4 Molecular Structure and Polarity

Predicting Molecular Shape

Molecules are three-dimensional, so a flat Lewis structure does not by itself show their full shape. To predict shape, first count the regions of electron density around the central atom. Each bond counts as one region, even if it is a double or triple bond, and each lone pair counts as one region.

VSEPR theory predicts that these regions arrange themselves as far apart as possible. In the notation AXₘEₙ, A is the central atom, X represents atoms bonded to it, and E represents lone pairs on it.

Electron-pair geometry describes the positions of all regions around the central atom. Molecular geometry describes the positions of the atoms only; lone pairs affect the shape but are not included in its name.

Common arrangements include:

• Two regions: linear electron-pair geometry; AX₂ is linear, with an ideal angle of $180^\circ$.

• Three regions: trigonal planar electron-pair geometry; AX₃ is trigonal planar, while AX₂E is bent, with an ideal angle of about $120^\circ$.

• Four regions: tetrahedral electron-pair geometry; AX₄ is tetrahedral, AX₃E is trigonal pyramidal, and AX₂E₂ is bent, with ideal angles of about $109.5^\circ$.

• Five regions: trigonal bipyramidal electron-pair geometry. Examples include AX₅ (trigonal bipyramidal), AX₄E (seesaw), AX₃E₂ (T-shaped), and AX₂E₃ (linear). Ideal reference angles include $90^\circ$, $120^\circ$, and $180^\circ$.

• Six regions: octahedral electron-pair geometry. Examples include AX₆ (octahedral), AX₅E (square pyramidal), and AX₄E₂ (square planar). Ideal reference angles include $90^\circ$ and $180^\circ$.

These angles are reference values; actual angles can differ. Lone pairs generally repel more strongly than bonding regions and can compress nearby bond angles. For example, $\mathrm{CH_4}$, $\mathrm{NH_3}$, and $\mathrm{H_2O}$ each have four electron-density regions around the central atom. They have different molecular geometries because they have different numbers of lone pairs.

Takeaway: Count bonding and lone-pair regions to find the electron-pair geometry, then use the atom positions and lone pairs to determine the molecular geometry.

From Bond Polarity to Molecular Polarity

A covalent bond is polar when its atoms attract the shared electrons unequally. The more electronegative atom carries a partial negative charge, written $\delta^-$, while the other carries a partial positive charge, written $\delta^+$. A bond dipole represents this separation of charge.

To determine molecular polarity, consider both the bond dipoles and the three-dimensional molecular shape. Because dipoles have direction, they add as vectors. If their vector sum is nonzero, the molecule is polar; if they cancel, the molecule is nonpolar.

Examples show why polar bonds alone do not guarantee a polar molecule:

• $\mathrm{CO_2}$ has polar C=O bonds, but its linear shape points the bond dipoles in opposite directions, so they cancel. The molecule is nonpolar.

• $\mathrm{H_2O}$ has polar O–H bonds and a bent shape, so its dipoles do not cancel. The molecule is polar.

• $\mathrm{CCl_4}$ has polar C–Cl bonds arranged symmetrically in a tetrahedron. Their dipoles cancel, so the molecule is nonpolar.

• $\mathrm{CH_3Cl}$ has a tetrahedral arrangement, but one substituent differs from the other three. Its bond dipoles do not cancel, so the molecule is polar.

Takeaway: Identify polar bonds, then consider whether their directions and arrangement cause the dipoles to cancel.

How Molecular Structure Affects Properties

A molecule’s shape and polarity affect where its electrons and partial charges are located. These features influence intermolecular attractions, which in turn affect properties such as boiling point and solubility.

Polar molecules can attract one another through Dipole–dipole forces. All molecules also experience London dispersion forces, which tend to become stronger as electron clouds become larger and more easily distorted. When hydrogen is directly bonded to N, O, or F, molecules can also form Hydrogen bonds, a particularly strong type of dipole–dipole attraction.

Stronger intermolecular attractions generally require more energy to overcome. When other factors are similar, stronger attractions often correspond to higher boiling points. A comparison of molecules with similar molar masses illustrates this pattern:

• Propane, $\mathrm{C_3H_8}$, is nonpolar and mainly experiences dispersion forces.

• Dimethyl ether, $\mathrm{CH_3OCH_3}$, is polar and also experiences dipole–dipole attractions.

• Ethanol, $\mathrm{C_2H_5OH}$, can form hydrogen bonds because it has an O–H bond.

Polarity also helps explain solubility. Substances tend to dissolve better when their particles can form favorable attractions with the solvent. This is a useful guideline, not an absolute rule: size, shape, and the strengths of all attractions involved also matter.

Takeaway: Shape and polarity influence the attractions between particles, but boiling point and solubility depend on the combined effects of those attractions and other factors.