Difficulty: Intermediate | Prerequisites: Functional groups (Ch. 8.1-8.2), Lewis structures, electronegativity
This section explains why different organic compounds have different boiling points, melting points and physical behaviour. The answer lies in noncovalent (intermolecular) forces: van der Waals attractions, dipole-dipole interactions and hydrogen bonding. Understanding these forces is essential for predicting solubility, designing separations and making sense of spectroscopy data later in the course. You should already know what functional groups are and how electronegativity works.
Molecules in a liquid or solid are held together by noncovalent attractions, not by breaking or forming covalent bonds. The three main types, in increasing strength, are van der Waals (London dispersion) forces, dipole-dipole attractions and hydrogen bonds. A molecule's size, shape and polarity determine which forces dominate, and that in turn sets its boiling point, melting point and solubility.
Noncovalent intermolecular interactions
Attractions between molecules that do not involve breaking or forming covalent bonds. Think of them as the "glue" that holds molecules near each other in a liquid or solid.
Van der Waals attractions (London dispersion forces)
The weakest intermolecular force. Arise from temporary, induced dipoles caused by momentary distortions in a molecule's electron cloud. Present in all molecules, and are the only intermolecular force available to nonpolar species such as alkanes.
Induced dipole
A temporary dipole created in a molecule when a nearby charge or dipole distorts its electron cloud. In simple terms, one molecule's electrons shift, and the neighbouring molecule's electrons shift in response.
Flickering dipoles
Another name for the rapidly forming and dissipating temporary dipoles involved in van der Waals attractions. These form and reform on a femtosecond timescale.
Dispersion interaction
The cohesive interaction that must be overcome to vaporise a liquid. Essentially another way to describe the net van der Waals attraction holding a liquid together.
Polarisability
A measure of how easily an external charge or dipole can distort the electron distribution of a molecule or atom. In simple terms, a molecule with a "soft" or easily deformed electron cloud is highly polarisable.
Permanent dipole
A dipole that exists all the time in a molecule because of differences in electronegativity between bonded atoms (e.g. the C-O bond in ethers). Not temporary like a van der Waals dipole.
Dipole-dipole attraction
The attraction between the positive end of one permanent dipole and the negative end of another. Raises the boiling point above what van der Waals forces alone would predict.
Hydrogen bond
The strongest common intermolecular force. An attraction that occurs when a hydrogen bonded to F, O or N interacts with a lone pair on another F, O or N atom. Think of it as a particularly strong dipole-dipole interaction, limited to the most electronegative elements.
Hydrogen-bond donor
The atom (F, O or N) to which the hydrogen is covalently bonded.
Hydrogen-bond acceptor
The atom bearing the lone pair that the hydrogen interacts with.
Boiling point
The temperature at which a liquid's vapour pressure equals atmospheric pressure. Serves as a proxy for the strength of intermolecular attractions in the liquid state.
Melting point
The temperature at which a solid transforms into a liquid (solid and liquid in equilibrium). Reflects noncovalent interactions in both the liquid and the crystalline solid state.
Symmetry number
The number of equivalent, indistinguishable orientations of a molecule related by rotations. Higher symmetry means better crystal packing and, typically, a higher melting point.
Enthalpy of fusion
The enthalpy difference between the solid and its liquid melt at the melting temperature. A direct measure of the strength of noncovalent interactions among molecules in the solid.
Entropy of fusion
The entropy difference between the solid and its liquid melt at the melting temperature. Reflects how much disorder is gained on melting.
Molecules can interact without breaking covalent bonds. These noncovalent interactions can raise or lower the energy of the molecules involved, and molecules will form attractive noncovalent interactions whenever possible. The molecular structure of a compound determines the magnitude of its noncovalent interactions.
Any condensed phase (liquid or solid) owes its existence to these attractions. Without them, every substance would be a gas.
Arise from temporary, induced dipoles. Even in a perfectly nonpolar molecule like an alkane, the electron cloud can momentarily become lopsided, creating a fleeting dipole.
When a second molecule is nearby, its electron cloud distorts to form a complementary dipole. The positive end of one temporary dipole attracts the negative end of the other.
These temporary dipoles form and dissipate on a femtosecond timescale ("flickering dipoles"). They can form and re-form many times during a single molecular collision (which itself lasts roughly a nanosecond).
Van der Waals attractions increase with the surface area of the interacting electron clouds. Larger molecules have more surface area, so they have stronger dispersion forces and higher boiling points.
Surface area vs. volume:
It is surface area, not volume, that controls boiling point.
Branching makes a molecule more compact (more spherical), reducing its effective surface area.
Heavily branched isomers therefore have lower boiling points than their unbranched counterparts of the same molecular formula.
Two key rules for boiling-point trends in a homologous series:
Boiling points increase with molecular weight, roughly 20-30 degrees C per additional carbon atom, because of increased van der Waals attractions.
Boiling points decrease with increased branching, because of reduced molecular surface area.
Even/odd carbon effect: alkanes with an even number of carbons have slightly higher boiling points than predicted, while odd-numbered alkanes come in slightly lower.
Polarisability measures how easily an external charge or dipole can distort a molecule's electron cloud.
More polarisable molecules have "softer" electron clouds that distort more readily, leading to stronger induced dipoles and greater van der Waals forces.
Factors that increase polarisability: many electrons, electrons far from the nucleus, delocalised electrons, atoms of lower electronegativity.
Factors that decrease polarisability: high effective nuclear charge, tightly held (localised) electrons, highly electronegative atoms.
A liquid made of more polarisable molecules will have a higher boiling point.
Molecules with permanent dipole moments (from polar bonds) show enhanced intermolecular attractions beyond van der Waals alone.
The negative end of one dipole aligns with the positive end of another, and this attraction can occur between many molecules simultaneously.
This raises the boiling point of a polar compound above that of a nonpolar molecule of similar size and shape.
The electronegativity of oxygen in ethers, for instance, creates a permanent dipole while reducing polarisability. The effects of surface area and polarity are inversely proportional to each other.
A hydrogen bond forms when a hydrogen covalently bonded to F, O or N interacts with a lone pair on another F, O or N atom.
These three elements are singled out because they are the most electronegative, so the X-H bond is extremely polar.
The hydrogen remains covalently bonded to the donor atom, but interacts weakly (electrostatically) with the acceptor atom. The covalent bond to H is shorter than the hydrogen bond itself.
Some atoms (e.g. oxygen in water or an alcohol) can act as both donor and acceptor.
Best hydrogen-bond donor atoms: oxygen, nitrogen, fluorine, and Bronsted acids.
Hierarchy of intermolecular forces (strongest to weakest effect on boiling point):
Hydrogen bonding: molecules that can hydrogen-bond have the highest boiling points among compounds of similar polarity.
Permanent dipole-dipole: molecules with permanent dipole moments boil higher than nonpolar molecules of similar size and shape.
Van der Waals (dispersion): influenced, in order of importance, by molecular surface area, molecular shape (extended vs. branched) and polarisability.
Strengths of noncovalent attractions are entirely separate from the stability of covalent bonds. Strong covalent bonds and weak intermolecular forces can coexist in the same molecule.
Melting points reflect noncovalent interactions in both the liquid state and the crystalline solid state.
Molecular symmetry plays a special role: more symmetric molecules pack more efficiently into a crystal lattice, raising the melting point.
The symmetry number is the count of equivalent, indistinguishable orientations a molecule can adopt through rotations. Higher symmetry number tends to mean higher melting point.
Enthalpy of fusion is a direct measure of the noncovalent interactions in the solid. Entropy of fusion reflects how much disorder is gained on melting.
Chemical structure --> determines how structures interact noncovalently --> determines physical properties (boiling point, melting point, solubility).
Boiling point as a proxy: BP = temperature at which vapour pressure = atmospheric pressure. Higher BP = stronger intermolecular attractions in the liquid.
Van der Waals strength scales with: molecular surface area > molecular shape (extended > branched) > polarisability.
Hydrogen bond geometry: X-H - - - Y, where X and Y are each F, O or N. The covalent X-H bond is shorter than the H - - - Y hydrogen bond.
The reason water has an unusually high boiling point for its molecular weight is extensive hydrogen bonding. This is why water is a liquid at room temperature while similar-sized molecules (e.g. methane, 16 g/mol) are gases.
Distillation, one of the most common laboratory and industrial separation techniques, works by exploiting differences in boiling points, which are ultimately differences in intermolecular forces. If you can predict relative boiling points from molecular structure, you can predict which component comes off the column first.
Students often confuse intermolecular forces with covalent bonds. Boiling a liquid does not break covalent bonds; it overcomes the noncovalent attractions between molecules. The molecules themselves remain intact.
Students sometimes think hydrogen bonding can occur with any hydrogen. It requires the hydrogen to be bonded to F, O or N, and the acceptor must also be F, O or N with a lone pair.
Students mix up molecular volume and molecular surface area when predicting boiling points. A compact, branched molecule may have the same volume as its linear isomer, but less surface area, and therefore weaker dispersion forces and a lower boiling point.
Students sometimes assume that higher molecular weight always means higher boiling point. Branching and polarity can override the molecular-weight trend within a class of compounds.
Ranking compounds by boiling point is a very common exam question. You will be given several structures and asked to order them from lowest to highest BP. Identify which intermolecular forces each compound can exhibit, then rank accordingly.
Know the hierarchy: hydrogen bonding > dipole-dipole > van der Waals. Within van der Waals, surface area and polarisability are the tiebreakers.
Be prepared to explain why branching lowers the boiling point.
Melting-point questions often test symmetry: a more symmetric molecule packs better and melts higher. The neopentane vs. pentane comparison is a classic.
The distinction between enthalpy and entropy of fusion may appear on exams. Enthalpy of fusion measures intermolecular interaction strength; entropy of fusion measures the disorder gained on melting.
True or false: Van der Waals forces are present only in nonpolar molecules. (False. They are present in all molecules, but they are the only intermolecular force available to nonpolar species.)
Fill in the blank: Hydrogen bonds require H to be bonded to ______, ______ or ______. (F, O or N.)
True or false: A heavily branched alkane boils higher than its straight-chain isomer. (False. Branching reduces surface area and lowers the boiling point.)
Fill in the blank: The measure of how easily an electron cloud can be distorted is called ______. (Polarisability.)
True or false: Melting a solid breaks its covalent bonds. (False. Melting overcomes noncovalent intermolecular attractions; covalent bonds remain intact.)
Q: Rank the following in order of increasing boiling point: pentane, 1-butanol, diethyl ether. Explain your reasoning.
A: Pentane (lowest) < diethyl ether < 1-butanol (highest). Pentane has only van der Waals forces. Diethyl ether has a permanent dipole (C-O bonds) adding dipole-dipole attractions on top of dispersion. 1-Butanol can form hydrogen bonds through its -OH group, giving it the strongest intermolecular attractions and the highest boiling point.
Q: Two alkanes have the same molecular formula C5H12. One is n-pentane, the other is neopentane (2,2-dimethylpropane). Which has the higher boiling point, and why?
A: n-Pentane. It has an extended, linear shape with more molecular surface area for van der Waals contact. Neopentane is nearly spherical, reducing its surface area and its dispersion interactions.
Q: Why does acetone (a ketone with a permanent dipole) have a lower boiling point than propan-1-ol, even though both have three carbons?
A: Propan-1-ol can act as both a hydrogen-bond donor (through its O-H) and acceptor (through the oxygen lone pairs). Acetone can only accept hydrogen bonds (it has no O-H); it relies on dipole-dipole and dispersion forces. The additional hydrogen-bonding capability of propan-1-ol gives it stronger intermolecular attractions and a higher boiling point.
Q: What property of a molecule determines how strong its London dispersion forces are?
A: Polarisability, which depends on the size and shape of the electron cloud, the number of electrons, how far the outer electrons are from the nucleus and how tightly they are held. Greater polarisability means stronger dispersion forces.
Q: Explain why molecular symmetry raises the melting point of a compound.
A: More symmetric molecules can pack more efficiently into a crystal lattice. Better packing means more noncovalent contacts per molecule in the solid, increasing the enthalpy of fusion and raising the temperature needed to disrupt the lattice.
These intermolecular forces directly control the solubility and solution behaviour covered in Ch. 8.6. The "like dissolves like" rule is a consequence of matching intermolecular interactions between solute and solvent.
Hydrogen bonding is central to biochemistry: it holds together DNA base pairs, stabilises protein secondary structure (alpha helices and beta sheets), and determines the unique properties of water.
Polarisability and dispersion forces come back in spectroscopy and chromatography, where differences in intermolecular interactions determine how molecules separate on a column or interact with electromagnetic radiation.
noncovalent interactions, intermolecular forces, van der Waals forces, London dispersion forces, induced dipole, temporary dipole, flickering dipoles, permanent dipole, dipole-dipole attraction, hydrogen bonding, hydrogen bond donor, hydrogen bond acceptor, polarisability, polarizability, boiling point trends, melting point trends, surface area vs volume, branching effect, symmetry number, enthalpy of fusion, entropy of fusion, crystal packing, organic chemistry chapter 8.4 8.5, CHEM 301, homologous series, noncovalent attractions