# Hydrogen peroxide preparation by electrolytic method

Hydrogen peroxide is largely manufactured by the electrolysis of 50% Sulphuric acid at $0^{\circ}&space;\textrm{}&space;C$ between platinum anode and lead cathode separated by a porous diaphragm.

50% Sulphuric acid solution is taken in a stoneware vessel. A platinum rod is dipped in the solution which acts as anode. A lead wire acts as the cathode. These two electrodes are separated of a porous diaphragm.

The vessel is placed in a trough of ice to maintain the temperature at $0^{\circ}&space;\textrm{}&space;C$.

On passing current the following reactions takes place.

$2\textrm{}&space;H_{2}&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}$ $\rightleftharpoons$  $2&space;\textrm{}&space;H^{+}$ + $2&space;\textrm{}&space;H&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}^{-}$ (Ionization)

Oxidizing properties :

$\textrm{}&space;H_{2}\textrm{}&space;O_{2}$ decomposes as $\textrm{}&space;H_{2}\textrm{}&space;O_{2}$ $\longrightarrow$ $\textrm{}&space;H_{2}&space;\textrm{}&space;O$ $+$ [O]. Because of this tendency, it acts as powerful oxidant.

Examples:

$\textrm{}&space;p\textrm{}&space;b&space;\textrm{}&space;S&space;+&space;4&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}&space;\longrightarrow&space;\textrm{}&space;p&space;\textrm{}&space;b&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;4&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O$

It oxidizes potassium iodide to iodine in presence of sulphuric acid.

$2&space;\textrm{}&space;K&space;\textrm{}&space;I&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;\longrightarrow&space;\textrm{}&space;K_{2}&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;\textrm{}&space;I_{2}&space;+&space;2&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O$

It oxidizes ferrous sulphate to ferric sulphate in presence of sulphuric acid.

$2&space;\textrm{}&space;F&space;\textrm{}&space;e&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}&space;\longrightarrow&space;\textrm{}&space;F&space;\textrm{}&space;e_{2}\left&space;(&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;\right&space;)_{3}&space;+&space;2&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O$

Reducing Properties :

$\textrm{}&space;H_{2}\textrm{}&space;O_{2}$ acts as a reducing agent because of the ease with which it accepts nascent oxygen from other oxidising agent.

$\mathrm{}&space;H_{2}&space;\mathrm{}&space;O_{2}&space;+&space;(\mathrm{}&space;O)&space;\longrightarrow&space;\mathrm{}&space;H_{2}&space;\mathrm{}&space;O&space;+&space;\mathrm{}&space;O_{2}$

It reduces moist silver oxide to metallic silver.

$\textrm{}&space;A&space;\textrm{}&space;g_{2}&space;\textrm{}&space;O&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}&space;\longrightarrow&space;2&space;\textrm{}&space;A&space;\textrm{}&space;g&space;+&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O&space;+&space;\textrm{}&space;O_{2}$

It reduces ozone to oxygen.

$\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}+&space;\textrm{}&space;O_{3}&space;\longrightarrow&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O&space;+2&space;\textrm{}&space;O_{2}$

It decolourises acidified potassium permanganate solution.

$2&space;\textrm{}&space;K&space;\textrm{}&space;M&space;\textrm{}&space;n&space;\textrm{}&space;O_{4}&space;+&space;3&space;\textrm{}&space;H_{2}&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;5&space;\textrm{}&space;H_{2}&space;\textrm{}&space;O_{2}&space;\longrightarrow&space;\textrm{}&space;K_{2}\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;2\textrm{}&space;M&space;\textrm{}&space;n&space;\textrm{}&space;S&space;\textrm{}&space;O_{4}&space;+&space;8&space;\textrm{}&space;H_{2}O&space;+&space;5&space;\textrm{}&space;O_{2}$.