WBCHSE Class 11 Chemistry Notes For Hydrogen Peroxide – Structure, Formula and Properties

Hydrogen Peroxide Preparation

In the laboratory, hydrogen peroxide is prepared by the following methods.

1. From sodium peroxide Calculated quantities of sodium peroxide are added in small portions to ice-cold dilute sulphuric acid (20%)

⇒ \(\mathrm{Na}_2 \mathrm{O}_2+\mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \mathrm{Na}_2 \mathrm{SO}_4+\mathrm{H}_2 \mathrm{O}_2\)

On cooling, most of the sodium sulphate separates out as Na₂SO4- 10H₂O and is filtered. By this process, a 30 per cent solution of hydrogen peroxide containing a bit of sodium sulphate is obtained.

2. From barium peroxide The reaction of hydrated barium peroxide with dilute sulphuric add gives hydrogen peroxide and barium sulphate, which is insoluble is filtered off. The method was first used to prepare hydrogen peroxide in 1818 by Thenard.

⇒ \(3 \mathrm{BaO}_2+2 \mathrm{H}_3 \mathrm{PO}_4 \longrightarrow \mathrm{Ba}_3\left(\mathrm{PO}_4\right)_2+3 \mathrm{H}_2 \mathrm{O}_2\)
Barium peroxide Phosphoric Acid Barium Phosphate

By electrolysis, This method used earlier involves the electrolysis of an equimolar mixture of sulphuric acid and r ammonium sulphate using a platinum anode and a graphite cathode.

⇒ \(\mathrm{H}_2 \mathrm{SO}_4 \rightleftharpoons \mathrm{H}^{+}+\mathrm{HSO}_4^{-}\)

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⇒ \(\underset{{\text { anmonium } \\ \text { sulphiate }}}{\left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4}+\mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \underset{{\text { amrmonium hydrogen } \\ \text { sulphate }}}{2 \mathrm{NH}_4 \mathrm{HSO}_4}\)

⇒ \(\mathrm{NH}_4 \mathrm{HSO}_4 \longrightarrow \mathrm{NH}_4^{+}+\mathrm{HSO}_4^{-}\)

⇒ \(2 \mathrm{HSO}_4^{-} \longrightarrow \mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_8+2 \mathrm{e}^{-} \text {(at anode) }\)

⇒ \(\mathrm{NH}_4 \mathrm{HSO}_4 \longrightarrow \mathrm{NH}_4^{+}+\mathrm{HSO}_4^{-}\)

The solution, on hydrolysis, produces hydrogen peroxide.

⇒ \(\mathrm{H}_2 \mathrm{~S}_2 \mathrm{O}_8+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{{\text { perowomono- } \\ \text { sulphuric acid }}}{\mathrm{H}_2 \mathrm{SO}_5}+\mathrm{H}_2 \mathrm{SO}_4\)

⇒ \(\mathrm{H}_2 \mathrm{SO}_5+\mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{H}_2 \mathrm{SO}_4+\mathrm{H}_2 \mathrm{O}_2\)

By the oxidation of 2-ethyl anthraquinol Hydrogen peroxide is nowadays produced on a commercial scale by a cyclic process.

In this method, 2-ethyl anthraquinol is oxidised by air to the corresponding quinone and hydrogen peroxide at a moderate temperature.

The anthraquinone is reduced back to anthraquinone with hydrogen using platinum, palladium or Raney nickel as a catalyst. The cycle is repeated.

The hydrogen peroxide formed is extracted with water and the aqueous solution (containing 1-2% H₂O₂) is concentrated under reduced pressure. It is sold as a 30 per cent (by weight) solution which has a pH of about 4.0. It is stored in plastic or wax-coated glass vessels.

The coating is done to prevent the decomposition of hydrogen peroxide on account of the traces of alkali present in glass, which catalyse the decomposition. Small amounts of urea or sodium stannate are added as stabilisers, to prevent the breakdown of H₂O₂ into water and oxygen.

For further concentration, the solution is cooled to about 263 K when crystals enriched in H₂O₂ separate out. The process of fractional crystallisation is repeated to obtain 100 per cent H₂O₂.

Hydrogen Peroxide Physical properties

Hydrogen peroxide is a colourless, odourless, syrupy liquid. It has a bluish tinge and is soluble in water, ether and alcohol in all proportions.

It is a liquid in which molecules are highly associated with each other due to the presence of intramolecular hydrogen bonds.

In fact, it is more hydrogen-bonded than water and has a boiling point of 423.2 K and a melting point of 273.6 K. Also, it has a dipole moment of 2.1 D, a little more than that of water (1.84 D).

Hydrogen Peroxide Chemical properties

Hydrogen peroxide is unstable and its rate of decomposition depends on temperature and concentration.

It V behaves as an oxidising as well as a reducing agent in both acidic and alkaline media. The oxidation state of oxygen in hydrogen peroxide is -1, compared to zero in O₂ and -2 in H₂O.

Therefore, hydrogen peroxide can be oxidised to or reduced to H₂O. H₂O₂ is a powerful oxidising agent.

1. Oxidation in an acidic medium: It oxidises acidified solutions of ferrous salts, iodides, sulphides and ferrocyanides to ferric salts, iodine, sulphates and ferricyanides respectively.

⇒ \(2 \mathrm{FeSO}_4+\mathrm{H}_2 \mathrm{SO}_4+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(2 \mathrm{KI}+\mathrm{H}_2 \mathrm{SO}_4+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{K}_2 \mathrm{SO}_4+\mathrm{I}_2+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{PbS}+4 \mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{PbSO}_4+4 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\underset{{\text { potassium } \\ \text { ferrocyanide }}}{2 \mathrm{~K}_4\left[\mathrm{Fe}(\mathrm{CN})_6\right]+\mathrm{H}_2 \mathrm{SO}_4}+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \underset{{\text { potassium } \\ \text { ferricyanide }}}{2 \mathrm{~K}_3\left[\mathrm{Fe}(\mathrm{CN})_6\right]+\mathrm{K}_2 \mathrm{SO}_4}+2 \mathrm{H}_2 \mathrm{O}\)

Hydrogen peroxide oxidises sulphites, nitrites and arsenites to sulphates, nitrates and arsenates respectively.

⇒ \(\underset{\text { sodium sulphite }}{\mathrm{Na}_2 \mathrm{SO}_3}+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \underset{\text { sodium sulphate }}{\mathrm{Na}_2 \mathrm{SO}_4}+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{KNO}_2+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{KNO}_3+\mathrm{H}_2 \mathrm{O}\)

⇒ \(\mathrm{Na}_3 \mathrm{AsO}_3+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{Na}_3 \mathrm{AsO}_4+\mathrm{H}_2 \mathrm{O}\)

2. Oxidation in an alkaline medium: Hydrogen peroxide oxidises chromium salts to chromates and manganese salts to manganese dioxide.

⇒ \(\mathrm{Cr}_2\left(\mathrm{SO}_4\right)_3+3 \mathrm{H}_2 \mathrm{O}_2+10 \mathrm{NaOH} \longrightarrow 2 \mathrm{Na}_2 \mathrm{CrO}_4+3 \mathrm{Na}_2 \mathrm{SO}_4+8 \mathrm{H}_2 \mathrm{O}\)

⇒ \(\underset{{\text { manganese } \\ \text { sulphate }}}{\mathrm{MnSO}_4}+\mathrm{H}_2 \mathrm{O}_2+2 \mathrm{NaOH} \longrightarrow \mathrm{Na}_2 \mathrm{SO}_4+\underset{{\text { manganese } \\ \text { dioxide }}}{\mathrm{MnO}_2}+2 \mathrm{H}_2 \mathrm{O}\)

3. Reducing action in an acidic medium: The following reactions depict the reducing action of H₂O₂ in an acidic medium.

⇒ \(\mathrm{H}_2 \mathrm{O}_2+\mathrm{Cl}_2 \longrightarrow 2 \mathrm{HCl}+\mathrm{O}_2\)

4. It reduces acidified potassium permanganate (pink) to manganese sulphate (colourless)

⇒ \(2 \mathrm{KMnO}_4+3 \mathrm{H}_2 \mathrm{SO}_4+5 \mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{K}_2 \mathrm{SO}_4+2 \mathrm{MnSO}_4+8 \mathrm{H}_2 \mathrm{O}+5 \mathrm{O}_2\)

It reduces the acidified potassium dichromate solution (yellow) to chromic sulphate (green).

⇒ \(\mathrm{K}_2 \mathrm{Cr}_2 \mathrm{O}_7+4 \mathrm{H}_2 \mathrm{SO}_4+3 \mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{K}_2 \mathrm{SO}_4+\mathrm{Cr}_2\left(\mathrm{SO}_4\right)_3+7 \mathrm{H}_2 \mathrm{O}+3 \mathrm{O}_2\)

It reduces silver oxide to silver with the liberation of oxygen gas, and lead dioxide to lead monoxide.

⇒ \(\mathrm{Ag}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow 2 \mathrm{Ag}+\mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

⇒ \(\mathrm{PbO}_2+\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{PbO}+\mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

5. Manganese dioxide is reduced to manganous sulphate.

⇒ \(\underset{{\text { Manganese } \\ \text { dioxide }}}{\mathrm{MnO}_2}+\mathrm{H}_2 \mathrm{O}_2+\mathrm{H}_2 \mathrm{SO}_4 \longrightarrow \underset{{\text { Manganous } \\ \text { sulphite (pink) }}}{\mathrm{MnSO}_4}+2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

4. Reducing action in an alkaline medium reduces an alkaline solution of potassium ferrocyanide to potassium ferrocyanide.

⇒ \(\underset{\text { ferricyanide }}{2 \mathrm{Fe}(\mathrm{CN})_6{ }^{3-}}+\mathrm{H}_2 \mathrm{O}_2+2 \mathrm{OH}^{-} \longrightarrow \underset{\text { ferrocyanide }}{2 \mathrm{Fe}(\mathrm{CN})_6{ }^{4-}}+2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2\)

It reduces ferric salts to ferrous salts.

⇒ \(\underset{\text { ferric ion }}{2 \mathrm{Fe}^{3+}}+\mathrm{H}_2 \mathrm{O}_2+2 \mathrm{OH}^{-} \longrightarrow \underset{\text { ferrous ion }}{2 \mathrm{Fe}^{2+}}+2 \mathrm{H}_2 \mathrm{O}+\underset{\text { dioxygen }}{\mathrm{O}_2}\)

Reaction with ozone Hydrogen peroxide reduces ozone to oxygen.

⇒ \(\mathrm{H}_2 \mathrm{O}_2+\mathrm{O}_3 \longrightarrow 2 \mathrm{O}_2+\mathrm{H}_2 \mathrm{O}\)

Bleaching action Hydrogen peroxide is used to bleach delicate articles like ivory, feathers, silk, etc. The bleaching action is due to the following reaction.

⇒ \(\mathrm{H}_2 \mathrm{O}_2 \longrightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{O}\)

The nascent oxygen obtained on the decomposition of hydrogen peroxide oxidises the colouring matter of the article to be bleached. When this happens the article loses colour, i.e., it gets bleached.

Hydrogen Peroxide Uses

1. Hydrogen peroxide is used as an antiseptic and germicide for washing wounds. It is also employed in mouthwashes.
2.  H₂O₂ is used as an oxidant for rocket fuel and as a propellant for submarines and torpedoes.
3. It is employed in the manufacture of many inorganic and organic compounds like sodium peroxoborate [Na2B2(02)2(0H)4 -6 H₂O₂ and epoxides. Sodium peroxoborate is used in washing powder. Organic peroxides are used to initiate addition polymerisation reactions (FVC, polyurethane and epoxy resins).
4. It is used in organic synthesis for the conversion of an aldehydic group (—CHO) into a hydroxyl group (—OH). The oxidation is carried out in an alkaline medium.
5. It is employed as an oxidising agent in the laboratory.
6. In the textile industry, it is used as what is known as an antichlor. This means that it helps remove excess chlorine after bleaching.

Structure Of Hydrogen Peroxide

The hydrogen peroxide molecule is nonplanar. The two oxygen atoms are linked to each other by a covalent bond, known as a peroxide bond in this case.

Also, each oxygen atom is linked to a hydrogen atom by covalent bonds. The two O—H bonds are in different planes.

The angle between the two planes is 111.5° in the gaseous phase. Because of hydrogen bonding in the crystalline state, however, this angle becomes 90.2° in the crystalline state.

The Structure Diagram Of Hydrogen Peroxide: