Boiling Temperature Trends
A concise revision guide to boiling temperature trends explained by intermolecular forces: alkanes with chain length and branching, alcohols compared with alkanes, and the hydrogen halides HF to HI, for Edexcel International A Level Chemistry.
- 7.5i
- 7.5ii
- 7.5iii
- 7.5iv
What these spec points say
- 7.5i understand, in terms of intermolecular forces, physical properties shown by substances, including: the trends in boiling temperatures of alkanes with increasing chain length
- 7.5ii understand, in terms of intermolecular forces, physical properties shown by substances, including: the effect of branching in the carbon chain on the boiling temperatures of alkanes
- 7.5iii understand, in terms of intermolecular forces, physical properties shown by substances, including: the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
- 7.5iv understand, in terms of intermolecular forces, physical properties shown by substances, including: the trends in boiling temperatures of the hydrogen halides HF to HI
GCSE Recap: Size, State and Crude Oil
Three quick questions on the GCSE ideas about states and alkanes that this page builds on.
How Intermolecular Forces Set the Boiling temperature
A substance boils when its molecules gain enough energy to overcome the intermolecular forces between them, so the stronger the forces, the higher the boiling temperature.
Every trend on this page is explained by asking which forces act between the molecules and what changes their strength: the number of electrons, the shape of the molecule, the presence of a permanent dipole and the presence of hydrogen bonds.
Students should always name the force, state what has changed and finish with the energy statement: “more energy is needed to separate the molecules”. A trend described without a force earns nothing.
Exam answer model: Identify the force → say why it is stronger or weaker → “so more (or less) energy is needed to overcome the forces between the molecules” → higher (or lower) boiling temperature.
Alkanes: Increasing Chain Length
Alkanes are non-polar, so the only force between their molecules is London forces.
As the chain lengthens, each molecule has more electrons and a larger surface area of contact with its neighbours. So the London forces strengthen and the boiling temperature rises steadily.
Methane, ethane, propane and butane are gases at room temperature; pentane to heptadecane are liquids; longer alkanes are waxy solids.
| Alkane | Formula | Electrons | boiling temperature / °C |
|---|---|---|---|
| Methane | CH₄ | 10 | -162 |
| Ethane | C₂H₆ | 18 | -89 |
| Propane | C₃H₈ | 26 | -42 |
| Butane | C₄H₁₀ | 34 | -1 |
| Pentane | C₅H₁₂ | 42 | 36 |
| Hexane | C₆H₁₄ | 50 | 69 |
| Heptane | C₇H₁₆ | 58 | 98 |
| Octane | C₈H₁₈ | 66 | 126 |
The boiling temperature of the alkanes rises with chain length because London forces strengthen; at a fixed formula, branching lowers it.
The rise is steepest at the start of the series, because adding one CH₂ group to methane increases its electron count by 80%, while adding one to heptane increases it by 14%. This is why the curve flattens as the chain lengthens.
Exam sentence: Hexane has a higher boiling temperature than butane because hexane molecules have more electrons and a larger surface area of contact, so the London forces between the molecules are stronger and more energy is needed to overcome them.
Quick Check: Two Alkanes You Have Not Met
Drag the words and numbers into the gaps to compare two longer alkanes.
Alkanes: The Effect of Branching
Isomers have the same number of electrons, so any difference in their boiling temperatures must come from shape.
Branched alkanes are more compact and nearly spherical, so neighbouring molecules cannot approach each other as closely or over as large an area. The London forces are weaker and the boiling temperature is lower.
| C₅H₁₂ isomer | Shape | boiling temperature / °C |
|---|---|---|
| Pentane | straight chain | 36 |
| 2-Methylbutane | one branch | 28 |
| 2,2-Dimethylpropane | four groups on one carbon, almost spherical | 10 |
Straight chains can lie side by side along their whole length, like pencils in a box, giving a large area over which instantaneous dipoles can induce dipoles in the neighbour. A ball-shaped molecule touches its neighbours only at a few points.
Common mistake: Saying branched isomers have “fewer electrons” or “weaker bonds”. They have the same electrons and the same C-C and C-H bonds; only the contact area between molecules differs.
Quick Check: Pick the Accurate Statement
Choose the one accurate statement about branched and straight-chain isomers in each round.
Alcohols Compared With Alkanes
Alcohols have much higher boiling temperatures, and so a much lower volatility, than alkanes with a similar number of electrons. The O-H group lets alcohol molecules form hydrogen bonds with each other, which alkanes cannot do.
The fair comparison is between molecules of similar size, so that their London forces are similar and the difference can be attributed to the hydrogen bonds.
| Compound | Formula | Electrons | Hydrogen bonding? | boiling temperature / °C |
|---|---|---|---|---|
| Propane | C₃H₈ | 26 | no | -42 |
| Ethanol | C₂H₅OH | 26 | yes | 78 |
| Butane | C₄H₁₀ | 34 | no | -1 |
| Propan-1-ol | C₃H₇OH | 34 | yes | 97 |
Alcohols boil about 100 °C higher than alkanes with the same number of electrons because hydrogen bonds act between alcohol molecules in addition to London forces.
The 120 °C gap between propane and ethanol is larger still than the 103 °C gap that a permanent dipole produced between F₂ and HCl.
This is one way to remember that hydrogen bonds are stronger than permanent dipole-permanent dipole interactions.
Alcohols also have permanent dipole-permanent dipole interactions from the polar C-O and O-H bonds, but these are minor next to the hydrogen bonds, which are several times stronger.
Exam sentence: Ethanol has a higher boiling temperature than propane because ethanol molecules form hydrogen bonds between the O-H groups, which are stronger than the London forces between propane molecules, so more energy is needed to separate ethanol molecules.
Quick Check: Same Electrons, Very Different Liquids
Choose the best explanation for the gap between two liquids with the same number of electrons.
The Hydrogen Halides HF to HI
The boiling temperatures of the hydrogen halides do not fall smoothly down the group. HF is anomalously high because it forms hydrogen bonds; from HCl to HI the boiling temperature rises because the number of electrons increases and the London forces strengthen, even though the permanent dipole becomes smaller.
| Hydrogen halide | Electrons | Electronegativity of halogen | Hydrogen bonding? | boiling temperature / °C |
|---|---|---|---|---|
| HF | 10 | 4.0 | yes | 20 |
| HCl | 18 | 3.0 | no | -85 |
| HBr | 36 | 2.8 | no | -67 |
| HI | 54 | 2.5 | no | -35 |
The HCl to HI trend is a useful test of understanding, because the dipole and the London forces pull in opposite directions.
HCl has the largest permanent dipole of the three, yet it has the lowest boiling temperature. The extra electrons in HBr and HI produce London forces that outweigh the shrinking permanent dipole-permanent dipole interactions.
Remember: For molecules with a large number of electrons, London forces usually dominate.
The Group 4 to 7 hydride graph on the Anomalous Properties of Water page shows the HF anomaly alongside those of water and ammonia; the same drawing serves both pages.
Exam sentence: HF has a much higher boiling temperature than HCl because HF molecules form hydrogen bonds; HI has a higher boiling temperature than HCl because HI molecules have more electrons, so the London forces between them are stronger.
Quick Check: Explain a Trend That Fights Itself
Write a short explanation for a series in which the dipole and the electrons pull in opposite directions.
Putting the Trends Together
When two substances are compared, work through the forces in order of strength and stop at the first difference.
| Question to ask | If the answer differs | Example |
|---|---|---|
| Does one form hydrogen bonds and the other not? | The hydrogen-bonded one boils much higher, provided the two molecules are of similar size | ethanol 78 °C versus propane -42 °C |
| Does one have a permanent dipole and the other not? | The polar one boils higher | HCl -85 °C versus F₂ -188 °C |
| Which has more electrons? | The one with more electrons boils higher | HI -35 °C versus HCl -85 °C |
| Which is less branched? | The straighter molecule boils higher | pentane 36 °C versus 2,2-dimethylpropane 10 °C |
The first question settles the order only when the two molecules are of a similar size. A long non-polar chain has so many more electrons that it can boil above a short hydrogen-bonded molecule, so check the sizes before you trust the hydrogen bonding step.
Worked example. Put propane, pentane and ethanol in order.
Step 1: Ethanol is the only one with an O-H group, so it forms hydrogen bonds and boils highest.
Step 2: Neither propane nor pentane has a permanent dipole, so compare electrons: pentane has 42 and propane 26, so pentane boils above propane.
Answer: The order is propane (-42 °C), pentane (36 °C), ethanol (78 °C).
Key idea: Hydrogen bonds outweigh dipoles, dipoles outweigh London forces for molecules of similar size, and London forces grow with electrons and contact area.
Compare two substances by checking hydrogen bonding, then a permanent dipole, then the number of electrons, then branching, and stop at the first difference.
Quick Check: Put Six Liquids in Order
Put the six substances in order of boiling temperature, lowest first.
Quick Check: Which Force Decides It?
Pick the one accurate statement in each round to check every trend on the page.
Common Exam Points
Explain the trend in boiling temperature of the alkanes
More electrons and larger surface contact, stronger London forces, more energy to separate molecules.
Explain why 2,2-dimethylpropane boils lower than pentane
Same electrons; more compact shape; smaller area of contact; weaker London forces.
Explain why alcohols are less volatile than alkanes
Hydrogen bonds between O-H groups, absent in alkanes, need more energy to overcome.
Explain the order HF > HI > HBr > HCl
HF: hydrogen bonds; HCl to HI: increasing electrons, stronger London forces, outweighing the falling dipole.
FAQs
Use these quick answers to check the boiling temperature comparisons that Edexcel International questions set.
Why does boiling temperature rise along the alkane series?
Each extra carbon adds electrons and surface area, so the London forces between molecules get stronger and more energy is needed to separate them. The rise flattens along the series because each extra CH₂ group is a smaller fraction of the total.
Why do branched alkanes boil lower than straight-chain isomers?
Isomers have the same number of electrons, so the difference is shape. A branched molecule is compact and touches its neighbours over a smaller area, so the London forces are weaker: pentane boils at 36 °C and 2,2-dimethylpropane at 10 °C.
Why are alcohols less volatile than alkanes?
The O-H group lets alcohol molecules hydrogen bond with each other, which alkanes cannot do. Ethanol and propane both have 26 electrons, yet ethanol boils 120 °C higher.
Why does HF boil higher than HCl, HBr and HI?
HF molecules form hydrogen bonds, which the other hydrogen halides cannot. From HCl to HI the trend then rises again as the number of electrons increases.
Which factor wins when they disagree?
Hydrogen bonding outweighs a permanent dipole, and a permanent dipole outweighs London forces for molecules of similar size. Once the electron count differs a great deal, London forces dominate, which is why HI boils higher than HCl.
Copyright notice: This OLS revision content, including the explanations, layout, diagrams, tables and embedded learning structure, is authored for Online Learning System by Dr. Mohammed Al-Fatah. It may not be copied, reproduced, redistributed or adapted without written permission.
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