-hydrogen skeleton, distribute hydrogens, and then check octets. If you need to adjust to satisfy the total electron count, consider forming additional double bonds or moving hydrogens around.
  • Example with a charge (anionic):
    • If the formula carries a negative charge, add one electron to the total count for each negative charge before constructing the structure.
    • If the formula carries a positive charge, remove one electron from the total count for each positive charge before constructing the structure.

From bond-line (skeletal) formulas to full Lewis structures

  • Bond-line (skeletal) formula basics:
    • Carbons are implied at every vertex and at the ends of lines.
    • Hydrogens on carbons are implicit; hydrogens on heteroatoms (O, N, S, etc.) are shown when they are bonded to those atoms (not implied if attached to a carbon).
    • Heteroatoms (O, N, S, etc.) appear explicitly; lone pairs on these atoms are often necessary to satisfy octets.
  • Translating a bond-line to a Lewis structure:
    • Identify the carbon skeleton from vertices/ends of lines (implicit C’s).
    • Count the required hydrogens on each carbon to satisfy its octet, given its degree (number of bonds to other atoms).
    • Add lone pairs on heteroatoms to satisfy octets where needed (e.g., O with two bonds typically has two lone pairs; O with one bond and three lone pairs carries a negative charge in the examples discussed).
    • Place any hydrogens on non-carbon atoms if they are explicitly shown bonded to heteroatoms; otherwise, hydrogens on carbons are implicit.
  • Example logic illustrated in the talk (paraphrased):
    • A bottom-line hydrocarbon (bond-line) can be converted into a full Lewis structure by: starting at the end nearest a branch, counting bonds, and adding hydrogens to complete octets; then accounting for lone pairs on oxygens or other atoms; if a heteroatom lacks an octet, add lone pairs or adjust bonds as needed.
    • The full Lewis structure sometimes appears unwieldy (e.g., amoxicillin) but the bond-line representation is a compact, communication-friendly shorthand that you can later translate back to a Lewis structure when needed.
  • Summary rule of thumb for bond-line to Lewis: the number of hydrogens on a carbon is the number of bonds needed to reach four total bonds for that carbon (unless it bears a formal charge or special context). For heteroatoms, ensure they have the correct number of lone pairs to complete their octet, and place hydrogens on heteroatoms only when necessary.

Practical example themes discussed

  • Amoxicillin and penicillin are given as real-world examples of complex structures where the Lewis diagram would be unwieldy, hence the bond-line shorthand is preferred for communication in organic chemistry.
  • The speaker notes that the bond-line approach is the standard in organic chemistry for communicating complex structures efficiently, with Lewis structures used selectively to highlight electron pairs, charges, and formal charges.
  • The instructor emphasizes that, while the bond-line approach is efficient, you should still be able to reconstruct a Lewis structure when needed to consider lone pairs and formal charges.

Strategies and tips from the lecture

  • Always perform electron counting and octet checks early to avoid building structures that violate fundamental rules.
  • Practice problems are essential for mastery; use the Canvas handout and additional problems to reinforce the skill.
  • When you think you’ve mastered it, practice more. Mastery requires consistent practice.
  • If you’re stuck, go back to the core rules: valence electrons, octet maintenance, and appropriate formal charge assignments, then try alternative bonding patterns (e.g., add a double bond, shift a hydrogen) to satisfy the electron count and octet rules.
  • If you want to check whether two structures are different constitutional isomers, compare their connectivity (the order in which atoms are connected). If connectivity differs, they are constitutional isomers; if connectivity is the same but atoms are rearranged without changing the order of connections, they are the same structure.

Quick reference formulas and concepts

  • Formal charge formula (as used in class):
    • ext{FC} = V - (B + L)
    • V = valence electrons of the atom (periodic table value).
    • B = number of bonds to the atom (each bond counts as 1; double bonds count as 2, etc.).
    • L = number of nonbonded electrons (lone pairs) on the atom.
  • Typical valence electrons (V) for common atoms:
    • C: V = 4
    • N: V = 5
    • O: V = 6
    • F: V = 7
    • P: V = 5 (in some contexts; treat as V = 5 for the counting scheme shown in class)
  • Common octet expectations:
    • Carbon, nitrogen, and oxygen atoms generally seek octets in neutral, common organic structures.
    • Boron and aluminum can be electron-deficient (less than an octet) in some stable structures; sulfur and phosphorus can exceed the octet in some contexts.
  • Charge handling in electron counting:
    • Negative charge (anion) → add one electron to the total electron count.
    • Positive charge (cation) → remove one electron from the total electron count.
  • Conceptual take-home:
    • The distribution of formal charges helps identify the most likely resonance forms and the sites where reactions are likely to occur.
    • The bond-line formula is a compact way to represent large organic molecules, while the Lewis structure reveals lone pairs and formal charges when needed.

Summary takeaways for exam prep

  • Be able to count electrons, assign hydrogens to satisfy octets, and verify that you do not violate the octet rule for main-group elements (with noted exceptions).
  • Be able to calculate formal charges on all non-hydrogen atoms and identify the atom(s) bearing charge in a given structure.
  • Distinguish constitutional isomers from identical connectivity structures; realize that moving hydrogens without changing connectivity does not create a new isomer.
  • Convert between bond-line (skeletal) formulas and full Lewis structures by filling in hydrogens and lone pairs according to octet rules and valence considerations.
  • Practice with several problems (as provided on Canvas) to solidify the habits of electron counting, octet verification, and formal charge assessment.

Note: The discussion includes several worked verbal examples and references to specific structures and counts (e.g., neutral vs charged species, C4H8O-like formulas, and charged species with O− or N+. These illustrate the general principles above. Values like total electrons (30, 32, 26, etc.) in the transcript depend on the exact formula and charges considered; the core method remains consistent: sum valence electrons, adjust for charge, build skeleton, satisfy octets, and compute formal charges to check.