Chemical Bonding & Lewis Model – Comprehensive Notes

Chapter Outline

• Ionic Bonding

• Covalent Bonding

• Lewis Symbols & Lewis Structures

• Formal Charge & Resonance

• Strength of Ionic & Covalent Bonds

• Molecular Structure & Polarity

Bonding Theories (General)

• Purposes

  • Explain how & why atoms attach to form compounds

  • Rationalise stability vs. instability of atom combinations

  • Predict molecular shape, physical & chemical properties

• Driving force: lowering of potential energy (PE)

• Octet rule = first-order predictor for bond formation

Lewis Model – Fundamentals

• Valence electrons represented as dots around element symbol

• Lewis structure (electron-dot structure) depicts:

  • Arrangement of atoms (skeletal structure)

  • Bonding pairs vs. lone (non-bonding) pairs

  • Allows prediction of stability, geometry, and some properties

Why Do Atoms Bond?

• Bonds form when PE(bonded) < PE(separate)

• PE determined by three electrostatic terms

  • Nucleus–nucleus repulsion

  • Electron–electron repulsion

  • Nucleus–electron attraction

Types of Bonds

(relationship between atom types)

• Metal + Non-metal → Ionic (electron transfer)

• Non-metal + Non-metal → Covalent (electron sharing)

• Metal + Metal → Metallic (electron pooling/delocalisation)

Ionic Bonding — Key Points

• Metals (low $IE$) lose e⁻ → cations

• Non-metals (high $EA$) gain e⁻ → anions

• Oppositely charged ions attract → ionic bond

• Crystal lattice = 3-D array of alternating anions & cations

• Properties predicted from strong, nondirectional electrostatic forces:

  • High $T<em>m$ & $T</em>b$ (usually $>300\,^{\circ}\text C$)

  • Hard, brittle solids; empirical formulas only (no discrete molecules)

  • Conduct electricity when molten or aqueous; not as solids

Lattice Energy ( $U_{lat}$ )

• Energy released when gaseous ions form crystal: always exothermic

• Depends on charge magnitude and ionic radius (Coulomb’s law)
  $U<em>{lat}\propto \dfrac{|q</em>1q_2|}{r}$

• Born–Haber cycle allows indirect calculation via Hess’s law

  • Example NaCl cycle yields $U_{lat}=−788\,\text{kJ·mol}^{−1}$

• Trends

  • Larger ion radius → less exothermic $U_{lat}$ (LiCl −834 kJ > CsCl −657 kJ)

  • Greater charge magnitude → far more exothermic (MgO ≫ NaCl)

Covalent Bonding — Essentials

• Occurs between non-metals (high $IE$) where e⁻ removal costly

• Sharing pairs lowers PE; electrons localised between nuclei

• Types

  • Single bond (1 pair / 2 e⁻)

  • Double bond (2 pairs / 4 e⁻)

  • Triple bond (3 pairs / 6 e⁻)

• Terminology

  • Bonding pair = shared

  • Lone pair = not shared

• Covalent bonds are highly directional → discrete molecules

• Typical properties: low $T<em>m$, $T</em>b$ (<300 °C); variable hardness; poor conductors

Metallic Bonding (brief)

• Metal atoms donate valence e⁻ to a communal “sea”

• Delocalised e⁻ glue positive ion cores together → malleability, conductivity, luster

Valence Electrons & Electron Counting

• Main-group column no. gives # valence e⁻

• Transition metals usually counted as 2 valence e⁻ in Lewis context

• Octet rule: atoms seek 8 e⁻ in outer shell (duet for H). Exceptions exist.

Writing Lewis Structures (Molecules/Ions)

1. Skeleton: least electronegative atom central; H always terminal.

2. Count total valence e⁻ (add/subtract charge for ions).

3. Form single bonds (2 e⁻ each).

4. Distribute remaining e⁻ to give terminal atoms octets, then central.

5. If central lacks octet, convert lone pairs on adjacent atoms into multiple bonds.

6. Check formal charges.

Formal Charge (FC)

\text{FC}=#\,\text{valence e}^- − [\text{nonbonding e}^- + \tfrac12(\text{bonding e}^-)]
Guidelines

• $\sum FC = 0$ for neutral, = net charge for ion

• Prefer minimal |FC|; negative FC on more electronegative atom

Resonance

• Delocalisation represented by multiple Lewis forms differing only in electron placement.

• Actual structure = resonance hybrid (e.g., $\text O_3$, carbonate).

Octet Rule – Exceptions

• Odd-electron species (free radicals) e.g., $\text{NO}$, $\text{ClO}$

• Incomplete octet: B, Al (e.g., $\text{BF}_3$)

• Expanded octet: period 3+ elements with d-orbitals (e.g., $\text{H}<em>3\text{PO}</em>4$)

Electronegativity (EN) & Bond Polarity

• EN = ability to attract bonding e⁻ (Pauling scale)

• Trend: ↑ across period, ↓ down group; $F$ highest (4.0)

Bond Type via $\Delta EN$

• $\Delta EN=0$ → pure (nonpolar) covalent

• $0.1≤\Delta EN≤0.4$ → nonpolar covalent

• $0.4<\Delta EN<2.0$ → polar covalent

• $\Delta EN≥2.0$ → ionic

Dipole Moment

$\mu = q r$ (units Debye)

• Larger charge separation (q) or bond length (r) → larger $\mu$

Percent Ionic Character

$\%\,IC = \dfrac{\mu<em>{exp}}{\mu</em>{ionic}}\times100$

• Graph vs. $\Delta EN$ shows gradual continuum between covalent & ionic.

Bond Energies & Enthalpy of Reaction

• Bond energy (D): energy required to break 1 mol bond in gas; always $>0$ (endothermic)

• Trends: multiple bonds stronger/shorter than singles; strength increases across period, decreases down group.

• Reaction enthalpy estimate
  $\Delta H<em>{rxn}\approx\sum D</em>{\text{broken}} - \sum D_{\text{formed}}$
  (Breaking = +, Making = –)

Bond Lengths

• Distance between nuclei; averages tabulated.

• More shared e⁻ → shorter bonds; increases down group; decreases across period.

• Inverse relation: longer bond → weaker.

Valence Shell Electron Pair Repulsion (VSEPR)

• Electron groups (lone pairs, single/double/triple bonds, single e⁻) arrange to minimise repulsion.

• Five fundamental electron geometries:

  1. 2 groups → Linear, $180^{\circ}$

  2. 3 groups → Trigonal planar, $120^{\circ}$

  3. 4 groups → Tetrahedral, $109.5^{\circ}$

  4. 5 groups → Trigonal bipyramidal, $120^{\circ}/90^{\circ}$

  5. 6 groups → Octahedral, $90^{\circ}$

• Lone pairs distort angles (LP–LP > LP–BP > BP–BP).

Derivative Molecular Geometries

• 4 e-groups: trigonal pyramidal (1 LP), bent (2 LP)

• 5 e-groups: seesaw (1 LP), T-shaped (2 LP), linear (3 LP)

• 6 e-groups: square pyramidal (1 LP), square planar (2 LP)

Multiple Central Atoms

• Describe geometry at each interior atom separately (e.g., glycine example: N = trigonal pyramidal, left C tetrahedral, etc.)

Molecular Polarity (Net Dipole)

Procedure

1. Draw Lewis & identify shape (VSEPR).

2. Mark each polar bond with vector toward more EN atom.

3. Add vectors (vector addition) → resultant $\vec \mu$

   • If vectors cancel → nonpolar.

   • If resultant ≠ 0 → polar.
     Special cases

• Trigonal planar & tetrahedral with identical surrounding atoms = nonpolar despite polar bonds.

• Bent, trigonal pyramidal etc. usually polar.

Effect on solubility

• Polar molecules & ions dissolve in polar solvents (e.g., water). Oil (non-polar) immiscible with water.

Representing 3-D Structures on Paper

• Straight line = bond in plane

• Solid wedge = bond coming out

• Hashed wedge = bond going behind

Limitations of Lewis Model

• Qualitative; cannot yield precise bond energies/lengths or exact bond angles

• Cannot handle extensive delocalisation (requires resonance hybrids & MO theory)

• Incorrect magnetic predictions (e.g., predicts $\text O_2$ diamagnetic; experimentally paramagnetic)

Key Equations & Data (selection)

• Coulomb’s Law: $E \propto \dfrac{q<em>1 q</em>2}{r}$

• Formal charge: $\text{FC}= \text{VE} - [\text{lone} + \tfrac12\text{shared}]$

• Dipole moment: $\mu=qr$

• Reaction enthalpy (bond energies): $\Delta H=\sum D<em>{broken}-\sum D</em>{formed}$

• Example lattice energies (kJ·mol$^{-1}$):
  LiCl −834, NaCl −788, KCl −701, CsCl −657

Practice/Conceptual Questions Highlighted in Slides

(answers provided in original slides)

• Identify ionic compound among CH₄, N₂O, MgF₂ (→ MgF₂)

• Lewis symbol for Si

• Higher $T_m$: NaCl vs. MgO (→ MgO, due to 2+ & 2− charges)

• Arrange EN: P, Na, N, Al (→ N > P > Al > Na)

• Free-radical molecule? (CO, CO₂, N₂O, ClO → ClO)

• Expanded octet candidate? (H₂CO₃, H₃PO₄, HNO₂ → H₃PO₄)

• Geometry with two double bonds/no lone pairs (→ linear)

• Molecular geometry for central atom with 3 BP + 1 LP (→ trigonal pyramidal)

• Location of 2 LPs in hexagonal planar (positions opposite: 1 & 4)

• Always true VSEPR statement: shape determined by repulsions among all electron groups.

End of Notes.