How do I break the hydrogen bond

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The most common are hydrogen bonds between hydrogen and oxygen. Their existence is also noticeable in the course of the non-metal-hydrogen compounds of the 6th main group. Like hydrogen fluoride, water has an abnormally high boiling point. The other anomalies of the water (melting point, increase in volume when solidifying) can also be traced back to the hydrogen bonds.

In the ice crystal, each oxygen atom is tetrahedrally surrounded by four hydrogen atoms. It forms covalent bonds with two of them (97) and with two more distant hydrogen bonds (178). Each H atom is in turn coordinated by two O atoms, in turn covalently bound with one and with the other via a hydrogen bond. This energetically favorable arrangement, which enables a maximum number of hydrogen bridges, results in a bulky structure with a relatively poor use of space. This structure is responsible for the lower density of the ice compared to liquid water. When the ice melts, however, the hydrogen bonds do not break completely, but rather ordered areas, so-called clusters, are retained in the above sense. This order decreases with increasing temperature, i.e. the clusters become smaller, the space filling better and the density increases until it reaches its maximum at 4. As the temperature continues to rise, the simultaneous increase in thermal expansion gains the upper hand, and the density decreases, as with other fabrics.

Ammonia molecules also associate in the liquid state, but the hydrogen bonds are weaker and their number per molecule is lower than with water. The increase in boiling points from to corresponds primarily to the increase in dispersion interactions.

The methane molecule has no dipole moment because of its regular tetrahedral structure. The interactions are dispersion forces. Accordingly, the boiling points increase from to with the size of the molecule.