Hydrogen bonding is a distinctive type of dipole-dipole interaction that occurs between molecules. Crucially, it’s not a covalent bond to a hydrogen atom itself. Instead, it arises from the attractive force linking a hydrogen atom covalently bonded to a highly electronegative atom – such as nitrogen (N), oxygen (O), or fluorine (F) – and another electronegative atom. The strength of hydrogen bonds typically falls in the range of 4 to 50 kJ per mole.
In molecules containing N-H, O-H, or F-H bonds, the considerable electronegativity difference between the hydrogen atom and the N, O, or F atom leads to a highly polar covalent bond, effectively creating a bond dipole.
Below is a table highlighting the electronegativity values of these key elements:
| Element | Electronegativity Value |
|---|---|
| H | 2.1 |
| N | 3.0 |
| O | 3.5 |
| F | 4.1 |
This difference in electronegativity causes the hydrogen atom to acquire a substantial partial positive charge, while the N, O, or F atom gains a significant partial negative charge.
Consequently, a hydrogen atom in one molecule experiences an electrostatic attraction to the N, O, or F atom in another molecule. This attraction is what we define as a hydrogen bond.
| = O = N = H |
|---|
| Hydrogen bonding between two water (H2O) molecules. Note that the O atom in one molecule is attracted to a H atom in the second molecule. |
Physical Consequences Explained by Hydrogen Bonding
Consider the contrasting states of nitrosyl fluoride (ONF) and water (H2O) at 25°C: ONF exists as a gas, while water is a liquid. Let’s examine why this difference occurs.
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ONF and water share a similar molecular shape.
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ONF has a higher molecular weight (49 amu) compared to water (18 amu). This rules out London dispersion forces as the primary cause for the difference in their physical states.
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ONF and water possess similar dipole moments. This eliminates dipole-dipole forces as the sole reason for the observed difference.
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Crucially, ONF cannot form hydrogen bonds, whereas water can. This difference in hydrogen bonding capability explains their distinct physical states.
| = O = F = N = H |
|---|
| The structure of ONF showing no hydrogen bonding. |
| Microscopic view of ONF at 25°C, showing gaseous state. |
In summary, hydrogen bonding significantly impacts the physical properties of substances, particularly those containing N-H, O-H, or F-H bonds. The capability to form hydrogen bonds explains why water exists as a liquid at room temperature, despite its relatively low molecular weight.
