AP® Chemistry review sheet from Aim for Five (aimforfive.com/chem/units/3/3-1)
Unit 3 · Topic 3.1
3.1 Intermolecular and Interparticle Forces
Intermolecular forces are the attractions between separate particles. They're much weaker than the covalent bonds inside molecules, but they decide boiling points, melting points, vapor pressures and solubility. London dispersion forces act between all particles, and polar molecules, hydrogen bonding and ions add stronger attractions on top.
Key terms
- London dispersion forces
- polarizability
- dipole-dipole forces
- hydrogen bonding
- ion-dipole forces
- dipole-induced dipole forces
Intermolecular versus intramolecular
Intramolecular forces are the bonds inside a molecule, like the O–H bonds in water. Intermolecular forces (IMFs) are attractions between neighboring molecules. When water boils, the molecules separate from each other, but no O–H bonds break. Boiling, melting and evaporating overcome IMFs only.
All IMFs are Coulombic: partial or full charges on one particle attract opposite charges on another.
London dispersion forces
Electrons are always moving. At any instant, a molecule's electron cloud can be lopsided, creating a temporary dipole. That temporary dipole pushes on the electrons in a neighboring molecule and creates an induced dipole there. The two attract. These fleeting attractions are London dispersion forces (LDFs), and they exist between all molecules and atoms, polar or not.
Two things make dispersion forces stronger. The first is polarizability: how easily a molecule's electron cloud can be pushed out of shape. Big electron clouds with many electrons distort easily, and so do the electrons in pi bonds. The second is contact area: long, straight molecules touch their neighbors along more of their length than compact, ball-shaped ones.
For large molecules, LDFs are often the strongest IMF present, even if the molecule is polar. Don't confuse 'London dispersion forces' with 'van der Waals forces', which is a broader term covering several kinds of IMFs.
Forces involving dipoles and ions
- Dipole-dipole: between polar molecules. The δ+ end of one molecule attracts the δ− end of another. These act in addition to LDFs, so a polar molecule usually has stronger IMFs than a nonpolar molecule of about the same size.
- Dipole-induced dipole: between a polar molecule and a nonpolar one. The polar molecule pushes the nonpolar molecule's electrons to one side, creating a dipole it can attract. The bigger the dipole, and the easier the other cloud is to distort, the stronger the pull.
- Hydrogen bonding: a strong special case. It forms when an H atom covalently bonded to N, O or F is attracted to a lone pair on an N, O or F atom in another molecule (or another part of the same large molecule).
- Ion-dipole: between an ion and a polar molecule, such as Na⁺ and water. Usually stronger than dipole-dipole forces.
Orientation matters
Dipole and ion attractions depend on which way the molecules point. A cation is attracted to the partially negative end of a polar molecule, so water molecules point their oxygen atoms toward Na⁺. An anion attracts the partially positive end, so water points its hydrogens toward Cl⁻.
Large biological molecules like proteins and DNA use these same attractions between different parts of the same molecule. Hydrogen bonds hold the two strands of DNA together and help fold proteins into their shapes.
Comparing substances
To compare boiling points, identify every IMF in each substance, then decide which substance has the stronger total attraction. Stronger IMFs mean more energy is needed to separate the particles, so a higher boiling point. If the molecules are similar in size, the one with hydrogen bonding or dipole-dipole forces wins. If one molecule is much bigger, its extra dispersion forces can outweigh another molecule's polarity.
Worked examples
Try each one yourself first, then open the solution.
- Example 1
Same formula, different IMFs
Ethanol (CH₃CH₂OH) and dimethyl ether (CH₃OCH₃) both have the formula C₂H₆O. Ethanol boils at about 78 °C and dimethyl ether at about −24 °C. Explain the difference.
Show the solutionHide the solution
- Step 1: Same formula means the same number of electrons and similar size, so their London dispersion forces are similar.
- Step 2: Ethanol has an H bonded directly to O, so ethanol molecules hydrogen bond to each other.
- Step 3: In dimethyl ether every H is bonded to C, so it can't hydrogen bond with itself. It is polar, so it has dipole-dipole forces, but those are weaker.
- Step 4: Stronger IMFs in ethanol take more energy to overcome, so ethanol boils at a higher temperature.
Answer: Ethanol molecules form hydrogen bonds with each other (H on O), while dimethyl ether has only dipole-dipole and dispersion forces, so ethanol's IMFs are stronger and its boiling point is higher.
- Example 2
When dispersion beats polarity (classic trap)
CH₃Cl is polar and boils at about −24 °C. CCl₄ is nonpolar and boils at about 77 °C. Explain why the nonpolar compound has the higher boiling point.
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- Step 1: CH₃Cl has dipole-dipole forces plus dispersion forces. CCl₄ has only dispersion forces.
- Step 2: CCl₄ has four Cl atoms (74 electrons in all) versus one Cl in CH₃Cl (26 electrons). Its electron cloud is much larger and more polarizable.
- Step 3: The much stronger dispersion forces in CCl₄ outweigh the dipole-dipole forces in CH₃Cl.
- Step 4: The rule 'polar molecules have higher boiling points' only works when the molecules are similar in size.
Answer: CCl₄ has many more electrons and a more polarizable electron cloud, so its London dispersion forces are stronger than CH₃Cl's combined dipole-dipole and dispersion forces.
- Example 3
Shape and contact area
Pentane (a straight five-carbon chain) boils at 36 °C. Neopentane, also C₅H₁₂, is nearly spherical and boils at 9.5 °C. Explain.
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- Step 1: Both are nonpolar with the same number of electrons, so they differ only in shape.
- Step 2: Straight pentane molecules can lie alongside each other with lots of surface in contact. Compact neopentane molecules touch only at small areas.
- Step 3: More contact area means stronger dispersion forces overall, so pentane needs more energy to boil.
Answer: Pentane's straight chains have more contact area, giving stronger London dispersion forces and a higher boiling point.
Common mistakes
- Saying boiling breaks covalent bonds. Boiling a molecular substance only overcomes IMFs between molecules.
- Counting any H in the molecule as hydrogen bonding. The H must be bonded directly to N, O or F.
- Forgetting that every molecule has London dispersion forces, including polar ones.
- Ranking by polarity alone when molecules differ a lot in size. Compare electron counts too.
On the exam
- Free-response questions ask you to explain a difference in boiling point, vapor pressure or solubility. Name the specific IMFs in both substances, say which is stronger and why (structure), and link that to the property. Using vague 'stronger bonds' language loses points.
- Expect particle diagrams where you identify a hydrogen bond. It's drawn between the H of one molecule and the N, O or F lone pair of another, not along the covalent bond.
Connected topics
Videos
Check yourself
4 questions on 3.1 Intermolecular and Interparticle Forces. Pick an answer to see if you got it, and why.
The normal boiling points of the halogens increase in the order F₂ < Cl₂ < Br₂ < I₂. Which of the following best explains this trend?
| Compound | Structure | Molar mass (g/mol) | Normal boiling point (°C) |
|---|---|---|---|
| Propane | CH₃CH₂CH₃ | 44.10 | −42 |
| Dimethyl ether | CH₃OCH₃ | 46.07 | −24 |
| Ethanol | CH₃CH₂OH | 46.07 | 78 |
Approximate boiling points at 1 atm
Which of the following best explains why ethanol has a much higher boiling point than dimethyl ether, even though they have the same molar mass?
Which of the following best explains why dimethyl ether boils at a higher temperature than propane?
In which of the following pure substances do hydrogen bonds form between molecules?
0 of 4 answered