Hydrogen Peroxide: Structure and Preparation

Hydrogen peroxide adopts a non-planar, open-book (skew) shape: the two O-H bonds sit on either side of the O-O bond at a dihedral angle, twisted specifically to minimise repulsion between the lone pairs and bond pairs of electrons around each oxygen. The two oxygen atoms sit close enough together that repulsion between their lone pairs also makes the O-O bond itself comparatively weak and susceptible to cleavage (part of why H₂O₂ decomposes/reacts readily). In the liquid state, H₂O₂ is noticeably more viscous than water, because strong intermolecular hydrogen bonding combines with a restricted rotation around the O-O bond (the molecule cannot easily twist to relieve steric/electronic strain while also forming and breaking hydrogen bonds with neighbours), making the liquid flow less easily than water. The dihedral angle itself also differs measurably between the solid state (a smaller angle, close to 90 degrees) and the gas phase (a larger angle, close to 111 degrees) — a difference driven by intermolecular hydrogen bonding constraints present in the solid state that are absent for an isolated gas-phase molecule.

The classic laboratory preparation of hydrogen peroxide reacts hydrated barium peroxide, BaO₂.8H₂O, with dilute sulfuric acid: BaO₂.8H₂O + H₂SO₄ → BaSO₄ (an insoluble precipitate, which is filtered off) + H₂O₂ + 8H₂O. This works because the peroxide ion in BaO₂ is simply protonated/displaced by the acid while barium is removed from solution as insoluble barium sulfate, leaving H₂O₂ behind in the filtrate. Other peroxide- or oxide-type compounds listed as distractors (PbO₂, MnO₂, TiO₂) do not release H₂O₂ this way with dilute acid.