Polarity & Intermolecular Forces

Bond Polarity vs. Molecular Polarity

Polarity is the uneven distribution of charge in a bond or molecule. It comes from differences in electronegativity (the Properties of Selected Elements table, legacy Table S). There are two levels to keep straight on the NY Regents exam: the polarity of a single bond and the polarity of the whole molecule.

Polar and Nonpolar Bonds

• A nonpolar covalent bond shares electrons equally; the electronegativity difference is near zero (0 to about 0.4). Example: the H-H bond in H2 or the bond between two identical atoms.
• A polar covalent bond shares electrons unequally because one atom is more electronegative. The more electronegative atom takes on a partial negative charge (written as the Greek delta-minus) and the other a partial positive charge. This separation is called a dipole. Example: H-Cl, where Cl pulls electrons toward itself.

Polar and Nonpolar Molecules

A molecule can contain polar bonds yet still be nonpolar overall if its shape is symmetrical so the dipoles cancel. Molecular shape is therefore essential.

Water (H2O) is the headline example: its bent shape means the two O-H dipoles do not cancel, so water is a polar molecule with a negative oxygen end and positive hydrogen ends. This polarity makes water an excellent solvent ("like dissolves like": polar solvents dissolve polar and ionic solutes, while nonpolar solvents dissolve nonpolar solutes such as oils).

To decide molecular polarity on the exam, ask two questions: (1) Are the bonds polar (different electronegativities)? (2) Is the shape symmetrical so the dipoles cancel? Only when bonds are polar and the shape is asymmetrical is the whole molecule polar. A diatomic molecule of two identical atoms (such as O2 or N2) is always nonpolar because its single bond is nonpolar.

Intermolecular Forces (IMFs)

Intermolecular forces are attractions between separate molecules. They are much weaker than the intramolecular covalent bonds within a molecule. The key idea for the Regents exam: IMFs determine physical properties such as boiling point, melting point, and vapor pressure, while the covalent bonds inside the molecule stay intact during melting or boiling. To boil water you overcome the IMFs between molecules, not the O-H bonds inside each molecule.

There are three IMFs, from weakest to strongest:

1. Dispersion forces (London forces) - the weakest, caused by temporary, random shifts in electron clouds. They exist between all molecules and are the only IMF in nonpolar substances. They grow stronger as molecules get larger (more electrons), which is why I2 is a solid but F2 is a gas.
2. Dipole-dipole attractions - between the permanent dipoles of polar molecules; the positive end of one molecule attracts the negative end of another.
3. Hydrogen bonding - the strongest IMF, a special, extra-strong dipole attraction that occurs only when hydrogen is bonded directly to nitrogen, oxygen, or fluorine (N, O, or F). The small, exposed hydrogen and the highly electronegative N/O/F create a powerful attraction.

To rank IMF strength for a set of molecules, first check for an N-H, O-H, or F-H bond (hydrogen bonding). If none, check whether the molecule is polar (dipole-dipole). If it is nonpolar, only dispersion forces act, and you then compare molecular size. Larger nonpolar molecules with more electrons have stronger dispersion forces, which is why the halogens go from gas (F2, Cl2) to liquid (Br2) to solid (I2) as you move down Group 17.

Why Hydrogen Bonding Matters

Hydrogen bonding explains several facts the Regents exam loves:

• Water's high boiling point (100 degrees C) is far higher than similar-sized molecules like H2S, because hydrogen bonds between water molecules take extra energy to break.
• Ice floats because hydrogen bonds hold water molecules in an open, hexagonal lattice when frozen, making solid water less dense than liquid water - rare among substances.
• Surface tension and water's role as the "universal solvent" also trace back to hydrogen bonding.

Comparing Boiling Points: A Worked Approach

Question: Why does water (H2O) boil at a much higher temperature than methane (CH4), even though both are small molecules?

1. Identify the IMF in each. CH4 is nonpolar (symmetrical tetrahedral) - only dispersion forces. H2O is polar with H bonded to O - hydrogen bonding.
2. Rank the strength: hydrogen bonding is much stronger than dispersion forces.
3. Stronger IMFs require more energy to separate molecules, so water boils much higher.

Common Exam Traps

• Confusing intermolecular and intramolecular forces. Boiling breaks IMFs (between molecules), not the covalent bonds inside.
• Forgetting shape. CO2 has polar bonds but is nonpolar because it is linear and symmetrical.
• Misapplying hydrogen bonding. It requires H bonded to N, O, or F - not just any hydrogen.
• Ignoring molecular size in dispersion forces. Larger molecules have stronger dispersion forces and higher boiling points.
• Saying ice sinks. Ice floats because hydrogen bonding makes solid water less dense than liquid water.

Link electronegativity to bond polarity, bond polarity plus shape to molecular polarity, and molecular polarity to IMF strength, and you can predict boiling points and solubility on demand.