Hydrogen Bonding — Definition

Definition

Imagine you have a hydrogen atom that's feeling a bit 'pulled' by a very strong, bossy neighbor. In chemistry, this 'bossy neighbor' is a highly electronegative atom like Fluorine (F), Oxygen (O), or Nitrogen (N).

When hydrogen is directly attached to one of these super-electronegative atoms, the electron pair in the covalent bond gets pulled much closer to the bossy neighbor. This leaves the hydrogen atom with a partial positive charge (we denote this as $\delta+$) and the electronegative atom with a partial negative charge ($\delta-$).

This is the first crucial step.

Now, imagine this partially positive hydrogen atom sees another highly electronegative atom (again, F, O, or N) nearby, either in the same molecule or in a different molecule. This second electronegative atom usually has at least one lone pair of electrons, which are negatively charged. Because opposites attract, the partially positive hydrogen atom gets attracted to the partially negative (or simply electron-rich) electronegative atom. This attractive force is what we call a 'hydrogen bond'.

It's not a true chemical bond like a covalent bond (where electrons are shared) or an ionic bond (where electrons are transferred). Instead, it's a strong intermolecular (between molecules) or intramolecular (within the same molecule) attractive force.

Think of it as a strong handshake between molecules or different parts of the same molecule. This 'handshake' is strong enough to significantly affect properties like boiling points (making them higher), solubility (making substances dissolve better in water), and even the shapes of complex biological molecules like DNA and proteins.

Without hydrogen bonding, life as we know it wouldn't exist, as water would be a gas at room temperature!