Comprehensive Lewis Structures & Molecular Properties

Periodic Table Facts & Valence-Electron Counts

• Second period non-metals (row 2)

  • \text{B} in Group 3A ⇒ 3 valence e⁻

  • \text{C} (4A) ⇒ 4 valence e⁻

  • \text{N} (5A) ⇒ 5 valence e⁻

  • \text{O} (6A) ⇒ 6 valence e⁻

  • \text{F} (7A) ⇒ 7 valence e⁻

• Below-row analogues share the same counts (e.g.

  • \text{Cl}\,, \text{Br} also 7 e⁻; \text{Be}=2, \text{Li}=1).

Ideal Bond Numbers & Octet Heuristics

• Non-metals seek an octet (8 e⁻) by acquiring electrons.

  • \text{F} needs 1 ⇒ likes 1 bond

  • \text{O} needs 2 ⇒ likes 2 bonds

  • \text{N} needs 3 ⇒ likes 3 bonds

  • \text{C} needs 4 ⇒ likes 4 bonds

• Metals (left side) prefer to lose e⁻.

  • \text{B} (3 e⁻) gives 3 ⇒ still 3 bonds but often incomplete octet.

  • \text{Be} gives 2 ⇒ 2 bonds

  • \text{Li} gives 1 ⇒ 1 ionic bond

Expanded Octet Elements (Row 3+)

• Energy level 3 has subshells 3s, 3p, 3d (capacity 2+6+10=18 e⁻).

• \text{P, S, Cl} may exceed 8 e⁻ (≥ 5 or 6 bonds).

  • e.g. \text{SF}6 (6 bonds), \text{PCl}5 (5 bonds), \text{ClO}_4^- (7 bonds overall on Cl).

Lone-Pair Trend for Second-Row OUTER Atoms

• 1 bond ⇒ 3 lone-pairs (halogen pattern)

• 2 bonds ⇒ 2 lone-pairs (O in \text{H}_2\text{O})

• 3 bonds ⇒ 1 lone-pair (N in \text{NH}_3)

• 4 bonds ⇒ 0 lone-pairs (C in \text{CH}_4)

Formal Charge Formula

\boxed{\text{FC}=V-(B+D)}

• V = valence e⁻

• B = number of bonds (each counts as 1, not 2)

• D = non-bonding electrons on atom.

• Deviation from ideal bond count ⇒ non-zero FC.

Multiple-of-Eight Rule (for non-H compounds)

• If total valence e⁻ is a multiple of 8 and no hydrogens are attached:

  • The central atom will carry 0 lone pairs.

  • Quick locator for expanded-octet patterns.

Bond Order & Definition

• For simple molecules the bond order equals the bond multiplicity.

  • \text{F}_2 single ⇒ \text{BO}=1

  • \text{O}_2 double ⇒ \text{BO}=2

  • \text{N}_2 triple ⇒ \text{BO}=3

Molecular vs. Electron-Pair Geometry (VSEPR)

Group count = # atoms attached + # lone pairs on center.

• 2 groups ⇒ Linear

• 3 groups ⇒ Trigonal planar

• 4 groups ⇒ Tetrahedral

• 5 groups ⇒ Trigonal bipyramidal

• 6 groups ⇒ Octahedral

If center owns lone pair(s) ⇒ molecular shape name changes (bent, trigonal-pyramidal, seesaw, square-pyramidal, square-planar, etc.).

Characteristic Bond Angles

Linear: 180^{\circ}

Trigonal planar: 120^{\circ}

Tetrahedral (no lone pair): 109.5^{\circ}

Trigonal pyramidal NH3: 107^{\circ}

Bent H2O: 104.5^{\circ}

Trigonal bipyramidal axes: axial–axial 180^{\circ}; axial–equatorial 90^{\circ}; eq–eq 120^{\circ}

Octahedral: 90^{\circ} (eq–eq & ax–eq); axial 180^{\circ}

Hybridisation Shortcut

Total groups → hybridisation set where exponents sum to groups.

• 2 groups ⇒ sp

• 3 groups ⇒ sp^2

• 4 groups ⇒ sp^3

• 5 groups ⇒ sp^3d

• 6 groups ⇒ sp^3d^2

Worked Diatomic Examples

• \text{F}2:7\times2=14 e⁻ → single bond, each F with 3 lone pairs ⇒ non-polar. • \text{O}2:6\times2=12 e⁻ → double bond, 2 lone pairs/atom ⇒ non-polar.
• \text{N}2:5\times2=10 e⁻ → triple bond, 1 lone pair/atom ⇒ non-polar. • \text{H}2 ⇒ single bond only; H satisfies duet.

Incomplete Octet Case: \text{BH}_3

• 6 total e⁻. Boron adopts 3 single bonds; octet incomplete (only 6 e⁻) but FC = 0.

• Molecular geometry: trigonal planar (120°), non-polar.

Water-Family Cases

• \text{H}_2\text{O} → bent; FC(O)=0; polar; \angle H–O–H=104.5^{\circ}; sp^3.

• \text{H}_3\text{O}^+ → trigonal pyramidal; FC(O)=+1; sp^3.

• \text{OH}^- → linear fragment; FC(O)=−1; 3 lone pairs on O.

Polarity & Dipole Moments Basics

• Polar bond: \Delta EN ≥ 0.5 (e.g.
  \Delta EN_{\text{C−O}} = 3.5-2.5=1.0 ⇒ polar).

• Whole-molecule polarity depends on vector sum of individual bond dipoles.

  • Symmetric shapes with identical outer atoms (e.g.
    \text{CO}2 linear, \text{BF}3 trigonal planar, \text{SF}_6 octahedral) ⇒ net dipole =0.

  • Lone-pair asymmetry or mixed atoms yields net polarity (e.g.
    \text{SO}2 bent, \text{NH}3 trigonal pyramidal, \text{SF}_4 seesaw).

Key Comparisons

• \text{CO}_2 (linear) vs. \text{CO} (diatomic):

  • Both have polar C–O bonds.

  • Symmetry cancels dipoles in CO$_2$ ⇒ non-polar.

  • CO has unequal dipole; FC diagram yields C$^{−}$–O$^{+}$; polar.

• \text{SO}2 is polar because of bent shape + lone pair; contrasts with non-polar CO$2$.

Expanded-Octet Examples

• \text{SF}_6: 48 e⁻ (multiple of 8), octahedral, sp^3d^2, non-polar, S with 6 bonds.

• \text{PCl}_5: trigonal bipyramidal, sp^3d, non-polar.

• \text{IF}_5: square pyramidal, 1 lone pair → polar, sp^3d^2.

• \text{XeF}_4: square planar, 2 lone pairs opposite → non-polar.

• \text{SF}_4 (34 e⁻) → seesaw, 1 lone pair, polar, sp^3d.

• \text{SF}_2 (20 e⁻) → bent, polar, sp^3.

• \text{I}_3^-: linear (I–I–I), 3 lone pairs on center I; electron-pair geometry trigonal bipyramidal, sp^3d; net charge –1.

Multiple-of-8 Rule in Practice

• For \text{SO}2\text{Cl}2: 32 e⁻ → no lone pair on S; four single bonds (2 O, 2 Cl). Tetrahedral sp^3.

• For \text{XeOF}_2: 28 e⁻ → 2 lone pairs, 3 bonds → T-shaped molecular geometry.

Resonance Essentials

• Occurs when electrons can shift without moving nuclei.

• Curved arrows show 2-e⁻ (full) or 1-e⁻ (half) motion.

• The resonance hybrid is the weighted average; bond order may be fractional.

Examples

1. Carbonate \text{CO}_3^{2-} → three equivalent structures with one C=O double bond rotating.

2. Nitrite \text{NO}_2^- → two equivalent structures (one N=O double bond each).

3. Nitrate \text{NO}_3^- → three equivalent.

4. \text{SO}_2 best structure has S=O double bond to each O (FC=0 on S).

5. \text{BF}_3 can draw B–F double-bond variant, but neutral all-single-bond form (incomplete octet) is more stable (all FC = 0).

Formal-Charge/Valence Rule for Center Atoms

To get FC = 0 on the center:
B = V - D
If dots found from 8-rule, number of bonds follows directly.

Selecting the Central Atom

• Choose the least electronegative or the atom capable of most bonds.

  • e.g.
    \text{SF}_2 → S central (can handle 4 bonds) instead of O or F.

• Hydrogen is never central; halogens rarely central unless expanded octet (e.g.
  \text{ClO}_4^-).

Polyatomic Ion Series (with Expanded Octet)

• Sulfate \text{SO}_4^{2-} → S with 6 bonds (2 double, 2 single), FC(S)=0, tetrahedral.

• Phosphate \text{PO}_4^{3-} → P with 5 bonds; trigonal bipyramidal e-pair geometry → tetrahedral molecular shape (all singles, charge on O’s).

• Chlorine oxy-anions

  • Perchlorate \text{ClO}_4^- → 7 bonds on Cl.

  • Chlorate \text{ClO}_3^- → 5 bonds; 1 lone pair.

  • Chlorite \text{ClO}_2^- → 3 bonds; 2 lone pairs.

  • Hypochlorite \text{ClO}^- → 1 bond; 3 lone pairs.

Radicals (Odd-Electron Species)

• \text{NO}_2 (17 e⁻) → N has 3 bonds + 1 single e⁻; bent; reactive.

• \text{NO} (11 e⁻) → N–O double; assign unpaired e⁻ to less-EN atom (N). Choose distribution giving smallest |FC|.

Resonance & Charge Delocalisation Example: \text{SCN}^-

• Two major forms:

  1. \text{S}^{-}–\text{C}\equiv\text{N}

  2. \text{S}=\text{C}=\text{N}^{-}

• Greater stability putting negative charge on larger atom (S) despite lower EN.

Organic Lewis-Structure Shortcuts (CxHyOz etc.)

Carbon family:

• Alkane C–C single (ethane \text{C}2\text{H}6).

• Alkene C=C (ethene \text{C}2\text{H}4).

• Alkyne C≡C (acetylene \text{C}2\text{H}2).
  Functional groups & drawing tips

• Alcohol \text{–OH}: methanol \text{CH}_3\text{OH} (O 2 bonds + 2 LP).

• Aldehyde: \text{CH}_3–\text{C}(=\text{O})H.

• Ketone: \text{CH}3–\text{C}(=\text{O})–\text{CH}3 (acetone).

• Ether: \text{CH}3–\text{O}–\text{CH}3 (dimethyl ether).

• Carboxylic acid: \text{CH}_3–\text{C}(=\text{O})\text{OH} (acetic acid).

• Ester: \text{CH}3–\text{C}(=\text{O})\text{O}–\text{CH}3.

• Amine: \text{C}2\text{H}5\text{NH}_2; N 3 bonds + lone pair.

• Amide: \text{CH}3–\text{C}(=\text{O})\text{NH}2.

• Nitrile: \text{R–C}\equiv\text{N} (triple bond).

Ethical & Practical Connections

• Molecular polarity informs solubility (e.g.
  non-polar \text{SF}6 as electrical insulator; polar \text{NH}3 dissolves in water).

• Formal charge & resonance guide reactivity sites in organic/biochemical reactions.

• Expanded-octet hypervalent species (e.g.
  \text{PCl}_5) are used in phosphorylation & chlorination industrial processes.

Study/Exam Tips

• Always count valence e⁻ first; write total.

• Identify central atom by minimal EN & maximal bonding capacity.

• Apply multiple-of-eight & bond-dot=valence rule for center FC = 0.

• Check octet/duet for each atom (remember B & Be exceptions).

• Calculate FC quickly; adjust with multiple bonds for neutrality/minimum charge.

• Determine group count → geometry → hybridisation.

• Sketch dipoles to predict polarity.

• For resonance, move only π or lone-pair electrons, never atoms.

• Practise with oxy-anions & expanded-octet examples; they recur in exams.