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Chemistry · Ch 6 — Chemical Bonding and Molecular Structure

Cause of Formation of Hydrogen Bond

6.9.1

Cause of Formation of Hydrogen Bond

The Cause of Hydrogen Bond Formation

The hydrogen bond is not a true chemical bond like a covalent or ionic bond. It is a special type of dipole-dipole attraction that arises from a very specific set of conditions. The entire phenomenon begins with the bond between hydrogen and a highly electronegative element.

When hydrogen is covalently bonded to a strongly electronegative atom (call it X, such as fluorine, oxygen, or nitrogen), the shared pair of electrons is pulled much closer to X than to hydrogen. Because hydrogen has only one electron to begin with, this displacement leaves the hydrogen nucleus (a single proton) largely exposed. The hydrogen atom therefore acquires a significant fractional positive charge, denoted as δ+\delta^+. Meanwhile, the electronegative atom X, having gained a greater share of the electron density, acquires a fractional negative charge, denoted as δ−\delta^-.

This separation of charge creates a polar covalent bond, Hδ+−Xδ−H^{\delta+} - X^{\delta-}. The molecule itself becomes a dipole. The key to hydrogen bonding is that the δ+\delta^+ hydrogen of one molecule is then strongly attracted to the δ−\delta^- electronegative atom (X) of a neighbouring molecule. This electrostatic attraction is the hydrogen bond.

The textbook represents this chain of interactions as:

Hδ+−Xδ−⋯Hδ+−Xδ−⋯Hδ+−Xδ−H^{\delta+} - X^{\delta-} \cdots H^{\delta+} - X^{\delta-} \cdots H^{\delta+} - X^{\delta-}

The dotted lines (⋯\cdots) represent the hydrogen bonds. They are much weaker than covalent bonds but are significantly stronger than typical dipole-dipole interactions.

Watch out

A common mistake is to think the hydrogen atom is "lost" or becomes a bare proton. It is not. The hydrogen atom remains covalently bonded to its own X atom. The hydrogen bond is an intermolecular attraction between the δ+\delta^+ hydrogen of one molecule and the δ−\delta^- atom of another.

Dependence on Physical State

The magnitude of hydrogen bonding is not constant. It depends heavily on how closely the molecules can approach each other, which is determined by the physical state of the substance.

  • Solid State: In a solid, molecules are packed in a fixed, ordered arrangement. The distances between the δ+\delta^+ hydrogen of one molecule and the δ−\delta^- atom of another are at a minimum. Consequently, the hydrogen bond strength is at its maximum in the solid state.
  • Liquid State: In a liquid, molecules have more kinetic energy and are in constant motion. The average distance between molecules is greater than in a solid, and the ordered arrangement is lost. The hydrogen bond strength is therefore intermediate in the liquid state.
  • Gaseous State: In a gas, molecules are far apart and move independently at high speeds. The electrostatic attraction between them is too weak to hold them together for any significant time. The hydrogen bond strength is at its minimum in the gaseous state, and for most practical purposes, it is negligible.
Important

This variation in hydrogen bond strength with physical state is the direct reason why many hydrogen-bonded compounds (like water) have unusually high melting and boiling points. A large amount of energy is required to overcome the strong hydrogen bonds in the solid or liquid state to transition to a less-ordered phase.

Influence on Structure and Properties

Because hydrogen bonds are strong enough to influence the arrangement of molecules, they have a profound effect on the properties of compounds. The textbook explicitly states that hydrogen bonds have a strong influence on the structure and properties of the compounds. This is not a list of properties to be memorised, but a direct consequence of the cause described above. …

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