Identify intermolecular forces and use their strengths to predict melting and boiling points, solubility and vapor pressure

Learning goal: Distinguish intermolecular forces of attraction from the bonds inside molecules, identify London dispersion forces, dipole-dipole forces and hydrogen bonding in a given substance and rank their strengths, and use the type and strength of the forces to predict and explain melting and boiling points, solubility and vapor pressure, including why water behaves unusually.

Between molecules, not within them. In middle school you explained solids, liquids and gases by the spacing, motion and attraction of particles. This lesson names the attractions. A chemical bond holds atoms together: covalent bonds hold the atoms of a molecule together (intramolecular bonds), while ionic and metallic bonding hold a whole lattice together with no separate molecules at all. An intermolecular force is a much weaker attraction between separate molecules, and it is what must be overcome to melt or boil a molecular substance. When water boils, the O–H bonds inside each molecule do not break; the molecules simply move apart. Intermolecular forces are generally weak compared with bonds, but they span a wide range, and that range is what keeps helium a gas down to about −269 °C (it boils at −268.9 °C) while water is a liquid at room temperature.

London dispersion forces: present in everything. Electrons in a molecule move constantly, so at any instant the charge may be slightly uneven, with a momentary partial negative side and partial positive side. That flicker induces a matching flicker in a neighbor, and the two attract for a moment. These London dispersion forces exist between all molecules and are the only intermolecular force between nonpolar ones: the noble gases, the halogens, hydrocarbons and oils. Their strength grows with the number of electrons and the size of the molecule, because larger electron clouds distort more easily, and with contact area, because more of the surface can interact. That is why the boiling points of the straight-chain alkanes climb steadily: methane −162 °C, ethane −89 °C, propane −42 °C, butane about −1 °C, pentane 36 °C. Each added carbon adds electrons and length. Branching works the other way: pentane, 2-methylbutane and 2,2-dimethylpropane all have the formula C₅H₁₂, but their boiling points are 36 °C, 28 °C and 9.5 °C, because a compact branched molecule has less surface in contact with its neighbors and the dispersion forces are weaker. Liquefied inert gases are held together by nothing but dispersion forces.

Dipole-dipole forces: polar molecules. A polar molecule has a permanent partial positive end and partial negative end. Neighboring polar molecules line up so that opposite ends attract, adding a dipole-dipole force on top of the dispersion forces every molecule already has. Whether a molecule is polar was settled in the previous lesson: polar bonds arranged so that their dipoles do not cancel. Comparing molecules of similar size, the polar one has the higher boiling point.

Hydrogen bonding: a special, strong dipole-dipole case. When hydrogen is bonded to fluorine, oxygen or nitrogen, the bond is so polar and hydrogen so small that the exposed hydrogen can approach very close to a lone pair on the F, O or N of a neighboring molecule. The resulting attraction, called a hydrogen bond, is the strongest common intermolecular force, though still far weaker than a covalent bond. In water it is an attraction between molecules, not a bond within one, though hydrogen bonds can also form between distant parts of one large molecule, as in folded proteins. The relative strengths are: hydrogen bonding > dipole-dipole > London dispersion, for molecules of comparable size; a very large nonpolar molecule can have dispersion forces that outweigh the dipole-dipole forces of a small polar one, so size always has to be considered.

Melting and boiling points. The temperature of a substance measures the average kinetic energy of its particles. To melt, and especially to boil, the particles need enough kinetic energy to overcome the attractions between them, so, in general, the stronger the intermolecular forces the higher the boiling point, and usually the melting point too. Melting also depends on how neatly the molecules pack into a crystal, so melting points follow the forces less reliably: compact, symmetrical 2,2-dimethylpropane boils lower than pentane (9.5 °C against 36 °C) yet melts far higher (about −16 °C against about −130 °C), because its ball-like molecules pack well into a solid. Substances with strong forces, or with three-dimensional networks of ionic or covalent bonds in which there are no separate molecules at all, are solids at room temperature with high melting points. Compare water, hydrogen sulfide and methane, all small molecules: methane is nonpolar and boils at −162 °C; hydrogen sulfide is polar but cannot hydrogen bond and boils at −60 °C; water hydrogen bonds and boils at 100 °C. Ammonia, which hydrogen bonds but less extensively, boils at −33 °C. Ethanol, C₂H₅OH, and dimethyl ether, CH₃OCH₃, have the same formula; ethanol has an O–H and boils at 78 °C while the ether, with no O–H, boils at about −24 °C.

Water is unusual, and it matters. Each water molecule can form hydrogen bonds through both of its hydrogens and both of its lone pairs, so water hydrogen bonds more extensively than any other common liquid. That gives it a boiling point far above what its size predicts, a high surface tension, so that drops bead up on wax and many drops fit on a coin, and a high specific heat. In ice the hydrogen bonds hold the molecules in an open hexagonal network with more empty space than in the liquid, so ice is less dense than water and floats. A frozen pond therefore freezes from the top down, and the ice layer insulates the water below, which stays liquid through winter. Hydrogen bonding between the two strands of DNA and the way lipid molecules arrange themselves into membranes, with their polar heads toward water and nonpolar tails away from it, are the same forces at work in living things.

Solubility: like dissolves like. A solute dissolves when the attractions it can form with the solvent are comparable to the attractions it gives up. Substances with similar intermolecular forces tend to mix; those with very different forces tend to stay apart, because each prefers to interact with its own kind. Water is polar and hydrogen bonds, so it dissolves many ionic and polar covalent substances (though not all: silver chloride and calcium carbonate barely dissolve): when an ionic solid dissolves, the attractions between ions in the lattice are replaced by attractions between each ion and the oppositely charged ends of many water molecules, and the ions end up surrounded by a shell of oriented water. Comparing these attractions is a useful qualitative guide, not a complete test: whether a salt dissolves, and how much, depends on the balance of the whole process, so measured solubility data decide particular cases. Sugar dissolves because its many O–H groups hydrogen bond with water. Oil is nonpolar; its molecules have only dispersion forces, water molecules would have to give up hydrogen bonds to surround them, and so oil and water separate. Nonpolar solutes dissolve instead in nonpolar solvents. Dissolving is a process that is genuinely hard to classify as a physical or a chemical change, and it is not classified either way here. Also keep solubility, the amount that can dissolve, separate from the rate of dissolving, which is a kinetics question.

Vapor pressure. In any liquid, a few molecules at the surface have enough kinetic energy to escape the attractions of their neighbors and become gas; that is evaporation. In a closed container the escaped molecules exert a pressure, the vapor pressure, which rises with temperature because more molecules have enough energy to escape. The stronger the intermolecular forces, the fewer molecules escape at a given temperature and the lower the vapor pressure. Water at 20 °C has a vapor pressure of about 2.3 kPa; ethanol, with weaker hydrogen bonding, has a higher vapor pressure at the same temperature and evaporates faster. A liquid boils when its vapor pressure equals the pressure of the atmosphere above it, which is why the normal boiling point of water is the temperature at which its vapor pressure reaches 101.3 kPa and why water boils at a lower temperature on a mountain. A graph of vapor pressure against temperature for several liquids ranks their intermolecular forces: the curve that reaches atmospheric pressure first has the weakest forces. Molecular substances with weak forces evaporate readily and are often detected by their smell; ionic and network solids have such strong attractions that they hardly vaporize and are usually odorless.

A worked model: rank three liquids. Three liquids are in identical open dishes: water, ethanol and vegetable oil. Water has extensive hydrogen bonding; ethanol has one O–H per molecule and hydrogen bonds less extensively; oil molecules are large and nonpolar, with dispersion forces only, but they are very large molecules. Predicted vapor pressure at room temperature: ethanol highest, then water, then oil, because oil's huge molecules have very strong dispersion forces in total. Predicted order of evaporation in the dishes: ethanol fastest, then water, oil barely at all. Predicted mixing: water and ethanol mix, since both hydrogen bond; oil mixes with neither. Every prediction names a force and compares its strength.

Where the rest of the standards is taught. Formula writing, naming and Lewis structures from C.PM.4 are taught in the formulas and shapes lesson; this lesson uses shape and polarity from there to predict intermolecular forces.

Investigation: forces you can see. Materials: water, 70 percent rubbing alcohol (isopropyl alcohol), vegetable oil, a dropper, a coin, wax paper, a glass plate, three shallow dishes, a timer, a small clear jar with a lid, goggles. Keep the alcohol away from any flame and work in a ventilated room. Steps: (1) count how many drops of water fit on the coin before it overflows, then repeat with alcohol; (2) place one drop of water on wax paper and one on glass and sketch the shapes; (3) put a single drop of each liquid on the plate and time how long each takes to disappear, repeating three times; (4) shake oil and water in the jar and watch for two minutes. Record counts, sketches, times and observations. A result that supports the claim: water forms a tall dome of many drops on the coin and beads on wax but spreads on glass; alcohol evaporates first and water second while oil stays; oil and water separate into layers. What the model cannot show: you are seeing the consequences of intermolecular forces, not the forces themselves, and drop counts and timing depend on room temperature and air movement, so trials are compared only against one another on the same day.

Worked example 1 — identify the forces present. For each substance, list the intermolecular forces. Bromine, Br₂: nonpolar, so dispersion forces only. Hydrogen chloride, HCl: polar (difference 0.96) but hydrogen is bonded to chlorine, not F, O or N, so dispersion plus dipole-dipole, no hydrogen bonding. Methanol, CH₃OH: has an O–H, so dispersion, dipole-dipole and hydrogen bonding. Carbon dioxide: polar bonds but a linear, symmetric molecule, so nonpolar and dispersion only. Every substance has dispersion forces; the question is what is added.

Worked example 2 — predict boiling point order. Rank CH₄, H₂O and H₂S from lowest to highest boiling point. All are small. CH₄ is nonpolar: dispersion only, lowest. H₂S is bent and polar but sulfur is not electronegative enough for hydrogen bonding: dispersion plus dipole-dipole, middle. H₂O hydrogen bonds extensively: highest. Predicted order CH₄ < H₂S < H₂O, and the data agree: −162 °C, −60 °C, 100 °C. The large gap between H₂S and H₂O, 160 degrees, is the signature of hydrogen bonding.

Worked example 3 — branching and chain length. Predict which has the higher boiling point: hexane (six carbons in a chain) or pentane (five), and pentane or 2,2-dimethylpropane. Hexane has more electrons and more length in contact, so stronger dispersion forces and a higher boiling point than pentane. Pentane's straight chain touches its neighbors along its whole length; 2,2-dimethylpropane is nearly spherical with less contact, so its dispersion forces are weaker and its boiling point is lower: 36 °C against 9.5 °C, although both are C₅H₁₂. Same formula, different shape, different forces.

Worked example 4 — choose a solvent. A student wants to dissolve a waxy nonpolar residue from a lab bench and, separately, some spilled potassium chloride. Wax has only dispersion forces, so it dissolves in a nonpolar solvent such as a hydrocarbon or oil, not in water, which would have to give up hydrogen bonds to surround it. Potassium chloride is ionic; water's polar molecules surround K⁺ with their oxygen ends and Cl⁻ with their hydrogen ends, and the many ion-water attractions make up for the lattice attractions given up, so water is the solvent (solubility data confirm that potassium chloride is very soluble). Like dissolves like is a summary of comparing forces, not a separate rule.

Representation task: three pictures of attraction. Draw (1) two water molecules with a dotted line from a hydrogen of one to a lone pair on the oxygen of the other, labeled hydrogen bond, and the O–H bonds labeled covalent; (2) two HCl molecules aligned positive end to negative end, labeled dipole-dipole; (3) two argon atoms with a momentary uneven electron cloud in each, labeled London dispersion. Under the pictures, rank the three forces by strength for molecules of similar size and draw a sketch graph of vapor pressure against temperature for water and ethanol, marking 101.3 kPa and each liquid's normal boiling point (100 °C and 78 °C).

Independent performance: forces investigation and prediction set. Carry out the four-part investigation and write a report. Success criteria: a data table with three trials for each measurement; an explanation for each observation that names the intermolecular force responsible; a paragraph on why water does not follow the trend expected from its size, using the words hydrogen bonding, surface tension, density of ice and boiling point; a prediction, with reasons, of the order of boiling points and of vapor pressures at room temperature for four substances your teacher names; and a statement of what the investigation shows directly and what it only lets you infer.