Lewis dot structures: counting bonds and shared electron pairs

Core idea from transcript

• When drawing a Lewis dot structure for a covalently bonded substance, the first step is to determine the number of bonds or the number of shared electron pairs in the molecule so that we are able to construct the structure correctly.
• This step sets up how many electron pairs are being shared between atoms, which directly shapes the connectivity and geometry of the molecule.

Step-by-step approach: determine the number of bonds or shared pairs

• Step 1: Determine the total valence electrons in the molecule.
  • Formula: $V = \sum<em>i v</em>i$ where $v_i$ is the valence electrons of atom i.
• Step 2: Choose a central atom (usually the least electronegative element, and never H).
• Step 3: Arrange the atoms in a skeletal structure (skeleton) and connect them with single bonds.
  • Each single bond uses 2 electrons (one shared electron pair).
• Step 4: Distribute the remaining valence electrons as lone pairs to satisfy the octet rule for the outer atoms first.
• Step 5: Place any leftover electrons on the central atom to complete its octet.
• Step 6: If the central atom does not have an octet after distribution, convert lone pairs from surrounding atoms into additional bonds (double or triple bonds) to satisfy the octet.
• Step 7: Check that the total number of valence electrons used matches the initial count $V$.
• Step 8: Count the number of bonds or shared electron pairs in the final structure.
  • Important distinction: a single bond contains 1 shared electron pair, a double bond contains 2 shared electron pairs, etc.
• Note on terminology:
  • The number of bonds can refer to the number of bond connections (e.g., C–O single bonds) and the number of shared electron pairs can refer to the total bonding pairs across all bonds (e.g., two C=O double bonds contain 4 shared electron pairs in total).

Key terms and definitions

• Covalent bond: a bond formed by sharing electrons between atoms.
• Shared electron pair (bonding pair): the pair of electrons that is shared between two atoms to form a covalent bond.
• Lone pair (nonbonding pair): a pair of valence electrons not involved in bonding.
• Octet rule: most main-group elements tend to surround themselves with eight electrons around them in covalent compounds.
• Bond order: the number of shared electron pairs between two atoms (e.g., a single bond has bond order 1, a double bond has bond order 2).

Worked example: carbon dioxide (CO₂)

• Step 0: Transcript idea applied — determine number of shared electron pairs first.
• Step 1: Valence electrons: $v<em>C = 4, \, v</em>O = 6$; Total $V = v<em>C + 2 v</em>O = 4 + 2(6) = 16$
• Step 2: Central atom and skeleton: O – C – O
• Step 3: Start with two single bonds in the skeleton: number of bonds (connections) $B_{initial} = 2$
• Step 4: Electrons used by these bonds: $2B_{initial} = 4$
• Step 5: Remaining electrons: $L = V - 2B_{initial} = 16 - 4 = 12$
• Step 6: Distribute remaining electrons to satisfy octets on oxygen first:
  • Each O needs 6 more electrons to complete its octet (3 lone pairs each): total 12 electrons allocated to outer atoms.
  • After this distribution, carbon has only 4 electrons around it (from the two bonds), so its octet is not satisfied.
• Step 7: Form additional bonds to satisfy octet: convert to two C=O double bonds.
  • Final structure: O = C = O (two double bonds).
• Step 8: Bond counting in final structure:
  • Each C=O double bond contains 2 shared electron pairs, so total shared pairs = $2 imes 2 = 4$.
  • The final number of bonds (bond connections) between atoms = 2 (two C–O connections), but each is a double bond.
• Summary for CO₂:
  • Total valence electrons: $V = 16$
  • Final bonds: two C=O bonds
  • Total shared electron pairs: $4$

Special cases and extensions

• Expanded octet: For elements in period 3 or higher, the central atom can have more than eight electrons (hypervalent species).
• Resonance: Some molecules have multiple valid Lewis structures; the true structure is a resonance hybrid with bond orders averaged across resonance forms.
• Practical tips:
  • Always verify that outer atoms (except H) have octets where possible.
  • If you cannot satisfy octets, reconsider the central atom choice or bonding pattern (single vs multiple bonds).

Connections to foundational principles

• Electron counting is grounded in valence electron theory; predicting structure from valence electrons links to chemical bonding concepts.
• The octet rule and bond order connect to molecular stability and reactivity.
• The idea of distributing electrons to satisfy octets ties to electron domain geometry and VSEPR considerations for predicting shapes.

Practical implications and conceptual takeaways

• Knowing the number of shared electron pairs helps determine how many bonds and what type (single, double, triple) will be present.
• The process bridges simple electron counting with actual Lewis structures used to predict molecular geometry and reactivity.
• Counting errors can lead to incorrect bond orders, wrong geometry, or violation of the octet rule for some atoms.

Quick recap

• The first step in drawing a Lewis structure is to determine the number of shared electron pairs (bonds) in the molecule.
• Use valence electrons, build a skeleton, satisfy octets with lone pairs, and upgrade bonds as needed to satisfy octets.
• Remember the difference between bond connections and shared electron pairs in counting (single vs double bonds).
• CO₂ serves as a clear example where two double bonds lead to 4 shared electron pairs in total.