No, dimethyl ether does not form hydrogen bonds with itself. Its oxygen atom has lone pairs of electrons, but none of its hydrogen atoms are bonded directly to that oxygen. Since hydrogen bonding requires a hydrogen attached to a highly electronegative atom like oxygen, nitrogen, or fluorine, pure dimethyl ether lacks this capability. The distinction matters because it explains why dimethyl ether behaves so differently from its structural twin, ethanol, despite having the exact same atoms.
Why the Structure Rules Out Hydrogen Bonding
Dimethyl ether has the formula CH₃OCH₃. The oxygen sits in the middle of the molecule with a carbon on each side, and all six hydrogen atoms are bonded to those carbons, not to the oxygen. For a molecule to form hydrogen bonds with other copies of itself, it needs at least one hydrogen directly attached to oxygen (or nitrogen or fluorine). Carbon-hydrogen bonds simply aren’t polar enough to participate in hydrogen bonding.
The molecule is bent at the oxygen, with a C-O-C bond angle of about 111.2 degrees. That bent shape makes dimethyl ether a polar molecule, meaning it has a positive end and a negative end. But polarity alone doesn’t create hydrogen bonds. It creates a weaker type of attraction called dipole-dipole interaction, which is the main intermolecular force holding liquid dimethyl ether together, along with the even weaker London dispersion forces that exist between all molecules.
The Ethanol Comparison Makes It Clear
The most convincing evidence comes from comparing dimethyl ether to ethanol. These two molecules are structural isomers, meaning they contain the exact same atoms (C₂H₆O) arranged differently. Ethanol has a hydrogen bonded directly to its oxygen (an O-H group), so it forms strong hydrogen bonds between molecules. Dimethyl ether does not.
The difference in physical properties is dramatic. Ethanol boils at 78°C, while dimethyl ether boils at -25°C, a gap of over 100 degrees. Ethanol melts at -117°C compared to -138°C for dimethyl ether. These numbers reflect how much harder it is to pull ethanol molecules apart from each other. Hydrogen bonds are significantly stronger than the dipole-dipole forces in dimethyl ether, so ethanol requires far more energy to vaporize.
The heat of vaporization tells the same story from a different angle. Dimethyl ether requires only about 21.5 kJ/mol of energy to transition from liquid to gas. Ethanol needs roughly 38.6 kJ/mol, nearly double. That extra energy goes toward breaking all those hydrogen bonds that dimethyl ether simply doesn’t have.
It Can Accept Hydrogen Bonds From Other Molecules
Here’s where it gets a little more nuanced. While dimethyl ether can’t donate a hydrogen bond (it has no O-H, N-H, or F-H bond to offer), it can accept one. The oxygen atom still has two lone pairs of electrons, which can interact with a hydrogen that is bonded to an electronegative atom on a different molecule. When you mix dimethyl ether with water, for example, water’s O-H hydrogens can form hydrogen bonds with dimethyl ether’s oxygen.
This is why dimethyl ether is reasonably soluble in water. The oxygen acts as a hydrogen bond acceptor, allowing water molecules to latch onto it. So the accurate statement is that dimethyl ether cannot hydrogen bond with itself, but it can participate in hydrogen bonding as an acceptor when paired with a hydrogen bond donor like water or an alcohol.
What Forces Hold Dimethyl Ether Together
In its pure liquid form, dimethyl ether is held together by two types of intermolecular forces. The stronger of the two is dipole-dipole attraction. Because the molecule is bent rather than symmetrical, the electron-rich oxygen creates a partial negative charge on one side while the carbon and hydrogen ends carry partial positive charges. Neighboring molecules orient themselves so that opposite charges face each other, creating an attractive pull.
The second force is London dispersion, the weak, temporary attraction that arises when electron clouds briefly shift and create fleeting charge imbalances. Every molecule experiences London dispersion forces, and for small molecules like dimethyl ether, they contribute a meaningful share of the total intermolecular attraction. Together, these two forces are enough to keep dimethyl ether liquid below -25°C, but they’re far weaker than the hydrogen bonding network found in ethanol or water, which is why dimethyl ether is a gas at room temperature.