Ethanol has a higher boiling point than methanol because it is a larger, heavier molecule with stronger attractive forces between its molecules. Methanol boils at 65 °C while ethanol boils at 78 °C, a 13-degree difference that comes down to one extra carbon and two extra hydrogen atoms in ethanol’s chain. That small structural addition increases the energy needed to pull ethanol molecules apart and send them into the gas phase.
The Role of Molecular Size and Mass
Methanol (CH₃OH) has a molecular weight of 32, while ethanol (CH₃CH₂OH) comes in at 46. That 44% jump in mass matters because heavier molecules with more atoms have more electrons, and more electrons means stronger temporary attractive forces between neighboring molecules. These attractions, called London dispersion forces, exist in every substance, but they grow stronger as molecules get bigger and heavier. The electron cloud around a larger molecule is easier to distort, creating fleeting charge imbalances that pull nearby molecules closer together.
This pattern holds across the entire alcohol family. Methanol boils at 65 °C, ethanol at 78 °C, and propanol (one more carbon) at 97 °C. Each added carbon-and-hydrogen unit increases the surface area where neighboring molecules can interact, stacking on additional London dispersion forces and pushing the boiling point higher in a predictable staircase.
Hydrogen Bonding in Both Molecules
Both methanol and ethanol can form hydrogen bonds because they each contain an oxygen-hydrogen (O-H) group. Hydrogen bonding is a strong type of intermolecular attraction that occurs when the hydrogen on one molecule’s O-H group is drawn toward the oxygen on a neighboring molecule. This is why both alcohols boil at much higher temperatures than similarly sized molecules that lack an O-H group.
What’s interesting is that hydrogen bonding isn’t identical in the two liquids. Research published in The Journal of Physical Chemistry B measured the energy of hydrogen bonds in both alcohols and found that the bonding enthalpy is about 12.8 kJ/mol in methanol and 16.8 kJ/mol in ethanol. So each hydrogen bond in ethanol is roughly 30% stronger than in methanol. The larger carbon chain in ethanol subtly affects how electrons are distributed around the oxygen atom, which slightly strengthens each hydrogen bond the molecule forms.
How These Forces Affect Vaporization
To boil a liquid, you need to supply enough energy to overcome all the intermolecular forces holding its molecules together. The total energy required to vaporize a substance is captured by a measurement called the heat of vaporization. For methanol, that value is 8.95 kcal/mol. For ethanol, it’s 10.11 kcal/mol. Ethanol requires about 13% more energy per molecule to escape into the gas phase, which is why you need to heat it to a higher temperature before it boils.
You can also see this difference at room temperature by comparing vapor pressures. In a classic chemistry demonstration, methanol injected into a sealed tube depresses a mercury column to 634 mm, while ethanol only depresses it to 688 mm. The smaller depression for ethanol means fewer ethanol molecules escape into the vapor at the same temperature. Methanol evaporates more readily because its weaker intermolecular forces make it easier for individual molecules to break free from the liquid surface.
Why Dispersion Forces Matter More Than You’d Think
It’s tempting to attribute the boiling point difference entirely to hydrogen bonding, since that’s the strongest intermolecular force present in alcohols. But the hydrogen bonding capability of methanol and ethanol is actually quite similar. Both molecules have one O-H group, and both can form the same types of hydrogen bond networks with their neighbors. The dominant reason ethanol boils higher is the cumulative boost in London dispersion forces from the extra CH₂ group.
This is a general principle in chemistry: for similar substances, London dispersion forces get stronger with increasing molecular mass and surface area. Heavier species are more polarizable, meaning their electron clouds shift more easily in response to neighboring molecules. Ethanol’s additional carbon gives it roughly 40% more polarizable electron density than methanol. While each individual dispersion interaction is weak, they add up across the entire molecular surface and collectively raise the amount of thermal energy needed to vaporize the liquid.
Putting It All Together
The 13 °C boiling point gap between methanol and ethanol comes from two reinforcing effects. First, ethanol’s larger size and greater electron count produce stronger London dispersion forces, which account for most of the difference. Second, ethanol’s hydrogen bonds are individually somewhat stronger than methanol’s, adding a smaller but real contribution. Both effects increase the total energy needed to separate ethanol molecules from each other, which is why ethanol stays liquid at temperatures where methanol has already started to boil. This same logic extends up the alcohol series: each additional carbon unit brings more dispersion forces, stronger molecular cohesion, and a higher boiling point.