Notes on Lewis Structures, Formal Charges, Electronegativity, and VSEPR

Stepwise approach to perfect Lewis structures

• Step 1: Arrange atoms in space based on connectivity given in molecular formula.

• Step 2: Add single bonds to all atoms that are connected to each other.

• Step 3: Identify all carbon atoms without a filled valence shell. For each such carbon atom, look for an adjacent atom that is also without a filled valence and connect with one or two multiple bonds.

• Step 4: Add lone pairs to fill all remaining unfilled valence shells.

• Step 5: Add any formal charges as identified by the table presented during the first lecture.

• Important caveat: This approach works for all but molecules with a carbocation. Do not worry about carbocations at this time.

Valence electrons in neutral atoms

• Hydrogen: $V_H = 1$

• Carbon: $V_C = 4$

• Nitrogen: $V_N = 5$

• Oxygen: $V_O = 6$

• Fluorine, Chlorine, Bromine, Iodine: $V_{F,Cl,Br,I} = 7$

• Note: These valence electron counts are the basis for counting electrons when constructing Lewis structures and for applying the formal charge calculation.

Formal charge identification

• General rule (formal charge):

  • For any atom, formal charge $FC = V - (L + frac{B}{2})$
  • Where:
  • $V$ = number of valence electrons of the neutral atom (from the values above),
  • $L$ = number of electrons in lone pairs on the atom,
  • $B$ = number of electrons involved in bonds around the atom (bonding electrons).

• Interpretation notes:

  • A neutral atom has $FC = 0$ in the Lewis structure.
  • A positive formal charge indicates the atom has fewer electrons than the neutral valence count in the structure.
  • A negative formal charge indicates the atom has more electrons than the neutral valence count in the structure.

• Example framework for counting (no specific molecule given):

  • If an atom has valence electrons $V$, is participating in $B$ bonding electrons and has $L$ nonbonding electrons, then the formal charge is
    $FC = V - (L + frac{B}{2}).$

• The typical table (as introduced in lectures) categorizes atoms by the distribution of bonds and lone pairs for neutral, positively charged, and negatively charged states, guiding how to assign charges when you complete the Lewis structure.

• Practical reminder: Use the octet rule (or duet for H) as a guide when assigning bonds and lone pairs, ensuring no atom exceeds its typical valence unless you intentionally form expanded octets (not typical for first-row elements in simple teaching problems).

Carbocation exception

• The described stepwise method (Step 1–Step 5) does not cover molecules containing a carbocation in the simplest form.
• When carbocations are involved, additional considerations are necessary beyond this basic approach.
• For now, you should be comfortable applying Steps 1–5 to neutral and typical anionic/covalent structures and note that carbocations require a separate treatment.

Pauling electronegativity values (conceptual)

• Pauling electronegativity (EN) assigns a relative pull of electron density toward an atom in a bond.
• On the Pauling scale, EN generally increases from left to right across a period and from bottom to top within a group.
• Typical representative values (illustrative, not exhaustively tabulated here):
  • Hydrogen: EN_H \n = 2.20
  • Carbon: $EN_C <br /> = 2.55$
  • Nitrogen: $EN_N <br /> = 3.04$
  • Oxygen: $EN_O <br /> = 3.44$
  • Fluorine: $EN_F <br /> = 3.98$
• For alkali/alkaline earth metals (examples from the transcript's partial table):
  • Lithium: $EN_{Li} <br /> ≈ 0.98$
  • Beryllium: $EN_{Be} <br /> ≈ 1.57$
  • Sodium: $EN_{Na} <br /> ≈ 0.93$
  • Magnesium: $EN_{Mg} <br /> ≈ 1.31$
• Practical implication: The EN values help explain bond polarity in heteronuclear bonds (difference in EN → dipole moment, partial charges in Lewis structures).
• Key takeaway: The greater the EN difference between two bonded atoms, the more polar the bond is likely to be; a very large difference can lead to ionic character in bonds.

VSEPR: Predicted molecular shapes

• Concept: Valence Shell Electron Pair Repulsion (VSEPR) predicts molecular geometry by considering regions of electron density around a central atom.
• Regions of electron density around central atom are counted as either bonding regions (single, double, or triple bonds) or lone-pair regions.
• Predicted distributions and bond angles:
  • 4 regions of electron density → Tetrahedral distribution: bond angles ≈ $109.5^ ext{°}$
  • Example: methane, $CH_4$ (central carbon with four C–H bonds).
  • 3 regions of electron density → Trigonal planar distribution: bond angles ≈ $120^ ext{°}$
  • Common examples include boron trifluoride, $BF_3$, and carbonate derivatives around a planar center.
  • 2 regions of electron density → Linear distribution: bond angles ≈ $180^ ext{°}$
  • Example: carbon dioxide, $O=C=O$ (central carbon with two double bonds to oxygens).
• Expressed succinctly:
  • Regions: 4 → tetrahedral (109.5°)
  • Regions: 3 → trigonal planar (120°)
  • Regions: 2 → linear (180°)
• Notes on application:
  • Actual molecular geometry may be distorted by lone pairs (e.g., ammonia NH₃ is trigonal pyramidal due to one lone pair, not a perfect trigonal planar arrangement around N).
  • The electron-pair geometry (based on regions) may differ from the molecular geometry (shape of the molecule) when lone pairs are present.

Connections to foundational principles and relevance

• Lewis structures provide a visual representation of electron distribution that underpins molecular geometry, reactivity, and polarity.
• Valence electron counts determine bonding capacity and the distribution of lone pairs; this links to octet (or duet) satisfaction and formal charges.
• Formal charge calculations help assess the plausibility of resonance forms and stability of the structure.
• Electronegativity values explain bond polarity and the likelihood of charge separation or dipole formation in molecules.
• VSEPR links electron-domain geometry to observable molecular geometries, which in turn influence physical properties and chemical behavior.

Practical implications and examples (conceptual)

• When designing or predicting a molecule's Lewis structure, start with the connectivity, satisfy valence requirements with single bonds, then introduce multiple bonds where needed to satisfy octets for carbons and other second-row elements.
• Use the formal charge formula to verify that your chosen resonance form or Lewis structure minimizes charge separation unless charge stabilization (e.g., in known ions) is desired.
• Consider electronegativity differences when predicting bond polarity and potential sites of nucleophilic or electrophilic attack.
• Apply VSEPR to anticipate molecular geometries from the number of electron domains around the central atom, keeping in mind lone-pair effects on actual shapes.

Note: All numerical references and formulas used above align with the content, including valence electron counts, the formal charge equation, and the VSEPR angle predictions, as discussed in the transcript. The examples provided (CH₄, BF₃, CO₂, etc.) illustrate how these concepts are applied in practice to interpret and predict chemical structures and properties.