Difficulty: Introductory | Prerequisites: General chemistry (atomic structure, electronegativity, polarity of bonds).
This topic sits at the intersection of structure and physical properties, which is one of the central themes in organic chemistry. You are learning to look at a molecule's structure and predict how it will behave in the real world, starting with something measurable: its boiling point. The key idea is that intermolecular forces (the attractions between molecules, not within them) govern how much energy you need to pull molecules apart from the liquid phase into the gas phase. If you are comfortable with electronegativity, bond polarity, and drawing Lewis structures, you have what you need.
Stronger intermolecular forces mean higher boiling points. Hydrogen bonding (alcohols) beats dipole-dipole (ketones), which beats London dispersion forces alone (alkanes). Within any one functional-group family, longer carbon chains raise the boiling point because more surface area means stronger London dispersion forces.
Boiling point
The temperature at which a liquid's vapour pressure equals the surrounding atmospheric pressure, causing a phase transition from liquid to gas. Think of it as the temperature where molecules have enough kinetic energy to escape the liquid entirely.
Intermolecular forces (IMFs)
Attractive forces that act between molecules (not within a molecule). In simple terms, these are the "glue" holding molecules near each other in the liquid phase.
Hydrogen bonding (HB)
A strong type of dipole-dipole attraction that occurs when hydrogen is bonded to a highly electronegative atom (N, O, or F) and interacts with a lone pair on another electronegative atom. Think of it as an especially strong dipole-dipole interaction, limited to molecules with O-H, N-H, or F-H bonds.
Dipole-dipole interactions (DD)
Attractive forces between the positive end of one polar molecule and the negative end of another. In simple terms, polar molecules line up positive-to-negative, like weak magnets.
London dispersion forces (LDF)
Weak, temporary attractions that arise from momentary fluctuations in electron density around any molecule, polar or nonpolar. Think of it as the baseline attraction that every molecule has, including nonpolar ones. The larger the molecule's surface area, the stronger these forces become.
Induced dipole (ID)
A temporary dipole created in a nonpolar molecule when a nearby polar molecule or ion distorts its electron cloud. In simple terms, a polar neighbour can temporarily "pull" the electrons in a nonpolar molecule, creating a short-lived attraction.
Vapour pressure
The pressure exerted by a vapour in equilibrium with its liquid at a given temperature. Liquids with strong IMFs have low vapour pressure because fewer molecules escape the surface.
Capillary tube method
A technique for determining boiling point by placing a liquid sample in a sealed capillary tube, heating until the liquid evaporates, then cooling and recording the temperature at which the vapour condenses and re-enters the tube.
The strength of intermolecular forces directly determines a compound's boiling point. Stronger forces require more energy to overcome, so the boiling point is higher.
In order of decreasing strength:
Hydrogen bonding (strongest of the three considered here). Requires an O-H, N-H, or F-H bond in the molecule.
Dipole-dipole interactions (moderate). Present in all polar molecules. Ketones have a polar C=O bond but no O-H bond, so they show dipole-dipole but not hydrogen bonding.
London dispersion forces (weakest). Present in all molecules, including nonpolar ones. This is the only IMF available to alkanes.
All three types operate simultaneously when they can. Alcohols, for example, exhibit hydrogen bonding, dipole-dipole interactions, and London dispersion forces at the same time.
Compare compounds with roughly the same number of carbons but different functional groups.
Pentane (alkane, bp 36 °C) < 2-butanone (ketone, bp 80 °C) < 1-butanol (alcohol, bp 118 °C).
Pentane has only LDF. 2-butanone adds dipole-dipole. 1-butanol adds hydrogen bonding on top of dipole-dipole and LDF.
The ranking pentane (1, lowest) < 2-butanone (2) < 1-butanol (3, highest) follows directly from the IMF hierarchy.
Within alkanes, alcohols, or ketones, boiling point increases with chain length.
Longer chains have more surface area, which increases the strength of London dispersion forces.
Alkanes: pentane 36 °C, hexane 69 °C, heptane 98 °C, octane 126 °C.
Alcohols: ethanol 78 °C, 1-propanol 97 °C, 1-butanol 118 °C, 1-pentanol 138 °C.
Ketones: 2-propanone (acetone) 56 °C, 2-butanone (MEK) 80 °C, 2-pentanone 102 °C, 2-hexanone 127 °C.
In every family, adding one CH₂ group raises the boiling point by roughly 20 to 30 °C.
At the same number of carbons, alcohols boil highest, ketones in the middle, alkanes lowest.
Example at five carbons: pentane 36 °C (LDF only), 2-pentanone 102 °C (DD + LDF), 1-pentanol 138 °C (HB + DD + LDF).
The gap between alkane and ketone is larger than the gap between ketone and alcohol, because adding dipole-dipole forces on top of LDF is a bigger jump than adding hydrogen bonding on top of existing dipole-dipole.
Compound | Functional group | IMFs present |
|---|---|---|
Pentane | Alkane | LDF (+ induced dipole near polar molecules) |
1-Butanol | Alcohol | Hydrogen bonding + dipole-dipole + LDF |
2-Butanone | Ketone | Dipole-dipole + LDF |
Fill a capillary tube with about 1 mL (or 20 drops) of sample.
Place the tube in a silicone oil bath on a hot plate/stirrer. Heat at roughly 5 °C per minute.
High end of boiling range: recorded when liquid cools and begins to re-enter the capillary tube after boiling.
Low end of boiling range: recorded when liquid stops rising inside the tube.
Each sample is measured twice: once heating quickly and once heating slowly, to bracket the true boiling point.
[embed: node/0d5a24a9-c8a0]
The chart makes both trends visible at once: each line rises with carbon count (LDF increasing), and the three lines sit at different heights (functional group IMFs).
This is why petrol (gasoline) evaporates so quickly at room temperature: it is mostly short-chain alkanes with weak IMFs and low boiling points. Conversely, the reason rubbing alcohol (isopropanol) lingers on your skin longer than nail-polish remover (acetone) is that the alcohol's hydrogen bonding keeps more molecules in the liquid phase. Engineers designing distillation columns for petrochemical refining rely on exactly these boiling-point differences to separate crude oil into useful fractions.
Students often confuse intermolecular forces with intramolecular bonds (covalent, ionic). Boiling does not break covalent bonds within the molecule. It only overcomes the attractions between molecules.
Students sometimes think London dispersion forces only exist in nonpolar molecules. Every molecule has LDF, including polar ones and those with hydrogen bonding. LDF is the baseline; other IMFs are added on top.
Students frequently say ketones can hydrogen bond. Ketones have a C=O group that can accept a hydrogen bond from a donor (like water), but pure ketones cannot donate a hydrogen bond because they lack an O-H or N-H bond. In a pure ketone sample, hydrogen bonding does not occur.
Students sometimes assume that higher molar mass always means a higher boiling point across different functional groups. This is only reliable within the same family. A lightweight alcohol (ethanol, 46 g/mol, bp 78 °C) boils well above a heavier alkane (pentane, 72 g/mol, bp 36 °C) because functional group IMFs matter more than mass alone.
⚠️ Expect a question asking you to rank compounds by boiling point and justify the ranking using IMF types. This is the most common exam format for this topic.
⚠️ Be ready to identify which IMFs are present in a given compound from its structure. Know the difference between "has a polar bond" and "can hydrogen bond."
⚠️ You may be asked to explain why two compounds with similar molar masses have different boiling points. The answer is always about which IMFs each compound can form.
⚠️ The lab practical may ask you to identify an unknown compound by comparing your measured boiling point range to literature values in the data tables.
True or false: London dispersion forces are present in ethanol. (True. Every molecule has LDF.)
Fill in the blank: The boiling point is the temperature at which a liquid's ______ equals atmospheric pressure. (vapour pressure)
True or false: A ketone can hydrogen bond with itself in a pure sample. (False. Ketones lack an O-H or N-H donor.)
Fill in the blank: Within a homologous series, increasing chain length increases boiling point because of stronger ______. (London dispersion forces)
True or false: Octane has a higher boiling point than pentane because octane has stronger dipole-dipole interactions. (False. Both are nonpolar. Octane's higher bp is due to stronger LDF from greater surface area.)
Q: Rank the following in order of increasing boiling point and explain: pentane, 1-butanol, 2-butanone.
A: Pentane (lowest) < 2-butanone < 1-butanol (highest). Pentane is nonpolar and has only LDF. 2-butanone is polar (C=O) and has dipole-dipole interactions plus LDF. 1-butanol has an O-H group, so it has hydrogen bonding, dipole-dipole, and LDF. Stronger IMFs require more energy to overcome, giving a higher boiling point.
Q: Hexane (molar mass 86.18 g/mol) boils at 69 °C, while ethanol (46.07 g/mol) boils at 78 °C. Explain why ethanol boils higher despite having a lower molar mass.
A: Ethanol can form hydrogen bonds through its O-H group, which are much stronger than the London dispersion forces available to hexane. The strength of IMFs matters more than molar mass when comparing across functional groups.
Q: Why does the boiling point of alkanes increase from pentane to octane?
A: All alkanes are nonpolar and rely on London dispersion forces. As the carbon chain gets longer, the molecule has more surface area for electron cloud contact, making LDF stronger. Stronger LDF means a higher boiling point.
Q: A student claims that 2-butanone can hydrogen bond because it contains an oxygen atom. Is this correct? Explain.
A: In a pure sample, no. Hydrogen bonding requires a hydrogen bonded directly to N, O, or F as a donor, plus a lone pair on another electronegative atom as an acceptor. 2-butanone has the acceptor (C=O oxygen with lone pairs) but no donor (no O-H or N-H bond). It can only act as a hydrogen bond acceptor when mixed with a donor molecule such as water.
Q: In the capillary tube method, what observation marks the upper end of the boiling point range?
A: The upper end is recorded when the liquid cools after boiling and begins to re-enter (condense back into) the capillary tube. The temperature at that moment is the high end of the boiling range.
This connects directly to solubility and miscibility (Chapter 4 in most organic texts): the same IMFs that govern boiling point also determine whether two liquids mix. "Like dissolves like" is a rule about matching IMF types.
It also ties into distillation techniques you will encounter later in the course. Fractional distillation separates mixtures precisely because components with different IMFs have different boiling points.
In physical chemistry and thermodynamics, boiling points relate to enthalpy of vaporisation (ΔH_vap), which quantifies the energy needed to overcome IMFs during the liquid-to-gas transition.
boiling point trends, intermolecular forces, IMF, hydrogen bonding, H-bonding, dipole-dipole interactions, London dispersion forces, van der Waals forces, induced dipole, vapour pressure, vapor pressure, boiling point determination, capillary tube method, alkane boiling points, alcohol boiling points, ketone boiling points, pentane, hexane, heptane, octane, ethanol, 1-propanol, 1-butanol, 1-pentanol, 2-butanone, MEK, acetone, organic chemistry lab, CHM 25502, physical constants, molar mass and boiling point, chain length and boiling point, functional group effect on boiling point, nonpolar molecules, polar molecules