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5B · The chemistry of life
Molecules and intermolecular interactions
How atoms bond into molecules (ionic, covalent, polarity, Lewis structures, VSEPR shape, hybridization) and how molecules stick to each other (the intermolecular forces that set boiling point, melting point, and solubility).
Chemical bonding and molecular structure
Bonds range from ionic (electron transfer, large electronegativity difference) to polar covalent to nonpolar covalent (equal sharing). Lewis structures + VSEPR give geometry; hybridization (sp, sp², sp³) follows from the number of electron domains.
Electronegativity difference sets bond character: large → ionic, intermediate → polar covalent, near-zero → nonpolar covalent. After drawing a valid Lewis structure (octets, minimized formal charge, resonance where delocalized), count electron domains around the central atom: VSEPR predicts the geometry (linear, trigonal planar, tetrahedral, etc.) and the matching hybridization. Molecular polarity then depends on whether bond dipoles cancel by symmetry (CO₂ is nonpolar despite polar bonds; H₂O is polar).
Ionic, covalent, and metallic bonds
Ionic: electrons transferred (metal + nonmetal), forming a lattice. Covalent: electrons shared (nonmetals); polar if shared unequally. Metallic: cations in a "sea" of delocalized electrons (conductivity, malleability).
Ionic compounds are high-melting, brittle, and conduct only when molten or dissolved (mobile ions). Covalent molecular substances are lower-melting with properties set by IMFs. Bond polarity within a covalent bond grows with the electronegativity difference, producing a dipole that may or may not cancel at the molecular level depending on geometry.
Lewis structures, formal charge, and resonance
Draw Lewis structures to satisfy octets and minimize formal charge; when more than one valid structure exists, the true picture is a resonance hybrid (delocalized electrons), more stable than any single form.
Formal charge = valence electrons − (lone-pair electrons + ½ bonding electrons); the best structure puts negative formal charge on the most electronegative atom and keeps charges minimal. Resonance delocalizes electrons over several atoms (carboxylate, benzene, the peptide bond), conferring extra stability and partial-bond character. Recognizing resonance explains acidity (a stabilized conjugate base is a stronger acid) and reactivity throughout organic chemistry. Watch the octet exceptions: hydrogen wants only 2 electrons; boron and beryllium are often electron-deficient (6 or 4); and period-3+ atoms (P, S, Cl) can hold expanded octets of 10–12 electrons.
VSEPR geometry and hybridization
Electron domains repel to maximize separation: 2 → linear (sp), 3 → trigonal planar (sp²), 4 → tetrahedral (sp³). Lone pairs occupy domains and bend the molecular shape (e.g., H₂O is bent, not linear).
Count all electron domains (bonds — single/double/triple each count once — plus lone pairs) to get the electron geometry and hybridization. The molecular geometry then describes just the atoms: with lone pairs present it differs from the electron geometry (NH₃ is trigonal pyramidal; H₂O bent). Hybridization tracks domain count: 2 = sp, 3 = sp², 4 = sp³. Double/triple bonds are one sigma plus one/two pi bonds.
Don't confuse
A molecule with polar bonds can still be nonpolar if symmetry cancels the dipoles (CO₂, CCl₄). Reading polarity off bond polarity alone — ignoring geometry — is the classic VSEPR trap.
Intermolecular forces
Ranked weakest → strongest: London dispersion (all molecules) < dipole–dipole (polar molecules) < hydrogen bonding (H to N/O/F) — with ion–dipole strongest of all (in solutions). Stronger IMFs mean higher boiling/melting points and shape solubility.
IMFs are attractions between molecules, far weaker than the covalent bonds within them, yet they dictate phase behavior. London dispersion forces exist in everything and grow with molecular size/surface area (more electrons → more polarizability), so larger nonpolar molecules boil higher. Dipole–dipole adds for polar molecules; hydrogen bonding is the strongest neutral IMF. Boiling point is essentially a contest of IMF strength; solubility follows "like dissolves like" (polar/H-bonding solutes in water, nonpolar in nonpolar solvents).
London, dipole–dipole, and hydrogen bonding
London dispersion: temporary induced dipoles, in all molecules, stronger for bigger/more-polarizable ones. Dipole–dipole: permanent dipoles align. Hydrogen bonding: H–N/O/F to a lone pair, the strongest neutral IMF.
To rank boiling points, find the strongest IMF present, then break ties by size (dispersion). Hydrogen bonding explains water's anomalously high boiling point versus H₂S, and why alcohols boil far above comparable alkanes. Among nonpolar molecules, the larger one always boils higher (more dispersion) — which is why hydrocarbon boiling points rise with chain length.
Don't confuse
Hydrogen bonding is an intermolecular force, not a covalent bond. The exam contrasts "breaking hydrogen bonds" (melting/boiling, denaturation — weak, reversible) with "breaking covalent bonds" (chemical reaction — far more energy). Conflating the two misranks energies badly.
IMFs and physical properties
Stronger IMFs raise boiling/melting point, viscosity, and surface tension, and lower vapor pressure. Solubility follows polarity matching ("like dissolves like").
Vapor pressure is inversely related to IMF strength (strong IMFs hold molecules in the liquid, so fewer escape) — and a liquid boils when its vapor pressure equals atmospheric pressure, which is why high-IMF liquids boil higher. Volatility, evaporative cooling, and how readily a compound partitions between water and an organic layer (key in 5C) all trace back to IMFs.
Phases and phase changes
Stronger intermolecular forces raise melting and boiling points. During a phase change the temperature stays constant (added heat breaks IMFs rather than raising kinetic energy). A phase diagram maps which phase is stable at each temperature and pressure.
Heating a solid raises its temperature until a phase change, where the temperature plateaus while heat goes into overcoming intermolecular forces (the latent heat of fusion or vaporization) — only once the change completes does temperature climb again. A phase diagram (pressure vs. temperature) shows the solid/liquid/gas regions, the lines where two phases coexist, the triple point (all three coexist), and the critical point (beyond which liquid and gas become indistinguishable). Water's solid–liquid boundary slopes negative (ice is less dense than liquid water), a notable exception.
Related
The heat absorbed at the temperature plateau is the latent heat quantified in calorimetry.
Worked question
Which of the following pure substances is expected to have the highest boiling point: ethane (C₂H₆), fluoromethane (CH₃F), ethanol (C₂H₅OH), or propane (C₃H₈)?