Molecular structure and hydrogen bonding

Oxygen's electron configuration is 1s² 2s² 2p⁴; it needs two electrons to complete its shell, so it shares them with two hydrogen atoms. The four electron pairs around oxygen are sp³ hybridised, but two of them are lone pairs, so the molecule is bent (V-shaped) with an H–O–H angle of 104.5°, not tetrahedral 109.5°.

• Oxygen (EN 3.5) attracts the shared electrons more strongly than hydrogen (EN 2.1), so oxygen carries a partial negative charge δ− and each hydrogen a partial positive charge δ+ — the molecule is a dipole.
• Each water molecule can form up to four hydrogen bonds with its neighbours: two through its hydrogen atoms and two through the lone pairs of oxygen.
• In the liquid the bonds break and reform about 10¹² times a second; in ice the network is fixed and open, which is why ice is less dense.
• Hydrogen bonding also explains the anomalously high boiling points of the hydrides of group 16: H₂O (100 °C) far above H₂S (−60 °C), H₂Se and H₂Te.