How IMFs explain physical properties — Chemistry

Before this

This leads to

A thousand times thicker

Pour four clear, colourless liquids side by side. Ethanol and water run off the spoon almost at once. Ethylene glycol, used in car antifreeze, pours more slowly. Glycerol, used in medicines and soaps, oozes like syrup. At 25 °C, glycerol is about a thousand times more resistant to flow than water.

Nothing about the colour or the look of these liquids explains this. The reason is invisible: the forces between their molecules. The last three lectures built three of these forces, London forces, dipole–dipole forces and hydrogen bonds, mostly to explain boiling points. The hydrogen bonding lecture ended with a promise: to bring all three together and use them to explain the physical properties of many kinds of substances. This lecture keeps it. Boiling, flowing, surface "skins" and dissolving all come from the same small set of attractions.

One chain of reasoning

Every explanation in this lecture follows the same chain. Learn the chain, not the list of facts.

From structure to property

Structure of the molecule → which intermolecular forces it has → how strong they are → the physical property.

Stronger attractions between molecules mean a higher boiling point, a higher viscosity and a higher surface tension.

To decide how strong the forces are, ask three questions, in this order. First, which forces does each substance have? Every molecule has London forces. Polar molecules add dipole–dipole forces. Molecules with H on N, O or F, and a lone pair on N, O or F, add hydrogen bonds. Second, are the molecules about the same size? Count the electrons. If one molecule has far more electrons, its London forces may outweigh everything else, as with iodine or CCl₄ in the earlier lectures. Third, if the sizes are similar, the extra forces decide: hydrogen bonds count for most, then dipole–dipole forces. For hydrogen-bonded substances, also count how many hydrogen bonds each molecule can make.

Stop and think

A student writes: "Water has a high boiling point, so it must have strong intermolecular forces." What is wrong with this as an explanation?

Show the answer

The conclusion is true, but the reasoning runs backwards. It uses the boiling point, the thing to be explained, as the evidence. A real explanation starts from the structure: water has O–H bonds and lone pairs on oxygen, so its molecules form hydrogen bonds, about two per molecule. Hydrogen bonds are strong, so a lot of energy is needed to separate the molecules. So water's boiling point is high. Structure first, property last.

Boiling points

To boil, a molecule must escape from the liquid completely, breaking free of all its neighbours. So the boiling point is the most direct measure of how strongly the molecules hold each other. Boiling never breaks the covalent bonds inside the molecules: steam is still H₂O.

Worked Example 1 — Pentane and butanone

Pentane, C₅H₁₂, boils at about 36 °C. Butanone, CH₃COCH₂CH₃, boils at about 80 °C. Explain the difference.

First count the electrons. Pentane: $5 \times 6 + 12 \times 1 = 42$ . Butanone, C₄H₈O: $4 \times 6 + 8 \times 1 + 8 = 40$ . Almost the same, so their London forces are similar.

Pentane is a hydrocarbon, so it is non-polar and has London forces only. Butanone has a C=O group, which is polar (difference in electronegativity $3.44 - 2.55 = 0.89$ ), so it has dipole–dipole forces as well. It cannot hydrogen-bond with itself, because all its hydrogens are on carbon, like propanone. With London forces matched, the extra dipole–dipole forces raise butanone's boiling point by about 44 degrees.

The same reasoning, with a stronger extra force, gives a far bigger gap. Ethanol and dimethyl ether have the same formula, C₂H₆O, but ethanol boils at about 78 °C and dimethyl ether at about −24 °C, a difference of about 100 degrees. Check the explanation below. One line of it is wrong.

Spot the error — ethanol and dimethyl ether

A student explains why ethanol, CH₃CH₂OH, boils so much higher than dimethyl ether, CH₃OCH₃. One line is wrong. Tap it. Every line has an explanation.

The explanation, with the mistake marked and corrected:

Line 1. Both are C₂H₆O, so both have 26 electrons and similar London forces. (Correct.)
Line 2. Ethanol has an O–H group, so one of its hydrogens is bonded directly to oxygen. (Correct.)
Line 3. Dimethyl ether cannot form hydrogen bonds because it has no oxygen atom. The mistake: it does have an oxygen, with lone pairs. It cannot hydrogen-bond with itself because all six hydrogens are bonded to carbon, so it has no donor.
Line 4. So ethanol molecules hydrogen-bond to each other, and more energy is needed to separate them. (Correct.)
Line 5. That is why ethanol boils at about 78 °C and dimethyl ether at about −24 °C. (Correct.)

Viscosity: how easily a liquid flows

Viscosity is a liquid's resistance to flowing. When a liquid flows, its molecules must slide past their neighbours, breaking old attractions and making new ones all the time. The stronger the attractions, the harder the sliding, and the more viscous the liquid.

Liquid (25 °C)	O–H groups per molecule	Viscosity in mPa·s
Octane, C₈H₁₈	0	0.508
Water, H₂O	2 (as O–H bonds)	0.890
Ethanol, C₂H₅OH	1	1.074
Ethylene glycol, C₂H₄(OH)₂	2	16.1
Glycerol, C₃H₅(OH)₃	3	934

Look down the alcohols. Ethanol has one O–H group, ethylene glycol two, glycerol three. Each extra O–H group lets a molecule make more hydrogen bonds, to more neighbours at once. So each molecule is tied into the liquid at more points, and sliding gets much harder. The viscosity rises by roughly 15 times, then about 60 times. Water has two O–H bonds but is a tiny molecule, so it slides more easily than ethylene glycol.

Size matters too, even without hydrogen bonds. Motor oil is made of long hydrocarbon chains, with London forces only, yet it is far more viscous than water. Long chains have many electrons and touch their neighbours along their whole length, and they can tangle. Now predict a closer contest.

Predict, then reveal — ethanol or octane?

Octane, C₈H₁₈, has 66 electrons. Ethanol, C₂H₅OH, has 26. Which is more viscous at 25 °C?

Ethanol, about twice as viscous: 1.074 mPa·s against 0.508 mPa·s. Octane has much stronger London forces, but ethanol's hydrogen bonds tie each molecule to its neighbours more firmly when they slide past each other.

This is a warning against one-rule thinking. Two forces are always in play, and different properties can weigh them differently. Here, ethanol flows more slowly, but octane boils at a higher temperature. Always name the forces in each substance before you compare.

Temperature changes viscosity a great deal. Glycerol's viscosity falls from 934 mPa·s at 25 °C to 152 mPa·s at 50 °C and to about 15 mPa·s at 100 °C. Hotter molecules move faster, so they break free from their neighbours more easily. This is why honey and syrup pour more easily when warm.

Surface tension: a skin on the liquid

A carefully placed steel needle can rest on the surface of water, and some insects walk on ponds. The surface behaves as if it had a thin, stretched skin. This effect is called surface tension.

Inside a liquid, a molecule is attracted equally in every direction, so the pulls cancel. At the surface, there are no liquid molecules above, so a molecule is pulled only sideways and downwards. The result is a net pull into the liquid, which makes the surface contract like a stretched skin.

Because surface molecules are pulled inwards, the liquid behaves as if it is trying to make its surface as small as possible. That is why small drops are round: a sphere has the smallest surface for its volume. Stronger attractions mean a stronger inward pull, so a higher surface tension. At 25 °C, water's surface tension is about 72 mN/m, ethylene glycol's about 48, ethanol's about 22 and octane's about 21. Water, with two hydrogen bonds per molecule in a three-dimensional network, is far ahead.

Notice ethanol and octane again: almost equal this time. Ethanol's hydrogen bonds and octane's larger London forces come out about level. The same two substances gave a clear winner for viscosity and a near draw for surface tension. That is not a contradiction. It shows that each property depends on the balance of forces in a slightly different way.

Dissolving: like dissolves like

When a solid or liquid dissolves, three things happen. Solute particles separate from each other. Solvent molecules move apart to make room. Then solute and solvent attract each other. The first two steps need energy. The third releases energy. Dissolving happens easily when the new solute–solvent attractions are about as strong as the ones that had to be broken.

That gives a useful general rule, "like dissolves like": polar and ionic substances dissolve in polar solvents, and non-polar substances dissolve in non-polar solvents.

Salt in water. The δ− oxygens of water molecules surround each Na⁺ ion, and the δ+ hydrogens surround each Cl⁻ ion. These ion–dipole attractions are strong enough to pull the ions out of the crystal.
Ethanol in water. Ethanol hydrogen-bonds to water in both directions, as the hydrogen bonding lecture showed, so they mix in any proportion.
Iodine in hexane. Both are non-polar, with London forces only. Iodine–hexane attractions are similar to the ones broken, so iodine dissolves, giving a purple solution.

Worked Example 2 — Why oil and water separate

Cooking oil is shaken with water. A few minutes later, the two have separated into layers, with the oil on top. Explain, starting from the structures.

Structures and forces. Oil molecules are long hydrocarbon chains. They are non-polar, so the only forces between them are London forces. Water molecules have O–H bonds and lone pairs, so they form a network of hydrogen bonds, about two per molecule.

What dissolving would need. For oil to dissolve, water molecules would have to move apart to make room for each oil molecule, breaking hydrogen bonds. In return, they could form only weak London attractions with the oil, because oil has no δ+ hydrogens and no lone pairs to offer.

The result. Mixing costs much more than it gives back. So water molecules stay attracted to each other, the oil molecules are pushed together, and the liquids separate. The oil floats because it is less dense than water. Notice that the oil is not repelled by water. Water simply attracts itself much more strongly than it attracts oil.

Stop and think

Methanol, CH₃OH, mixes with water in any proportion. Long-chain alcohols, with eight or more carbon atoms, barely dissolve in water at all. Both have an O–H group. Why the difference?

Show the answer

The O–H group can hydrogen-bond to water, but the hydrocarbon chain cannot. In methanol, the O–H group is most of the molecule, so it fits easily into water's network. In a long-chain alcohol, a long non-polar tail must push into the water, breaking many water–water hydrogen bonds for only weak London forces in return. The longer the tail, the more the molecule behaves like oil. Solubility depends on the balance of polar and non-polar parts, not just on whether an O–H group is present.

Three classic mistakes

Trap 1 — Explaining backwards

"It has a high boiling point, so it has strong forces" uses the result as the reason. Start from the structure, name the forces, then reach the property.

Trap 2 — Ranking the forces without checking size

Hydrogen bonds beat London forces only between molecules of similar size. Octane (66 electrons) boils above ethanol (26 electrons), even though ethanol hydrogen-bonds. Count electrons first.

Trap 3 — Breaking covalent bonds

Boiling, flowing and dissolving a molecular substance overcome forces between molecules. The covalent bonds inside each molecule stay whole.

Try it yourself

Work each one on paper first. Open a hint only when you are stuck.

Propane, CH₃CH₂CH₃, dimethyl ether, CH₃OCH₃, and ethanol, CH₃CH₂OH, each have 26 electrons. Put them in order of boiling point, lowest first, and explain.

Hint 1
Same electron count, so similar London forces. What else does each have?

Hint 2
Which is non-polar, which is polar with no donor, and which has an O–H group?

Full answer
Propane, then dimethyl ether, then ethanol. All three have similar London forces. Propane is non-polar, so that is all it has. Dimethyl ether is polar (its C–O–C is bent), so it adds dipole–dipole forces, but it has no O–H donor. Ethanol adds hydrogen bonds. Each step adds a stronger attraction, so each step raises the boiling point. (Dimethyl ether boils at about −24 °C and ethanol at about 78 °C; propane boils lower than both.)
Explain why glycerol, with three O–H groups, is far more viscous than ethanol, with one, even though both are alcohols.

Hint 1
What does each O–H group let a molecule do?

Hint 2
What must happen to neighbouring molecules when a liquid flows?

Full answer
Each O–H group can donate and accept hydrogen bonds, so a glycerol molecule can be hydrogen-bonded to many neighbours at once, in several directions. To flow, molecules must slide past each other, repeatedly breaking and remaking these attractions. With three O–H groups, each glycerol molecule is held at many more points than an ethanol molecule, so flow is much harder. At 25 °C, glycerol is about 934 mPa·s and ethanol about 1.07 mPa·s, roughly 870 times different.
Which of these dissolve well in water, and which in hexane? Sodium chloride, iodine, glucose (a sugar with five O–H groups), and octane. Explain each briefly.

Hint 1
Decide whether each is ionic, polar with many O–H groups, or non-polar.

Hint 2
Like dissolves like.

Full answer
Water: sodium chloride (ions attracted strongly by water's δ− and δ+ ends) and glucose (its five O–H groups hydrogen-bond to water). Hexane: iodine and octane (non-polar, London forces only, just like hexane). Neither salt nor glucose dissolves well in hexane, because hexane can offer only weak London forces in place of the strong attractions that hold them together.
Glycerol's viscosity is 934 mPa·s at 25 °C and about 14.8 mPa·s at 100 °C. By what factor does it fall? Explain the fall in terms of molecules, and say why its boiling point does not fall when it is heated.

Hint 1
Divide the two viscosities.

Hint 2
The forces between molecules are the same at both temperatures. What changes is the molecules' energy.

Full answer
$934 \div 14.8 \approx 63$ , so it falls by a factor of about 63. Heating gives the molecules more kinetic energy, so they break free from their hydrogen-bonded neighbours more often and slide past each other more easily. The attractions themselves are unchanged; the molecules simply overcome them more often. The boiling point is a fixed property of the substance at a given pressure: it is the temperature at which the molecules have enough energy to escape completely. Heating the liquid moves it towards its boiling point; it does not change where the boiling point is.

The bigger idea

Three small attractions explain a remarkable range of everyday things: why water is a liquid, why syrup is slow, why insects can stand on a pond, why oil and vinegar separate. Medicine depends on the same balance. A drug must dissolve in blood, which is mostly water, but it must also pass through cell membranes, which are oily. Chemists tune each new drug's balance of polar and non-polar parts. One of the first things they measure is how the drug divides itself between an oily layer and a watery one.

This completes the forces between molecules. The next topic, metallic bonding, is a different kind of attraction altogether: not between separate molecules, but throughout a whole piece of metal.

Key concepts in this lesson

Explain every physical property in the order structure → intermolecular forces → strength → property.
Stronger forces between molecules give a higher boiling point, a higher viscosity and a higher surface tension.
Compare molecules of similar size first; with similar electron counts, hydrogen bonds outweigh dipole–dipole forces, which outweigh London forces alone.
More O–H groups per molecule mean more hydrogen bonds and much higher viscosity, as in ethanol, ethylene glycol and glycerol.
Like dissolves like: a substance dissolves when the new solute–solvent attractions can replace the ones broken.