Hydrogen Peroxide: Preparation, Structure and Properties

Hydrogen peroxide, $\text{H}_2\text{O}_2$, is a pale blue, syrupy liquid, completely miscible with water, and one of the most important industrial oxidizing/bleaching agents.

Laboratory preparation. In the laboratory, hydrogen peroxide is conveniently obtained by treating a metal peroxide with a dilute acid, for example by acidifying sodium peroxide or barium peroxide:

the insoluble barium sulfate being removed by filtration to leave a solution of the peroxide.

Industrial preparation. Two methods are used industrially. The older, electrolytic route oxidizes a cold, concentrated solution of ammonium bisulfate (or sulfuric acid) at the anode of an electrolytic cell to peroxodisulfate:

and the peroxodisulfate is then hydrolysed, under reduced pressure to avoid decomposing the product, to give hydrogen peroxide:

The modern, large-scale route is the auto-oxidation (anthraquinone) process: a solution of 2-ethylanthraquinol in a suitable organic solvent is agitated with air (or oxygen), which oxidizes the quinol to 2-ethylanthraquinone while liberating hydrogen peroxide; the anthraquinone so formed is then catalytically reduced (with $\text{H}_2$/Pd or Ni) back to the quinol, regenerating the starting material so the whole cycle can be run continuously, with hydrogen peroxide simply extracted from the reaction mixture at each pass.

Structure. In both the gas phase and the solid state, hydrogen peroxide has a non-planar, "open book" or "skew" structure: the two O-H bonds do not lie in the same plane as each other. Picture the two oxygen atoms joined by the O-O bond as the spine of a partly-opened book, with each O-H bond lying in one of the two "pages" -- the dihedral angle between the two O-O-H planes is about $90.2^\circ$ in the gas phase, opening out to about $111.5^\circ$ in the solid state, where hydrogen bonding between neighbouring molecules widens the angle.

Oxidizing and reducing behaviour. The oxygen atoms in $\text{H}_2\text{O}_2$ sit in an intermediate oxidation state of $-1$ (compared with $0$ in $\text{O}_2$ and $-2$ in $\text{H}_2\text{O}$ or a normal oxide), and it is exactly this intermediate state that lets hydrogen peroxide act as either an oxidizing agent (being itself reduced to $\text{H}_2\text{O}$, oxidation state falling to $-2$) or a reducing agent (being itself oxidized to $\text{O}_2$, oxidation state rising to $0$), depending on what it reacts with. As an oxidizing agent, for example, it oxidizes $\text{Fe}^{2+}$ to $\text{Fe}^{3+}$ in acidic solution: $2\text{Fe}^{2+} + 2\text{H}^{+} + \text{H}_2\text{O}_2 \rightarrow 2\text{Fe}^{3+} + 2\text{H}_2\text{O}$. As a reducing agent, it reduces acidified potassium permanganate, itself being oxidized to oxygen gas: $2\text{MnO}_4^{-} + 6\text{H}^{+} + 5\text{H}_2\text{O}_2 \rightarrow 2\text{Mn}^{2+} + 8\text{H}_2\text{O} + 5\text{O}_2\!\uparrow$. …