C. Structure and Properties of Matter · Structure and Properties of MatterIntermolecular Forces: Predicting Boiling Points and Behaviour
C3.4 C2.4 A1.8About 75 minutes
Learning goals- I can distinguish intramolecular bonds from intermolecular forces.
- I can identify London dispersion, dipole-dipole, hydrogen bonding and ion-dipole forces in a substance.
- I can use intermolecular forces to rank and explain boiling points.
- I can explain unusual properties of water, such as floating ice, using hydrogen bonding.
Words to know
intermolecular force
An attraction between molecules (or between molecules and ions).
London dispersion force
A weak attraction between temporary and induced dipoles; present between all molecules.
dipole-dipole force
An attraction between the oppositely charged ends of polar molecules.
hydrogen bond
A strong dipole-dipole attraction between H bonded to N, O or F and a lone pair on N, O or F of another molecule.
ion-dipole force
An attraction between an ion and the oppositely charged end of a polar molecule.
polarizability
How easily an electron cloud can be distorted to form a temporary dipole; it increases with the number of electrons.
Why is water a liquid?
Water (H₂O, 18 g/mol) boils at 100 °C. Carbon dioxide (44 g/mol) is a gas until it is cooled to −78 °C, when it turns straight into solid "dry ice". Methane (16 g/mol) boils at −162 °C. Molecules that are similar in size can have wildly different boiling points. The difference comes from the forces between molecules.
Remember from Grade 11: in SCH3U (Gases and Atmospheric Chemistry) you met three intermolecular forces: London dispersion, dipole-dipole and hydrogen bonding, and saw that melting or boiling a molecular substance separates molecules without breaking covalent bonds. This lesson uses your new skills (shape and polarity) to predict which forces act, compare them, and rank substances.
Intramolecular versus intermolecular
- Intramolecular forces act inside a molecule: covalent bonds (typically about 150 to 1000 kJ/mol).
- Intermolecular forces (IMFs) act between molecules. They are much weaker (typically less than about 40 kJ/mol) and are all electrostatic: attractions between partial positive and partial negative charges. Together they are often called van der Waals forces, after the Dutch physicist Johannes van der Waals.
The intermolecular forces
1. London dispersion forces (LDF) act between all molecules. Electrons are always moving, so at any instant a molecule's electron cloud may be lopsided, making a temporary dipole. It distorts (polarizes) the electron clouds of neighbours, creating induced dipoles that attract. LDFs are the only IMF between non-polar molecules. They are stronger when:
- molecules have more electrons (larger, heavier molecules have bigger, more easily distorted clouds), and
- molecules have more surface contact (long, straight molecules touch more than compact, round ones).
2. Dipole-dipole forces act between polar molecules: the δ+ end of one attracts the δ− end of another. They add to the London forces already present.
3. Hydrogen bonding is an especially strong dipole-dipole attraction. It occurs when H is covalently bonded to N, O or F in one molecule and is attracted to a lone pair on N, O or F of another. N, O and F are small and highly electronegative, so the H becomes strongly δ+ and is almost a bare proton that can get very close to the lone pair.
4. Ion-dipole forces act between an ion and a polar molecule, for example Na⁺ surrounded by the δ− oxygen ends of water molecules when salt dissolves.
| Force | Acts between | Relative strength | Example |
|---|
| London dispersion | all molecules | weak per contact, but large for big molecules | I₂, CH₄, C₈H₁₈ (octane) |
| Dipole-dipole | polar molecules | moderate | HCl, CH₂Cl₂, propanone (acetone) |
| Hydrogen bonding | molecules with H–N, H–O or H–F | strongest IMF between neutral molecules | H₂O, NH₃, HF, ethanol |
| Ion-dipole | ions and polar molecules | strong | NaCl dissolved in water |
Using IMFs to predict boiling points
Stronger IMFs → more energy needed to separate molecules → higher boiling point, higher melting point and lower vapour pressure.
- Compare molecules of similar size (similar number of electrons): hydrogen bonding > dipole-dipole > London only. Ethanol (CH₃CH₂OH, bp 78 °C) and methoxymethane (CH₃OCH₃, bp −24 °C) have the same formula, C₂H₆O, but only ethanol has an O–H to form hydrogen bonds.
- Compare molecules with the same kind of IMF: the one with more electrons has stronger London forces. Halogens: F₂ (−188 °C), Cl₂ (−34 °C), Br₂ (59 °C, a liquid), I₂ (a solid that boils at 184 °C).
- Watch for shape: pentane (a straight chain, bp 36 °C) boils higher than 2,2-dimethylpropane (a compact ball with the same formula, C₅H₁₂, bp about 10 °C).
- Size can win over polarity: London forces in a very large non-polar molecule can outweigh dipole forces in a small polar one. Iodine (I₂, non-polar, 106 electrons) boils at 184 °C, far higher than polar HCl (−85 °C).
Water: the hydrogen-bonding champion
Each water molecule has two H atoms and two lone pairs, so it can form up to four hydrogen bonds. This explains water's high boiling point, its high specific heat capacity (it takes a lot of energy to warm a lake), and something every Canadian sees each winter: in ice, hydrogen bonds hold the molecules in an open hexagonal network that takes up more space than the liquid. Ice is less dense than water, so it floats, and lakes freeze from the top down, leaving liquid water underneath where fish can survive the winter. Hydrogen bonds also hold the two strands of DNA together and help proteins keep their shapes.
Worked exampleExample 1: Identify the forces
List all the intermolecular forces in (a) CH₄ (b) HCl (c) NH₃ (d) CO₂.
- (a) CH₄ is non-polar: London only.
- (b) HCl is polar (no H–N/O/F): London + dipole-dipole.
- (c) NH₃ is polar and has H bonded to N: London + dipole-dipole + hydrogen bonding.
- (d) CO₂ is linear and non-polar: London only.
CH₄ and CO₂: London only; HCl: London and dipole-dipole; NH₃: all three.
Worked exampleExample 2: Rank by boiling point
Rank in order of increasing boiling point: C₃H₈ (propane), CH₃OCH₃ (methoxymethane), CH₃CH₂OH (ethanol). All have similar numbers of electrons (26, 26 and 26).
- Propane: non-polar → London only.
- Methoxymethane: bent at O, polar → London + dipole-dipole (no O–H, so no hydrogen bonding between its molecules).
- Ethanol: has O–H → London + dipole-dipole + hydrogen bonding.
- With similar London forces, the extra forces decide the order.
C₃H₈ (−42 °C) < CH₃OCH₃ (−24 °C) < CH₃CH₂OH (78 °C).
Worked exampleExample 3: Explaining a trend
HF boils at 20 °C, HCl at −85 °C, HBr at −67 °C and HI at −35 °C. Explain the trend from HCl to HI and why HF does not fit it.
- HCl → HBr → HI: the number of electrons increases (18, 36, 54), so London forces increase and so do boiling points, even though polarity decreases.
- HF is much smaller, but H is bonded to F, so HF molecules form hydrogen bonds.
- Hydrogen bonds are much stronger than the dipole-dipole forces in HCl, so HF boils far higher than the trend predicts.
London forces explain the rise from HCl to HI; hydrogen bonding explains HF's unusually high boiling point.
Just for funGecko feet can grip a glass window thanks partly to London dispersion forces. Each foot has millions of tiny hairs that split into even tinier tips, making so much close contact that these weak forces add up enough to hold the lizard's weight. Scientists are copying the idea to make dry, reusable adhesives.
Hands-on activityEvaporation race
You need: an adult helper, a medicine dropper or a teaspoon, a dark-coloured smooth plate or a piece of glass, a timer, and three liquids from the house: water, rubbing alcohol (isopropanol, 70% or 99%) and, if available, nail polish remover that lists acetone as its main ingredient. Safety (grown-up note): rubbing alcohol and acetone are flammable and their vapours should not be breathed for long: work in a well-ventilated room, away from flames and heaters, wear safety glasses if you have them, keep the bottles closed when not in use, and wash your hands afterward.
- Put one drop (or a small spot, the same size each time) of each liquid on the plate, well apart. Label them on a piece of tape.
- Start the timer. Record how long each spot takes to disappear completely.
- Rub a drop of each on the back of your hand. Which feels coldest?
What to look for: acetone usually evaporates fastest, then rubbing alcohol, then water (water may take several minutes). Water molecules form the most hydrogen bonds; isopropanol has one O–H per molecule; propanone (acetone) is polar but has no O–H to hydrogen bond with its neighbours. Weaker IMFs mean molecules escape more easily. The fastest-evaporating liquid feels coldest, because evaporating molecules carry heat away from your skin. Write a conclusion linking evaporation time to the IMFs present.
Your turn
1What IMFs act between Br₂ molecules?
London dispersion only (Br₂ is non-polar).
2What IMFs act between CH₃Cl molecules?
London + dipole-dipole (polar, no H on N, O or F).
3Which molecules can hydrogen bond with each other: CH₃OH, CH₃F, HF, CH₄?
CH₃OH (O–H) and HF (H–F). CH₃F has F but its H atoms are bonded to C, not F; CH₄ cannot.
4Explain why I₂ is a solid at room temperature while F₂ is a gas.
I₂ has far more electrons (106 versus 18), so its London forces are much stronger, holding molecules together as a solid.
5Rank in order of increasing boiling point: H₂O, H₂S, H₂Se.
H₂S (−60 °C) < H₂Se (−41 °C) < H₂O (100 °C). H₂Se has more electrons than H₂S; water forms hydrogen bonds.
6Why does ice float on water?
Hydrogen bonds hold ice molecules in an open hexagonal network, which is less dense than liquid water.
7Which has the higher boiling point, Br₂ (160 g/mol) or ICl (162 g/mol)? Explain.
ICl (97 °C versus 59 °C): similar numbers of electrons give similar London forces, but ICl is polar and also has dipole-dipole forces.
More practice: warm-up
1Which IMF exists between all molecules?
London dispersion forces.
2Hydrogen bonding requires H bonded to which three elements?
3Are IMFs stronger or weaker than covalent bonds?
4When water boils, are O–H bonds broken?
No, only intermolecular forces (mainly hydrogen bonds) are overcome.
5What IMF holds Na⁺ and water molecules together in salt water?
More practice: on your own
1Explain why pentane boils at 36 °C but 2,2-dimethylpropane (same formula) boils at about 10 °C.
Pentane is a long chain with more surface contact between molecules, so its London forces are stronger. 2,2-dimethylpropane is compact and nearly spherical, with less contact.
2Rank by boiling point: CH₄, C₂H₆, C₃H₈, C₄H₁₀ (all non-polar).
Increasing electrons → stronger London forces: CH₄ < C₂H₆ < C₃H₈ < C₄H₁₀ (−162, −89, −42, about −1 °C).
3Ammonia (NH₃) boils at −33 °C but phosphine (PH₃) at −88 °C, even though PH₃ has more electrons. Explain.
NH₃ forms hydrogen bonds (H on N); PH₃ cannot (P is not N, O or F), so it has only weaker London and dipole forces.
4Which is more soluble in water, methanol (CH₃OH) or methane (CH₄)? Explain.
Methanol: it is polar and forms hydrogen bonds with water. Methane is non-polar and attracts water only through weak London forces.
5Explain why lakes in Canada freeze from the top down, and why this matters to living things.
Water is densest at about 4 °C, and ice is less dense than liquid water because of its hydrogen-bonded open structure. Ice forms and floats on top, insulating the water below, which stays liquid so fish and other organisms survive.
Challenge
1Propanone (acetone, CH₃COCH₃, 58 g/mol) boils at 56 °C, while butane (C₄H₁₀, 58 g/mol) boils at about −1 °C. Explain, naming every IMF present in each.
Butane is non-polar: London only. Propanone has a polar C=O bond and a trigonal planar carbon, so it is polar: London + dipole-dipole. With similar sizes (similar London forces), the dipole-dipole forces give propanone a much higher boiling point. Propanone has no O–H, so its molecules do not hydrogen bond with each other.
2Water (bp 100 °C) has a higher boiling point than hydrogen fluoride (bp 20 °C), although an H–F hydrogen bond is stronger than an O–H···O hydrogen bond. Explain.
Each H₂O molecule has 2 H atoms and 2 lone pairs, so on average it can form about twice as many hydrogen bonds (up to 4) as each HF molecule (1 H atom), which forms chains. More hydrogen bonds per molecule in total outweighs the slightly stronger individual H–F bonds.
3Glycerol, C₃H₅(OH)₃, boils at 290 °C, while propan-1-ol, C₃H₇OH, boils at 97 °C. Explain the difference.
Both have similar carbon chains, but glycerol has three O–H groups per molecule, so it forms many more hydrogen bonds than propan-1-ol (one O–H). It also has more electrons. Much stronger total IMFs give a much higher boiling point.
Quiz
1. Which intermolecular force is present in ALL molecular substances?
2. Which substance forms hydrogen bonds between its molecules?
3. Which halogen has the highest boiling point?
4. When liquid water boils, what is overcome?
5. Why does H₂O boil much higher than H₂S?
6. Ethanol and methoxymethane both have the formula C₂H₆O. Which statement is correct?
Big idea (C3.4): the forces between particles decide many physical properties. Students should identify each IMF from a molecule's structure, then compare: similar size → the type of IMF decides; same type → number of electrons and shape decide.
How to help: the evaporation race uses household liquids; please supervise closely, ventilate, and keep flames away. Ask your teen to explain the result using IMFs.
Common difficulties:
- Saying covalent bonds break when a liquid boils.
- Claiming any molecule with H and O can hydrogen bond (the H must be on N, O or F).
- Forgetting that London forces exist in polar molecules too, and can be the largest contribution in big molecules.