Part of the reason why these other atoms aren't commonly recognized as engaging hydrogen bonding is that compared to the traditional atoms, their hydrogen bonding is quite weak. Usually this means atoms like fluorine, oxygen, and nitrogen, but this can also include atoms like carbon, chlorine, and sulfur. You also want these electronegative atoms to be small the that donor and acceptor can approach each other more closely adding to the strength of the hydrogen bond. Well for hydrogen bonding to occur you want the hydrogen atom bonded to an electronegative atom (acting as the hydrogen bond donor) interacting with another electronegative atom (acting as the hydrogen bond acceptor). So fundamentally it's the same interaction which occurs in dipole-dipole, but because of the atom's sizes this attraction is enhanced and is stronger than other conventional dipole-dipole interactions. Third, since these atoms all quite small, the hydrogen atom on one molecule can approach the small electronegative atom (again N, O, and F) on another molecule very closely. Second the electronegativity difference between hydrogen and small electronegative atoms is large enough to be considered polar covalent, but not so much that it becomes ionic. First, it's simply a commonplace element in many crucial molecules like water. The reason we focus on hydrogen specifically is a result of several reasons. So we'll start having ionic interactions instead of dipole-dipole interactions which includes hydrogen bonding. And so pairing them with the electronegative atoms commonly associated with hydrogen bonding (nitrogen, oxygen, and fluorine) will have a large enough electronegativity difference that their bonding is no longer considered polar covalent, but rather ionic. The sphere on the right is oriented the same way with the positive side attracted to the "F" side of the first sphere.Well atoms like cesium or francium have very small electronegativity values. Most of the sphere corresponds to cloud of "F". The left sphere has a bulge on its left side corresponding to "H" and a partially positive charge. In each HF molecule, the hydrogen nucleus is rather poorly shielded by a thin electron cloud (only two electrons), and much of that electron cloud has been distorted toward the highly electronegative fluorine atom.ĭiagram of two distorted spheres. Suppose that two HF molecules approach each other, as shown in the following figure. In order to see why this happens, let us consider the simplest second-row hydride-HF. Clearly these second-row hydrides must have particularly strong intermolecular forces. In the second period, however, the polar hydrides NH 3, H 2O, and HF all have boiling points more than 100☌ above that of the nonpolar compound CH 4. Similar behavior occurs among the hydrides of elements in the fourth and third periods. SbH 3, H 2Te and HI, all of which are polar, have somewhat higher boiling points, but all lie within a range of 50☌. SnH 4, which consists of nonpolar molecules, boils at the lowest temperature. Hydrides of elements in the fifth period behave as we might predict. Note the anomalously high boiling points of H2O, HF, and NH3 in the second period. \) The boiling points of the hydrides of the nonmetals plotted against the period in which they occur in the periodic table.
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