Hydrogen bonding is an important topic in chemistry, and it underpins the behavior of many of the substances we interact with on a day-to-day basis, especially water. Understanding hydrogen bonding and why it exists is an important step in understanding intermolecular bonding and chemistry more generally. Hydrogen bonding is ultimately caused by the difference in net electric charge in some parts of specific molecules. These charged sections attract other molecules with the same properties.

The property of electronegativity ultimately causes hydrogen bonding. When atoms are covalently bonded to each other, they share electrons. In a perfect example of covalent bonding, the electrons are shared equally, so the shared electrons are about halfway between one atom and the other. However, this is only the case when the atoms are equally effective at attracting electrons. The ability of atoms to attract the bonding electrons is known as electronegativity, so if electrons are shared between atoms with the same electronegativity, then the electrons are roughly halfway between them on average (because electrons move continuously).

If one atom is more electronegative than the other, the shared electrons are more closely drawn to that atom. However, electrons are charged, so if they’re more prone to congregate around one atom than the other, this affects the balance of charge of the molecule. Rather than being electrically neutral, the more electronegative atom gains a slight net negative charge. Conversely, the less electronegative atom ends up with a slight positive charge. This difference in charge produces a molecule with what’s called a permanent dipole moment, and these are often called polar molecules.

Polar molecules have two charged sections within their structure. In the same way as the positive end of a magnet attracts the negative end of another magnet, the opposite ends of two polar molecules can attract each other. This phenomenon is called hydrogen bonding because hydrogen is less electronegative than molecules it often bonds with such as oxygen, nitrogen or fluorine. When the hydrogen end of the molecule with a net positive charge comes close to the oxygen, nitrogen, fluorine or another electronegative end, the result is a molecule-molecule bond (an intermolecular bond), which is unlike most other forms of bonding you encounter in chemistry, and it is responsible for some of the unique properties of different substances.

Hydrogen bonds are about 10 times less strong than the covalent bonds that hold the individual molecules together. Covalent bonds are hard to break because doing so requires a lot of energy, but hydrogen bonds are weak enough to be broken relatively easily. In a liquid, there are lots of molecules jostling around, and this process leads to hydrogen bonds breaking and reforming when the energy is sufficient. Similarly, heating the substance breaks some hydrogen bonds for effectively the same reason.

Water (H₂O) is a good example of hydrogen bonding in action. The oxygen molecule is more electronegative than hydrogen, and both of the hydrogen atoms are on the same side of the molecule in a “v” formation. This gives the side of the water molecule with the hydrogen atoms a net positive charge and the oxygen side a net negative charge. The hydrogen atoms of one water molecule, therefore, bond to the oxygen side of other water molecules.

There are two hydrogen atoms available for hydrogen bonding in water, and each oxygen atom can “accept” hydrogen bonds from two other sources. This keeps the intermolecular bonding strong and explains why water has a higher boiling point than ammonia (where the nitrogen can only accept one hydrogen bond). Hydrogen bonding also explains why ice occupies more volume than the same mass of water: The hydrogen bonds become fixed in place and give the water a more regular structure than when it is a liquid.