Organic Chemistry · Polar Covalent Bonds; Acids and Bases
Noncovalent Interactions between Molecules
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In 30 seconds
Molecules attract one another through a set of weaker forces — the noncovalent (intermolecular) interactions — that act between molecules rather than within them. There are four main types: ion–dipole, dipole–dipole, hydrogen bonding, and London dispersion forces (plus ion–ion interactions between charged species). Each is electrostatic in origin: interactions between permanent charges, permanent dipoles, and induced (temporary) dipoles. Individually these forces are far weaker than covalent bonds — a typical Hydrogen bond Strong dipole–dipole attraction between X-H (X = N, O, F) and a lone pair Full entry → is about 20 kJ/mol versus roughly 350 kJ/mol for a C–C bond — but they act between every pair of molecules and add up. Noncovalent interactions determine boiling points, melting points, solubility, surface tension, and viscosity, and they are the forces that hold proteins folded, DNA base-paired, and drugs bound to their receptors. This topic closes Chapter 2 by connecting polarity (the dipole moments of earlier topics) to the observable physical properties of organic compounds.
Why this matters
- Physical properties: whether a substance is a gas, liquid, or solid at room temperature is decided by its intermolecular forces. Butane (dispersion only) is a gas; butanol (hydrogen bonding) is a liquid that boils near 118 °C.
- Solubility and "Like dissolves like Rule that solubility matches polarity Full entry →": predicting whether a drug dissolves in water or crosses a lipid membrane requires judging which noncovalent interactions dominate.
- Biomolecular structure: protein folding, the DNA double helix, enzyme–substrate binding, and antibody–antigen recognition are all governed by hydrogen bonds, ionic interactions, and dispersion forces.
- Drug design and pharmacology: binding of a drug to a receptor is a balance of noncovalent interactions — stronger, more specific binding means lower dose and fewer side effects.
- Separation science: chromatography, distillation, and recrystallization all exploit differences in intermolecular forces.
The college version
Core Concepts
The four main noncovalent interactions, in order of typical strength
| Interaction | Typical strength (kJ/mol) | What attracts what |
|---|---|---|
| Ion–ion | 100–350+ | Full positive and negative charges |
| Ion–dipole | 40–600 (gas phase); ~10–50 in solution | An ion and a polar molecule's dipole |
| Hydrogen bond | 10–40 | A H bonded to N/O/F and a lone pair on N/O/F |
| Dipole–dipole | 5–25 | The positive end of one dipole and negative end of another |
| London dispersion | 0.05–40 (grows with size) | Temporary (induced) dipoles in any molecules |
(Strengths are typical textbook ranges; actual values depend strongly on distance, geometry, and medium.)
Hydrogen bonding: the special dipole–dipole case
A hydrogen bond is a strong Dipole–dipole interaction Attraction between the opposite ends of two molecular dipoles Full entry → between a hydrogen covalently bonded to a highly electronegative atom — N, O, or F — and a lone pair on another N, O, or F atom. The hydrogen acts as a "bridge": O-H ⋯O. Two conditions are required: a donor (X-H where X = N, O, F) and an acceptor (a lone pair on N, O, or F). Water, alcohols, carboxylic acids, amides, and amines all hydrogen bond; that is why ethanol (CH3CH2OH, bp 78 °C) boils far above dimethyl ether (CH3OCH3, bp −24 °C) even though both have the formula C2H6O.
London dispersion forces: universal and size-dependent
London dispersion forces arise from instantaneous fluctuations in electron distribution that create temporary dipoles, which induce opposite dipoles in neighboring molecules. They operate between all molecules, polar or not, and grow with Polarizability How easily an electron cloud distorts Full entry → — the ease with which electron clouds distort. Bigger atoms (I > Br > Cl > F) and longer carbon chains are more polarizable, which is why boiling points rise steadily along a homologous series (methane −162 °C → ethane −89 °C → propane −42 °C → butane −1 °C): each CH2 unit adds surface area and dispersion attraction.
Ion–dipole interactions and solvation
An ion dissolved in a polar solvent attracts the solvent's dipoles: the cation's positive charge draws the negative ends of nearby water molecules. This ion–dipole interaction is what pays the energy cost of breaking the ionic lattice when NaCl dissolves in water. The interaction energy follows Coulomb's law:
E = -q μ4πε0 r2
where q is the ion's charge, μ the dipole moment of the solvent molecule, r the ion–dipole separation, and ε0 the vacuum permittivity. The force falls off as 1/r2, so it is shorter-ranged than ion–ion interactions but much stronger than dipole–dipole at contact distance.
Consequences: boiling points and solubility
- Boiling point ranking logic: first count hydrogen-bond donors/acceptors; then compare dipole moments; then compare size/polarizability. More (and stronger) interactions → higher boiling point.
- "Like dissolves like": polar solvents (water, ethanol) dissolve polar and ionic solutes via dipole and ion–dipole interactions; nonpolar solvents (hexane, diethyl ether) dissolve nonpolar solutes via dispersion forces. Oil and water separate because water's hydrogen bonds to itself are far stronger than any interaction it can offer a hydrocarbon.
- Amphiphiles: molecules with both a polar head and a nonpolar tail (soaps, phospholipids) form micelles and bilayers — the basis of membranes and detergency.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Hydrogen bond | Covalent bond to hydrogen | A hydrogen bond is an intermolecular attraction (O-H ⋯O, ~20 kJ/mol); the O-H covalent bond is ~460 kJ/mol. |
| Dipole–dipole forces | Hydrogen bonding | H-bonding is a special, stronger dipole–dipole case requiring H on N/O/F plus an N/O/F lone-pair acceptor. Ethers have dipoles but no H-bond donor. |
| Dispersion forces "only for nonpolar molecules" | Dispersion forces | Dispersion operates between all molecules; it is just relatively more important for nonpolar ones. |
| Boiling point follows molar mass alone | Boiling point follows interactions | Mass matters only through polarizability; functional groups (OH, NH) can add 100+ °C regardless of mass (ethanol 46 g/mol, 78 °C vs propane 44 g/mol, −42 °C). |
| "Like dissolves like" as an absolute rule | Rule of thumb | It predicts trends, not exact solubilities; amphiphiles dissolve in both kinds of solvent via different parts of the molecule. |
| Intramolecular interactions | Intermolecular interactions | Intramolecular (within one molecule, e.g., intramolecular H-bond in salicylic acid) affects shape and properties differently from between-molecule forces. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Molecules have sticky patches, like Velcro. Some patches are strongly charged, some are slightly charged on one end, and some are just wobbly electron clouds that make neighbors wobbly too. Even the weakest stickiness — the wobbly-cloud kind — happens between all molecules, and it gets stronger as molecules get bigger. Hydrogen bonds are like super-strong Velcro between molecules that have hydrogen attached to oxygen or nitrogen. These stickiness forces decide whether something is a gas, a liquid, or a solid — and they're how your DNA stays zipped together.
Worked example
Example 1: Ranking boiling points — propane vs dimethyl ether vs ethanol
Problem: All three compounds have molar mass near 46 g/mol (C3H8, C2H6O, C2H6O). Rank their boiling points and explain. (Literature values: propane −42 °C, dimethyl ether −24 °C, ethanol 78 °C.)
Plan and formula: three checks in order — (1) can the molecule hydrogen bond? (2) does it have a permanent dipole? (3) how big/polarizable is it?
- Propane, CH3CH2CH3: no O or N, so no H-bonding; C–H bonds are nearly nonpolar, so little dipole; only dispersion forces.
- Dimethyl ether, CH3OCH3: has a permanent dipole (C–O bonds are polar; molecule is bent) → dipole–dipole in addition to dispersion. No O–H bond, so no hydrogen bonding.
- Ethanol, CH3CH2OH: same formula as dimethyl ether, but has an O–H bond → hydrogen bonding on top of dipole–dipole and dispersion.
Answer: CH3CH2OH (78 °C) > CH3OCH3 (−24 °C) > CH3CH2CH3 (−42 °C). The ~100 °C gap between ethanol and dimethyl ether — constitutional isomers with identical formulas — comes entirely from ethanol's ability to hydrogen bond.
Example 2: How strong is the ion–dipole interaction that dissolves table salt?
Problem: Estimate the energy of a single Na+–water ion–dipole contact. Use q = 1.60 × 10-19 C, the dipole moment of water μ= 1.85 D (= 6.17 × 10-30 C·m), a contact distance r = 2.4 × 10-10 m, and 1/(4πε0) = 8.99 × 109 N·m²/C².
Plan and formula:
E = -q μ4πε0 r2
Substitution:
E = -(1.60 × 10-19 C)(6.17 × 10-30 C·m)4πε0 (2.4 × 10-10 m)2
First compute the denominator: (2.4 × 10-10)2 = 5.76 × 10-20 m².
E = -9.87 × 10-495.76 × 10-20 × (8.99 × 109) = -1.54 × 10-19 J per ion–dipole pair
Convert to per mole:
E = -1.54 × 10-19 J × 6.022 × 1023 mol-1 ≈ -9.3 × 104 J/mol = -93 kJ/mol
Interpretation: a single Na⁺–water contact is worth roughly -90 kJ/mol — far larger than a typical hydrogen bond (~20 kJ/mol) and comparable to a weak covalent bond, yet still far below the ~350 kJ/mol C–C bond. Multiply this by the several water molecules surrounding each ion, and it is clear why the ion–dipole interaction can pay the lattice energy cost of dissolving salt. (This is an order-of-magnitude estimate; real solvation energies also include the reorganization of bulk water.)
Example 3: Which has the higher boiling point — butane or butanol?
Problem: Compare CH3CH2CH2CH3 (butane, molar mass 58) and CH3CH2CH2CH2OH (1-butanol, molar mass 74).
Plan and formula: apply the three-check logic. Both have four-carbon chains with similar dispersion contributions (butanol slightly larger). But 1-butanol has an O–H bond: it hydrogen bonds. Butane cannot.
Butane: dispersion only 1-butanol: dispersion + dipole–dipole + hydrogen bonding
Answer: 1-butanol boils at 118 °C; butane boils at −1 °C — a ~120 °C difference from one O–H group. This is the classic demonstration that hydrogen bonding dominates boiling point when a molecule can do it, and it explains why alcohols are liquids while alkanes of similar size are gases.
Key takeaways
- Noncovalent interactions (weak, between molecules) vs covalent bonds (strong, within molecules): H-bond ≈ 20 kJ/mol vs C–C bond ≈ 350 kJ/mol.
- Four types: ion–dipole, dipole–dipole, hydrogen bonding, London dispersion. All electrostatic; H-bonding is a strong dipole–dipole subset.
- H-bond needs a donor (X-H, X = N, O, F) and an acceptor (lone pair on N, O, F). No H on N/O/F → no hydrogen bonding.
- Dispersion forces are universal and grow with polarizability (bigger atoms, longer chains): bp rises ~20–30 °C per CH2 in a homologous series.
- Ion–dipole energy: E = -qμ/ (4πε0 r2) — the force that dissolves salts in water.
- Boiling point: hydrogen bonding > dipole–dipole > dispersion for similar-size molecules; among nonpolar molecules, bigger = higher bp.
- "Like dissolves like": polar/ionic solutes need polar solvents; nonpolar solutes need nonpolar solvents.
- Biology: protein folding, DNA base pairing, membranes, and drug–receptor binding all run on noncovalent interactions.
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
List the four main noncovalent interactions in order of typical strength, and state which one operates between all molecules.
Show answer
Ion–ion > ion–dipole > hydrogen bond > dipole–dipole > London dispersion (typical strengths: 100–350+, 10–50 in solution, 10–40, 5–25, 0.05–40 kJ/mol). London dispersion operates between all molecules.
Why does ethanol boil 100 °C higher than its isomer dimethyl ether?
Show answer
Ethanol has an O–H group, so it can donate and accept hydrogen bonds; dimethyl ether has only a dipole (C–O bonds) and no O–H, so it cannot hydrogen bond. The extra H-bonding raises ethanol's boiling point by ~100 °C.
Which is the stronger interaction per pair: a hydrogen bond (~20 kJ/mol) or a C–C covalent bond (~350 kJ/mol)? Roughly how many hydrogen bonds equal one C–C bond?
Show answer
A C–C bond is roughly 15–20 times stronger: 350/20 ≈ 17. This is why molecules held together only by noncovalent forces are easily separated (evaporation, melting) without breaking chemical bonds.
Why does the boiling point of alkanes increase by roughly 20–30 °C per additional CH2 unit?
Show answer
Each CH2 adds electron density and surface area, increasing polarizability and hence London dispersion forces; more dispersion attraction means a higher boiling point. The effect is roughly additive along the homologous series.
Water and hexane do not mix. Explain using noncovalent interactions.
Show answer
Water molecules hydrogen bond strongly to one another; hexane has no H-bond donors or acceptors and offers only weak dispersion interactions to water. Mixing would require breaking water's hydrogen bonds without replacing them with comparable interactions, which is energetically unfavorable — so the liquids separate.
Name two conditions required for hydrogen bonding.
Show answer
(1) A hydrogen covalently bonded to N, O, or F (donor); (2) a lone pair on an N, O, or F atom of another molecule (acceptor). Without both, there is no hydrogen bond.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Noncovalent interaction
- Attraction between molecules, weaker than covalent bonds
- Ion–dipole interaction
- Attraction between a charge and a polar molecule's dipole
- Dipole–dipole interaction
- Attraction between the opposite ends of two molecular dipoles
- Hydrogen bond
- Strong dipole–dipole attraction between X-H (X = N, O, F) and a lone pair
- London dispersion force
- Attraction from temporary induced dipoles
- Polarizability
- How easily an electron cloud distorts
- Amphiphile
- Molecule with polar head + nonpolar tail
- Like dissolves like
- Rule that solubility matches polarity
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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