Chemical Bonding: Ionic, Covalent, Metallic, and Intermolecular Forces

Electronegativity Determines Everything

The nature of a chemical bond between two atoms is not a discrete category — it is a point on a continuous spectrum determined by the difference in electronegativity between the bonded atoms. Electronegativity, a concept formalized by Linus Pauling in 1932, measures the tendency of an atom to attract electrons in a chemical bond. The Pauling scale runs from 0.7 (francium) to 3.98 (fluorine). When two atoms form a bond, the more electronegative atom pulls the shared electron density toward itself — and the magnitude of this pull determines the bond character.

• Electronegativity difference < 0.5: Nonpolar covalent bond — electrons shared nearly equally (H-H, C-H).
• Electronegativity difference 0.5–1.7: Polar covalent bond — electrons shared unequally, creating a partial dipole (O-H, N-H, C-O).
• Electronegativity difference > 1.7: Ionic bond — electron transfer produces discrete ions (NaCl: Na 0.93, Cl 3.16, ΔEN = 2.23).

These thresholds are approximate. Chemical bonds rarely fit neatly into categories. Bond character is always on a spectrum.

Ionic Bonds and Lattice Energy

Ionic bonds form through electron transfer from a metal to a nonmetal, producing oppositely charged ions held together by electrostatic attraction. Sodium chloride forms when sodium transfers its sole valence electron to chlorine: Na → Na+ + e−; Cl + e− → Cl−. The resulting ions adopt a face-centered cubic lattice structure in which each sodium ion is surrounded by six chloride ions and vice versa.

Lattice energy — the energy released when gaseous ions combine to form an ionic crystal — is quantified through the Born-Haber cycle, which applies Hess's law to the sequential steps from elemental sodium and chlorine to NaCl(s). Lattice energies are large negative values (NaCl: −787 kJ/mol) reflecting the strong electrostatic attractions in ionic solids. Higher charges and smaller ionic radii produce larger lattice energies: MgO (Mg2+, O2−) has a lattice energy of −3791 kJ/mol, nearly five times that of NaCl. This explains MgO's far higher melting point (2852°C vs. 801°C for NaCl).

Covalent Bond Energy and Length

In covalent bonds, atoms share electrons rather than transfer them. Bond strength and length depend on the bond order (single, double, or triple) and the sizes of the bonded atoms.

Higher bond order produces shorter, stronger bonds — a direct consequence of the increased electron density between nuclei. The N≡N bond (945 kJ/mol) is among the strongest in chemistry, which is why nitrogen gas is so unreactive and why nitrogen fixation — breaking N≡N to make fertilizer — requires such extreme industrial conditions.

VSEPR: Predicting Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory, developed by Gillespie and Nyholm in 1957, predicts molecular geometry from the principle that electron pairs (bonding and lone) adopt positions that minimize repulsion. The geometry reflects the number of electron domains around the central atom.

Lone pairs repel more strongly than bonding pairs, compressing bond angles: NH3 (107°, not 109.5°) and H2O (104.5°) demonstrate this compression relative to the ideal tetrahedral angle.

Hybrid Orbital Model

Hybrid orbitals explain why carbon forms four equivalent bonds in methane despite having only two unpaired electrons in its ground state. Valence bond theory proposes that atomic orbitals hybridize — mathematically mix — to form new equivalent orbitals optimized for bonding:

• sp hybridization: one s + one p orbital → two sp orbitals at 180°. Remaining two p orbitals form two π bonds. Result: linear geometry with one sigma and two pi bonds (alkynes, CO2, BeH2).
• sp2 hybridization: one s + two p orbitals → three sp2 orbitals at 120°. One remaining p orbital forms one π bond. Result: trigonal planar, one double bond character (alkenes, BF3, benzene's delocalized system).
• sp3 hybridization: one s + three p orbitals → four equivalent sp3 orbitals at 109.5°. Result: tetrahedral geometry, all single bonds (alkanes, water, ammonia).

Metallic Bonding: The Electron Sea

In metals, valence electrons are not localized between specific atom pairs but delocalize throughout the entire crystal structure — forming a "sea" of mobile electrons surrounding fixed positive metal ion cores. This electron sea model explains metallic properties: electrical conductivity (electrons flow freely under applied voltage), thermal conductivity (electrons transfer kinetic energy rapidly), malleability and ductility (ion cores can shift positions without breaking bonds), and metallic luster (free electrons absorb and re-emit visible light). Transition metals with partially filled d orbitals show the strongest metallic bonding — tungsten's delocalization across both s and d electrons explains its record-setting melting point of 3422°C.

Intermolecular Forces: The Hidden Hierarchy

Intermolecular forces — attractions between molecules rather than within them — determine boiling points, solubility, surface tension, and viscosity.

Water's anomalous properties follow directly from its hydrogen bonding network. Water's boiling point (100°C) is dramatically higher than predicted for a molecule of its size by London dispersion alone — H2S, the analog with only dispersion forces, boils at −60°C. Each water molecule forms up to four hydrogen bonds (two donated from O-H groups, two accepted at the lone pairs of oxygen), creating a dynamic, structured liquid with high surface tension, high heat capacity, and solid form (ice) that is less dense than the liquid — properties essential to life on Earth. Hydrogen bonds are 10× weaker than covalent bonds. But their collective effect shapes everything.

Coordinate Covalent (Dative) Bonds

In conventional covalent bonds, each atom contributes one electron to the shared pair. In coordinate covalent (dative) bonds, both electrons in the shared pair come from a single atom — the donor — to an atom with an empty orbital — the acceptor. The resulting bond is indistinguishable from a conventional covalent bond once formed; only the origin of electrons differs. Lewis acid-base chemistry is built on coordinate bonding: BF3 (Lewis acid, empty p orbital) accepts an electron pair from NH3 (Lewis base, lone pair) to form F3B:NH3. Transition metal coordination compounds — where ligands donate lone pairs to metal centers — are an entire chemistry domain founded on coordinate bonding.