Water should boil at about −80 °C if it followed the trend of its group. It boils at 100 °C because of hydrogen bonding, the strongest force between molecules. This lesson also links bond length to bond energy.
A hydrogen bond forms between a hydrogen atom bonded to a very electronegative atom (N, O or F) in one molecule and a lone pair on N, O or F in a neighbouring molecule. The H carries a large δ+ charge and is tiny, so the lone pair can approach it closely. Hydrogen bonds are the strongest intermolecular forces, which explains the unusually high boiling points of H2O, HF and NH3 and the open, less dense structure of ice. Inside molecules, bond energy is the energy needed to break one mole of a covalent bond in the gas phase; shorter bonds are generally stronger, so C≡C > C=C > C–C.
Load water and then hydrogen fluoride in the intermolecular forces simulator and switch on the hydrogen bond view. Notice that each hydrogen bond runs in a straight line, O–H···O, with the lone pair pointing at the hydrogen.
Fig. 1 shows the boiling points of the hydrides rising down each group as the number of electrons rises. The period 2 hydrides of N, O and F break the trend.
| Property | What happens | Why |
|---|---|---|
| boiling point of H2O | 100 °C, not about −80 °C | extensive hydrogen bonding must be overcome |
| density of ice | ice floats on water | hydrogen bonds hold molecules in an open lattice, further apart than in the liquid |
| solubility | ethanol, ammonia dissolve in water | they form hydrogen bonds with water molecules |
| surface tension, viscosity | high for water and glycerol | strong attractions between molecules |
Four quick checks. Each correct answer earns XP and records this skill on your profile.
A hydrogen bond can form between H bonded to which elements and a lone pair on another molecule?
Which substance forms hydrogen bonds between its own molecules?
Why is ice less dense than liquid water?
Which bond is the strongest?
Inside a molecule, the covalent bonds have measurable lengths and strengths. Multiple bonds pull the atoms closer together and hold them more strongly.
| Bond | Length / nm | Mean bond energy / kJ mol−1 |
|---|---|---|
| C–C | 0.154 | 350 |
| C=C | 0.134 | 610 |
| C≡C | 0.120 | 840 |
| H–Cl / H–Br / H–I | 0.127 / 0.141 / 0.161 | 431 / 366 / 299 |
Bond energies for bonds such as C–H are mean values averaged over many compounds, because the exact value depends on the rest of the molecule. The falling H–X bond energies down Group 17 explain why HI decomposes on heating far more easily than HCl.
A hydrogen bond is an intermolecular force, not a covalent bond; the covalent O–H bond inside water is about 20 times stronger. When drawing hydrogen bonding you must show the lone pair and the partial charges, and the dashed line must go from the lone pair to the hydrogen, not to the other atom.
Unlocks once the four checks above are done. Worth more XP, written to AS Paper 1 and Paper 2 standard.
The boiling point of water (100 °C) is much higher than that of hydrogen sulfide (−60 °C). The main reason is that:
Water boils at 100 °C and hydrogen fluoride at 20 °C, although the H–F bond is more polar. Why?
Which diagram feature is essential when showing a hydrogen bond between two water molecules?
The H–X bond energies are H–Cl 431, H–Br 366 and H–I 299 kJ mol−1. Which explains this trend?
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