Lewis Structures & Chemical Bonding – Key Vocabulary

Lewis Symbols & Valence Electrons

• Only the valence-shell electrons (outer-most n level) are shown in a Lewis representation.

• Each valence electron = one dot placed on the four cardinal positions (left, right, top, bottom).
  • NO diagonal placement.
  • Dots may be placed singly (for 1–4 e⁻) and then paired.

• Examples
  • \text{Na} : 1\;e^- → Na·
  • \text{Cl} : 7\;e^- → ·Cl: with a total of seven dots, arranged 2-2-2-1.

Octet Rule – The Driving Force for Bonding

• Main-group atoms tend to attain 8 valence electrons (the noble-gas configuration).
  • They may lose, gain, or share electrons to do so.
  • Mnemonic: “oct” as in octopus (8 legs).

• Consequences
  • Metals (e.g. Na) tend to lose electrons → cations.
  • Non-metals (e.g. Cl) tend to gain electrons → anions.
  • Two non-metals often share electrons → covalent bonds.

Ionic vs. Covalent Bonding

• Ionic
  • Complete transfer of e⁻ (e.g. Na 3s electron → Cl 3p).
  • No physical bond line drawn; attraction is purely electrostatic.
  • Partner can change if a stronger opposite charge appears.

• Covalent
  • Sharing of e⁻ pairs between two non-metals.
  • Shared pair = one line (–) replacing two dots.
  • After sharing, both atoms count the pair toward their octet (roommate metaphor).

Bond Vocabulary & Trends

• Lone pair / unshared pair: non-bonding pair; potential donor.

• Single bond = 1 line = 2\;e^-
  Double = 2 lines = 4\;e^-
  Triple = 3 lines = 6\;e^-

• Bond length: single > double > triple.
  Bond strength (bond dissociation): triple > double > single (bundle-of-sticks analogy).

Universal Algorithm for Lewis Structures (Must Show ALL Math!)

1. Count total valence electrons
   E\text{total}=\sum E\text{valence}(\text{each atom}) \;\pm\;\text{ion charge}

2. Draw the skeleton with single bonds:
   • Central atom = least electronegative (never H).

3. Count electrons used in the skeleton.
   E_\text{used}=2\times(\text{number of bonds})

4. Electrons left
   E\text{left}=E\text{total}-E_\text{used}

5. Compute electrons needed to complete all octets
   For every atom: E\text{need}=8-E\text{owned}
   (H only needs 2.)

6. Compare need vs. left
   • If E\text{left}=E\text{need} → just distribute lone pairs.
   • If E\text{left}\text{need} → you have an electron deficiency → create multiple bonds until counts match.
   – Deficiency of 2 → add one double bond.
   – Deficiency of 4 → one triple OR two doubles, etc.

7. Verify total electron count in the final drawing.

Worked Example 1 – \text{PCl}_3

1. Valence: 5 + 3(7)=26\;e^-

2. Skeleton: P central; three P–Cl single bonds.

3. E_\text{used}=3\times2=6\;e^-

4. E_\text{left}=26-6=20\;e^-

5. Need:
   • P owns 6 → needs 8-6=2\;e^-.
   • Each Cl owns 2 → needs 6\;e^-; three Cl → 18\;e^- total.
   • Overall need = 2+18=20\;e^-

6. Need equals left → distribute lone pairs.

7. Check: 13 pairs = 26 e⁻. ✓

Worked Example 2 – \text{NO}^+ (cation)

1. Valence: N:5 + O:6 - 1 = 10\;e^-

2. Skeleton: N–O.

3. E_\text{used}=2\;e^-

4. E_\text{left}=8\;e^-

5. Need:
   • Each atom owns 2 → each needs 6 → total 12.

6. E\text{left}\text{need} → 4-electron deficiency → either add one triple bond or two doubles.
   Best option with only two atoms: create a triple bond N≡O⁺.

7. Distribute remaining lone pairs; verify 10 e⁻.

Electron Deficiency & Multiple-Bond Options (Money Metaphor)

• Think of deficiency as \$\$\text{needed} to “buy” double (\$2) or triple (\$4) bonds.
  • $4 →$ 1 triple or 2 doubles.
  • $6 →$ 3 doubles or 1 triple + 1 double, etc.

Worked Example 3 – \text{CNS}^- & Resonance

Valence count = 4+5+6+1 = 16\;e^-
Four-electron deficiency → several valid structures (same skeleton SCN):

1. S–C≡N⁻

2. S≡C–N⁻

3. S=C=N⁻ (two double bonds)
   All satisfy octet & 16-e⁻ rule → resonance (drawn with double-headed arrows).

Resonance vs. Isomers

• Resonance structures
  • Same molecular

Lewis Symbols & Valence Electrons

• Only the valence-shell electrons (outer-most n level) are shown in a Lewis representation.

• Each valence electron = one dot placed on the four cardinal positions (left, right, top, bottom).

  • NO diagonal placement.

  • Dots may be placed singly (for 1–4 e⁻) and then paired.

• Examples

  • \text{Na} : 1\;e^- → Na·

  • \text{Cl} : 7\;e^- → ·Cl: with a total of seven dots, arranged 2-2-2-1.

Octet Rule – The Driving Force for Bonding

• Main-group atoms tend to attain 8 valence electrons (the noble-gas configuration).

  • They may lose, gain, or share electrons to do so.

  • Mnemonic: “oct” as in octopus (8 legs).

• Consequences

  • Metals (e.g. Na) tend to lose electrons → cations.

  • Non-metals (e.g. Cl) tend to gain electrons → anions.

  • Two non-metals often share electrons → covalent bonds.

Ionic vs. Covalent Bonding

• Ionic

  • Complete transfer of e⁻ (e.g. Na 3s electron → Cl 3p).

  • No physical bond line drawn; attraction is purely electrostatic.

  • Partner can change if a stronger opposite charge appears.

• Covalent

  • Sharing of e⁻ pairs between two non-metals.

  • Shared pair = one line (–) replacing two dots.

  • After sharing, both atoms count the pair toward their octet (roommate metaphor).

Bond Vocabulary & Trends

• Lone pair / unshared pair: non-bonding pair; potential donor.

• Single bond = 1 line = 2\;e^-

  • Double = 2 lines = 4\;e^-

  • Triple = 3 lines = 6\;e^-

• Bond length: single > double > triple.

• Bond strength (bond dissociation): triple > double > single (bundle-of-sticks analogy).

Universal Algorithm for Lewis Structures (Must Show ALL Math!)

1. Count total valence electrons
   E_\text{total}=\sum E_\text{valence}(\text{each atom})\;\pm\;\text{ion charge}

2. Draw the skeleton with single bonds:

   • Central atom = least electronegative (never H).

3. Count electrons used in the skeleton.
   E_\text{used}=2\times(\text{number of bonds})

4. Electrons left
   E_\text{left}=E_\text{total}-E_\text{used}

5. Compute electrons needed to complete all octets
   For every atom: E_\text{need}=8-E_\text{owned}
   (H only needs 2.)

6. Compare need vs. left

   • If E_\text{left}=E_\text{need} → just distribute lone pairs.

   • If E_\text{left}<E_\text{need} → you have an electron deficiency → create multiple bonds until counts match.

     • Deficiency of 2 → add one double bond.

     • Deficiency of 4 → one triple OR two doubles, etc.

7. Verify total electron count in the final drawing.

Worked Example 1 – \text{PCl}_3

1. Valence: 5 + 3(7)=26\;e^-

2. Skeleton: P central; three P–Cl single bonds.

3. E_\text{used}=3\times2=6\;e^-

4. E_\text{left}=26-6=20\;e^-

5. *Need:

   • P owns 6 → needs 8-6=2\;e^- .

   • Each Cl owns 2 → needs 6\;e^-; three Cl → 18\;e^- total.

   • Overall need = 2+18=20\;e^-*

6. Need equals left → distribute lone pairs.

7. Check: 13 pairs = 26 e⁻. ✓

Worked Example 2 – \text{NO}^+ (cation)

1. Valence: N:5 + O:6 - 1 = 10\;e^-

2. Skeleton: N–O.

3. E_\text{used}=2\;e^-

4. E_\text{left}=8\;e^-

5. *Need:

   • Each atom owns 2 → each needs 6 → total 12.*

6. E_\text{left}

7. Distribute remaining lone pairs; verify 10 e⁻.

Electron Deficiency & Multiple-Bond Options (Money Metaphor)

• *Think of deficiency as \text{needed} to “buy” double (2) or triple (4) bonds.

  • 4 \to 1 triple or 2 doubles.

  • 6 \to 3 doubles or 1 triple + 1 double, etc.*

Worked Example 3 – \text{CNS}^- & Resonance

Valence count = 4+5+6+1 = 16\;e^-$$
Four-electron deficiency → several valid structures (same skeleton SCN):

1. S–C≡N⁻

2. S≡C–N⁻

3. S=C=N⁻ (two double bonds)
   All satisfy octet & 16-e⁻ rule → resonance* (drawn with double-headed arrows).*

Resonance vs. Isomers

• Resonance structures

  • Same molecular*