2.12 Noncovalent Interactions between Molecules

Noncovalent Interactions between Molecules

Noncovalent interactions are the forces that hold molecules near each other without sharing electrons in a bond. They're much weaker than covalent bonds individually, but collectively they determine boiling points, solubility, and the shapes of biological molecules like proteins and DNA. In organic chemistry, recognizing these forces helps you predict physical properties and understand why molecules behave the way they do in solution.

Types of Noncovalent Molecular Interactions

Three main types of noncovalent forces show up repeatedly in organic chemistry. They differ in origin and strength, but all arise from electrostatic attractions.

Dipole-dipole forces occur between polar molecules that have permanent dipoles. The partially positive end of one molecule is attracted to the partially negative end of another. For example, HCl molecules align so that the $\delta^+$ hydrogen of one molecule faces the $\delta^-$ chlorine of a neighbor. The strength of these interactions depends on the size of the dipole moments and how well the molecules can align.

Dispersion forces (London forces) are present in all molecules, polar or not. They arise because electrons are constantly moving, which creates temporary, uneven distributions of charge. That momentary dipole in one molecule induces a complementary dipole in a neighboring molecule, producing a brief attraction. A few things to remember:

• Dispersion forces increase with molecular size (more electrons = more polarizable)
• They also increase with surface area. Long, straight-chain alkanes have stronger dispersion forces than compact, branched isomers because more of the molecule's surface can interact with neighbors
• For small nonpolar molecules like $CH_4$, dispersion forces are the only intermolecular force present

Hydrogen bonds are a special, stronger type of dipole-dipole interaction. They form when a hydrogen atom bonded to a highly electronegative atom ($N$, $O$, or $F$) interacts with a lone pair on another electronegative atom. The hydrogen in this arrangement carries a large $\delta^+$ because $N$, $O$, and $F$ pull electron density away from it so effectively. Hydrogen bonds are stronger than typical dipole-dipole forces but still much weaker than covalent bonds (roughly 5–25 kJ/mol vs. 350–400+ kJ/mol for a covalent bond).

Intermolecular Forces and Molecular Properties

Van der Waals forces is a collective term that includes dipole-dipole interactions, dispersion forces, and induced dipole interactions. You'll sometimes see this term used loosely, so know that it covers all the weak intermolecular attractions except ionic bonds and hydrogen bonds (though some textbooks include hydrogen bonds under this umbrella).

Polarity of a molecule determines which types of intermolecular forces dominate. A molecule with a net dipole moment will experience dipole-dipole forces on top of dispersion forces, giving it higher boiling points and greater solubility in polar solvents compared to a nonpolar molecule of similar size.

Induced dipoles explain how polar and nonpolar molecules can still attract each other. A polar molecule's electric field distorts the electron cloud of a nearby nonpolar molecule, creating a temporary dipole in it. This is why nonpolar gases like $O_2$ still dissolve (slightly) in water.

Effects of Hydrogen Bonding

Hydrogen bonding has outsized effects on physical properties, especially in water and biological molecules.

Water is the classic example. Each water molecule can form up to four hydrogen bonds (two through its hydrogens, two through its oxygen lone pairs). This extensive hydrogen bonding network explains several unusual properties:

1. High surface tension from strong cohesive forces between surface molecules
2. High specific heat capacity, meaning water resists temperature changes because energy goes into disrupting hydrogen bonds before raising the temperature
3. Ice is less dense than liquid water because hydrogen bonds lock water molecules into an open, hexagonal lattice in the solid state, taking up more volume than the liquid

DNA relies on hydrogen bonding for both structure and function. Adenine pairs with thymine through two hydrogen bonds (A–T), while guanine pairs with cytosine through three (G–C). These complementary base pairs hold the two strands of the double helix together. The specificity of this pairing is what makes accurate DNA replication and transcription possible. Similar hydrogen bonding stabilizes secondary structures in RNA, such as hairpin loops and stem-loops.

Hydrophilic vs. Hydrophobic Substances

The concept of "like dissolves like" comes down to noncovalent interactions. Substances interact best with solvents that offer similar types of intermolecular forces.

• Hydrophilic ("water-loving") substances contain polar or charged functional groups that can form hydrogen bonds with water. Examples: sugars (many $OH$ groups), salts (ionic), ethanol.
• Hydrophobic ("water-fearing") substances are nonpolar and lack groups capable of hydrogen bonding with water. Examples: alkanes like hexane, fats, oils. These molecules interact with each other primarily through dispersion forces.
• Amphiphilic substances have both a hydrophilic region and a hydrophobic region in the same molecule. Phospholipids are a key example: their polar head groups face water while their nonpolar fatty acid tails cluster together, forming the bilayer structure of cell membranes. Soaps work on the same principle, surrounding grease droplets with their nonpolar tails while their polar heads face the surrounding water.

Solvation is the process by which solvent molecules surround and stabilize a dissolved solute. When you dissolve NaCl in water, water molecules orient around each ion ($\delta^-$ oxygens face $Na^+$, $\delta^+$ hydrogens face $Cl^-$). For organic molecules, solvation depends on matching the polarity of the solute with the solvent. This is why you'll see polar reactions run in polar solvents and nonpolar compounds extracted with nonpolar solvents.