8.3 Chemical Bonding & Molecular Properties

The Octet Rule and Bonding Driving Forces

Chemical bonding is the fundamental process by which atoms combine to form stable compounds. The primary driving force behind chemical bond formation is the lowering of potential energy. Unbonded atoms possess relatively high potential energy; by forming bonds, atoms attain a more stable, lower-energy electronic arrangement.

Central to chemical bonding is the Octet Rule. The octet rule states that main-group atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons, achieving the exceptionally stable, filled s- and p-subshell configuration of a noble gas (or two electrons for hydrogen and helium).

Chemical forces are divided into two primary categories:

1. Intramolecular Forces: Strong chemical bonds (ionic, covalent, metallic) holding individual atoms together within a molecule or compound.
2. Intermolecular Forces (IMFs): Weaker attractive forces operating between neighboring molecules.

---

Ionic Bonding: Electron Transfer and Crystal Lattices

Ionic bonding occurs when valence electrons are completely transferred from one atom to another, typically between a metal (low electronegativity, low ionization energy) and a nonmetal (high electronegativity, high electron affinity). The large electronegativity difference ($\Delta\text{EN} > 1.7$) drives complete electron transfer:

• The metal atom loses one or more electrons to become a positively charged cation.
• The nonmetal atom gains those electrons to become a negatively charged anion.

Crystal Lattice and Physical Properties

Oppositely charged cations and anions attract each other through powerful, non-directional electrostatic (Coulombic) forces. Rather than forming isolated molecules, ionic compounds organize into a rigid, repeating three-dimensional crystal lattice structure (e.g., $NaCl$, $CaCl_2$).

Properties of ionic compounds include:

• High Melting and Boiling Points: Immense thermal energy is required to disrupt the strong electrostatic forces within the crystal lattice.
• Brittleness: Mechanical stress shifts lattice layers, placing like-charged ions adjacent to one another, causing strong repulsion and shattering.
• Electrical Conductivity: Solid ionic crystals cannot conduct electricity because ions are locked in place. However, when molten or dissolved in water (forming an aqueous solution), ions dissociate and move freely, acting as strong electrolytes.

---

Covalent Bonding: Electron Sharing and Polarity

Covalent bonding occurs when two nonmetal atoms share one or more pairs of valence electrons to achieve mutual octets. Covalent bonds form between atoms with similar, high electronegativities ($\Delta\text{EN} < 1.7$).

Bond Orders

• Single Covalent Bond: 1 shared pair of electrons ($2\text{ e}^-$ total, e.g., $H_2$, $H-Cl$).
• Double Covalent Bond: 2 shared pairs of electrons ($4\text{ e}^-$ total, e.g., $O_2$, $O=C=O$).
• Triple Covalent Bond: 3 shared pairs of electrons ($6\text{ e}^-$ total, e.g., $N_2$, $N\equiv N$).

As bond order increases from single to triple, bond length decreases and bond strength (bond dissociation energy) increases.

Electronegativity Difference ($\Delta\text{EN}$) and Bond Polarity

Electronegativity differences dictate how evenly shared electrons are distributed between bonded atoms:

1. Non-Polar Covalent Bonds ($\Delta\text{EN} < 0.5$): Electrons are shared equally between atoms with identical or nearly identical electronegativities (e.g., $H_2$, $O_2$, $C-H$ bonds in methane $CH_4$). No partial charges develop.
2. Polar Covalent Bonds ($0.5 \le \Delta\text{EN} \le 1.7$): Shared electrons are pulled closer to the more electronegative atom. This unequal electron distribution creates a bond dipole, giving the more electronegative atom a partial negative charge ($\delta^-$) and the less electronegative atom a partial positive charge ($\delta^+$) (e.g., $H-O$ in water, $H-Cl$).

---

Metallic Bonding: The Electron Sea Model

Metallic bonding accounts for the unique properties of pure metals and metal alloys. Metal atoms possess low electronegativities and loosely held valence electrons. In metallic bonding, positive metal cations are fixed in a lattice while their valence electrons become delocalized, moving freely throughout the entire structure in an "electron sea."

The mobile electron sea accounts for key metallic properties:

• High Electrical and Thermal Conductivity: Delocalized electrons move rapidly under electric fields or thermal gradients.
• Malleability and Ductility: Metal cations can slide past each other without shattering because the sea of electrons continuously adjusts to maintain bonding.

---

Lewis Dot Structures and VSEPR Theory

Lewis dot structures visually represent molecules by displaying valence electrons as dots surrounding atomic symbols. Single lines represent shared bonding electron pairs, while non-bonding pairs are drawn as lone pairs.

VSEPR Theory (Valence Shell Electron Pair Repulsion)

VSEPR theory predicts the three-dimensional geometric shapes of molecules. VSEPR states that electron pairs surrounding a central atom—both bonding pairs and non-bonding lone pairs—repel one another electrostatically and arrange themselves as far apart in space as possible.

Common VSEPR Geometries:

• Linear: 2 bonding domains, 0 lone pairs around central atom; bond angle $180^\circ$ (e.g., $CO_2$, $BeF_2$).
• Trigonal Planar: 3 bonding domains, 0 lone pairs; bond angle $120^\circ$ (e.g., $BF_3$).
• Tetrahedral: 4 bonding domains, 0 lone pairs; bond angle $109.5^\circ$ (e.g., $CH_4$).
• Trigonal Pyramidal: 3 bonding domains, 1 lone pair; bond angle $\approx 107^\circ$ (e.g., $NH_3$). Lone pair repulsion compresses bonding angles.
• Bent (Angular): 2 bonding domains, 2 lone pairs; bond angle $104.5^\circ$ (e.g., $H_2O$).

---

Intermolecular Forces (IMFs) and Physical Properties

Intermolecular forces are attractive forces between independent covalent molecules. Although significantly weaker than intramolecular chemical bonds, IMFs determine physical properties such as melting point, boiling point, viscosity, and surface tension.

The Three Primary IMFs

1. London Dispersion Forces (LDF): Temporary, instantaneous dipoles created by random fluctuations in electron density. LDFs exist between all molecules (both polar and non-polar). LDFs are the weakest IMF, but their strength increases with larger electron clouds (greater polarizability / higher molar mass).
2. Dipole-Dipole Interactions: Electrostatic attraction between the permanent partial positive ($\delta^+$) pole of one polar molecule and the partial negative ($\delta^-$) pole of an adjacent polar molecule (e.g., $HCl$).
3. Hydrogen Bonding: An extraordinarily strong subtype of dipole-dipole force occurring strictly when hydrogen is directly bonded to a small, highly electronegative atom: Nitrogen ($N$), Oxygen ($O$), or Fluorine ($F$). Hydrogen bonding gives water ($H_2O$) its high boiling point ($100^\circ\text{C}$), high surface tension, and expanded solid ice density.