WBBSE Solutions For Class 8 School Science Chapter 3 Some Common Gases Laboratory Experiments On Preparation of Oxygen And Hydrogen

Preparation of Oxygen

1. Laboratory Preparation Of Oxygen From Potassium Chlorate

Oxygen is usually prepared in the laboratory by heating carefully a mixture of potassium chlorate (KCIO₃) and manganese dioxide (MnO₂).

Four parts of solid potassium chlorate are intimately mixed with one part of solid manganese dioxide and taken in a hard glass test tube.

The test tube is fitted in such a way that it is tilted downwards. A delivery tube is fixed at the mouth of the test tube with the help of the bore of the cork.

The other end of the delivery tube is Potassium introduced into the gas jar filled with water. chlorate + The test tube is then heated strongly by a dioxide Bunsen burner.

Potassium chlorate melts and decomposes, evolving oxygen. The gas is collected in the gas jar by the downward displacement of water.

⇒ \(2 \mathrm{KClO}_3+\left[\mathrm{MnO}_2\right] \stackrel{\text { heat }(\Delta)}{\longrightarrow} 2 \mathrm{KCl}+3 \mathrm{O}_2+\left[\mathrm{MnO}_2\right]\)

The gas is collected by the downward displacement of water, because,

1. The solubility of oxygen in water is low
2. Oxygen is almost as heavy as air, so it cannot be collected by the downward displacement of air

In this reaction, MnO₂ acts as a catalyst. If KCIO₃ is heated alone, oxygen is produced at a temperature higher than 610°C.

In presence of a little amount of MnO₂, KCIO₃ decomposes at about 250°C to produce oxygen.

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Actually, when KCIO₃ is heated alone, it melts at 357°C and rapidly gives off oxygen at 380°C. But the mass becomes pasty as the reaction proceeds due to the formation of potassium perchlorate (KCIO₄ )whose melting point is 610°C.

⇒ \(4 \mathrm{KClO}_3 \rightarrow 3 \mathrm{KClO}_4+\mathrm{KCl}\)

When heated above 610°C, it decomposes to produce oxygen and a residue of KCI is left.

⇒ \(\mathrm{KClO}_4 \rightarrow \mathrm{KCl}+2 \mathrm{O}_2\)

In presence of a little amount of MnO₂, KCIO₃ smoothly decomposes at about 250°C to produce oxygen without the formation of KCIO₄  in the intermediate stage.

 

WBBSE Class 8 Preparation of Oxygen Notes

Thus MnO₂ accelerates the reaction and acts as a true catalyst. Moreover, both KCIO₃ and KCIO₄ are explosive substances. Any probable danger of explosion in the act of heating them to high temperatures is avoided by performing the reaction at a lower temperature with the help of a catalyst.

Precaution

1. KClO₃ and MnO₂ should be mixed intimately
2. MnO₂ must not be contaminated with charcoal or antimony sulphide
3. The hard glass test tube must be tilted downwards
4. Heating should be done slowly and should continue from the front to the back side of the test tube.

2. Preparation Of Oxygen From Hydrogen Peroxide At Room Temperature

Oxygen is obtained when solid manganese dioxide is added to a dilute aqueous solution of hydrogen peroxide.

⇒ \(2 \mathrm{H}_2 \mathrm{O}_2+\left[\mathrm{MnO}_2\right] \rightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2+\left[\mathrm{MnO}_2\right]\)

Materials and Apparatus Required: A round bottom flask, a long-necked funnel, some solid manganese dioxide (MnO₂), a dilute solution of hydrogen peroxide, a gas jar, and a jute stick.

Experiment: The round bottom flask is closed with a cork, fitted with a long-necked flask and a delivery tube, keeping some manganese dioxide solids inside.

Now, a dilute aqueous solution of hydrogen peroxide is poured in a round bottom flask through the funnel. The other end of the delivery tube is introduced into an inverted gas jar filled with water.

WBBSE Class 8 Preparation of Hydrogen Notes

Observation: Brisk effervescence due to the evolution of oxygen is observed. No heating is required to prepare oxygen by this method,

If a glowing jute stick is held on the mouth of the gas the rate of decomposition of H₂O If MnO₂ is not jar, it burns with a flame,

Confirming that the gas added to an aqueous solution of H₂O₂, oxygen is not produced in this reaction is oxygen. Here, MnO₂ acts as a catalyst and enhances.

3. Preparation of Oxygen from Sodium Peroxide at room temperature

Oxygen is produced easily when water is added to solid sodium peroxide. No heating is required for this process.

⇒ \(2 \mathrm{Na}_2 \mathrm{O}_2+2 \mathrm{H}_2 \mathrm{O} \rightarrow 4 \mathrm{NaOH}+\mathrm{O}_2\)

Materials and Apparatus Required: Solid sodium peroxide (Na₂O₂), distilled water, a conical flask, a cork with two holes in it, a dropping funnel, a bent delivery tube, and a gas jar.

Experiment: Solid Na₂O₂ is taken in the conical flask and the mouth of the conical flask is fitted with a cork.

Through one of the holes, a dropping funnel is attached and through another hole, one end of the bent delivery tube is inserted.

The other end of the bent delivery tube is introduced into the gas jar filled with water. Now water is added

Observation: Oxygen gas is evolved. The gas dropwise to solid Na₂O₂ through the dropping is collected in the gas jar by downward funnel displacement of water.

4. Preparation of Oxygen by Electrolysis of Water

Electrolysis of water acidified with dilute sulphuric acid can produce hydrogen at the cathode and oxygen gas at the anode.

A platinum electrode is used as an anode and cathode in a rectangular tank.

High voltage is passed through the tank to carry out the electrolysis. Here, hydrogen gas is obtained as a by-product.

Key Terms Related to Gas Preparation in Chemistry

Manufacture of Oxygen by Fractional Distillation of air Industrially oxygen is produced in bulk quantity by a process known as a fractional distillation of liquid air.

Air is composed of nitrogen and oxygen, in which oxygen forms about 21% by volume. The two gases can be separated from one another by liquefaction of air followed by fractional distillation.

Removal of water vapour, CO₂ & dust particles: Air is first freed from water vapour and carbon dioxide by passing them over fused calcium chloride and slaked lime, respectively.

Dust particles are removed from the air by passing them through an electric precipitator.

Liquefaction of air: Air is then liquefied by applying high pressure, followed by its sudden expansion into a region of very low pressure.

As a result, the temperature of the air falls and it is liquefied. Liquid air is a mixture of liquid oxygen (boiling point: – 183°C) and liquid nitrogen (boiling point: -196°C).

Fractional distillation of liquid air: Liquid air is then fractionally distilled. When the liquid air is evaporated, nitrogen (having a lower boiling point) will evolve first and oxygen.

(having a higher boiling point) will remain behind. The evolved gases are allowed to pass through a tall fractionation column.

Nearly pure nitrogen leaves through the top of the column and oxygen in the gas condenses and is collected at the bottom of the column as liquid oxygen.

When this liquid oxygen is evaporated, pure oxygen gas is obtained.

Physical Properties of Oxygen

1. Oxygen is a colourless, odourless and tasteless gas.
2. It is slightly heavier than air. The density of oxygen at normal temperature and pressure is 1.428 grams per litre.
3. It condenses to a pale blue liquid, which freezes to a blue solid if cooled in liquid hydrogen.
4. The freezing point of liquid oxygen is – 218°C and the boiling point of liquid oxygen is – 183°C.
5. Oxygen is slightly soluble in water. The solubility of oxygen at 0°C and 1 atmospheric pressure is 1438 mg/lit. The ‘ dissolved oxygen sustains the life of aquatic plants and animals.
6. The respiration of aquatic animals is dependent on the dissolved oxygen in the water. Since oxygen is more soluble in water than nitrogen, water is richer in oxygen than ordinary air.

Oxygen has three naturally occurring isotopes, \({ }_8^{16} \mathrm{O},{ }_8^{17} \mathrm{O} \text { and }{ }_8^{18} \mathrm{O}\) But the natural abundance of the last two is very low.

Isotope	168 O	178O	188O
Natural abundance	99.763%	0.037%	0.02%

Oxygen exhibits allotropy. Its allotropic modification is ozone (03).

Chemical Properties Of Oxygen

1. The oxygen molecule is diatomic. At high temperatures, oxygen molecule dissociates to produce atomic oxygen.

This is an endothermic reaction. Atomic oxygen is a powerful oxidizing agent.

⇒ \(\mathrm{O}_2 \rightarrow \mathrm{O}+\mathrm{O} \text { – heat }(\mathrm{Q})\)

2. Oxygen is non-combustible i.e. it itself does not burn. But it is a supporter of combustion. It rekindles a glowing splint. A glowing splinter is burst into flames when introduced in a gas jar filled with oxygen.

3. Oxygen is chemically very reactive and forms compounds with practically all elements except inert gases (such as helium, neon, argon etc.).

It combines with most of the elements except halogens (i.e. fluorine, chlorine, bromine, iodine, etc.) and a few noble metals (such as gold, platinum, etc.).

The reactivity of oxygen increases at high temperatures and in the presence of a suitable catalyst.

The compound produced in this reaction is called oxide. Actually, the oxide is a compound of two elements, one of which is oxygen.

1. Reaction Of Oxygen With Non-Metals

Oxygen reacts with non-metals such as carbon, sulphur, phosphorous, etc. to produce oxides. Generally, most oxides of non-metals are acidic.

Their aqueous solution produces acid. A few examples are given below.

1. When a piece of glowing charcoal, taken in a deflagrating spoon is introduced in a jar of oxygen, charcoal burns more brightly throwing sparks.

The product of this reaction is carbon dioxide.

⇒ \(\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2\)

When a moist blue litmus paper is held near the mouth of the gas jar, it turns red. If some clear lime water {Ca(OH)2)is poured into the gas jar and is shaken well, then the clear lime water turns milky due to the formation of insoluble calcium carbonate (CaCO₃) which remains suspended in solution.

⇒ \(\mathrm{Ca}(\mathrm{OH})_2+\mathrm{CO}_2 \rightarrow \mathrm{CaCO}_3+\mathrm{H}_2 \mathrm{O}\)

CO₂ reacts with water and forms carbonic acid (H₂CO₃), which is a weak acid. When a blue litmus paper is immersed in this solution, it turns red.

Examples of Chemical Reactions for Gas Preparation

A small quantity of sulphur Is taken in a deflagrating spoon and heated in a flame. When it melts and begins to burn feebly, it is then introduced in a jar filled with oxygen.

It is observed that the burning takes place brilliantly producing a blue flame and sulphur dioxide (SO₂) is produced. When a moist blue litmus paper is held near the mouth of the gas jar, it turns red.

⇒ \(\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{CO}_3\)

When SO₂ dissolves in water, it forms sulphurous acid (H₂SO₃). When a blue litmus paper is immersed in this solution, it turns red.

⇒ \(\mathrm{S}+\mathrm{O}_2 \rightarrow \mathrm{SO}_2\)

3. A piece of white phosphorous when introduced in a jar of oxygen, burns brightly with white flames and forms white fumes of phosphorous pentoxide (P2O5) which solidifies on cooling.

⇒ \(4 \mathrm{P}+5 \mathrm{O}_2 \rightarrow 2 \mathrm{P}_2 \mathrm{O}_5\)

It forms phosphoric acid when dissolved in water. The solution turns blue litmus paper red.

⇒ \(\mathrm{P}_2 \mathrm{O}_5+3 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{H}_3 \mathrm{PO}_4\)

2. Reaction of Oxygen with Metals

Some metals burn in oxygen on heating producing metal oxides. Metal oxides are mostly basic oxides.

Some basic oxides form hydroxides when they react with water. Hydroxides of some metals such as magnesium, calcium, etc. are soluble in water and turn red litmus paper blue. Some examples are given below.

1. When a piece of hot, dry sodium is introduced in a jar of oxygen, the metal burns spontaneously producing a golden yellow flame and forming sodium oxide (Na₂O).

When dissolved in water, it forms sodium hydroxide (NaOH) which turns red litmus paper blue.

⇒ \(4 \mathrm{Na}+\mathrm{O}_2 \rightarrow 2 \mathrm{Na}_2 \mathrm{O}\)

⇒ \(\mathrm{Na}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{NaOH}\)

When a burning magnesium ribbon is introduced in a gas jar filled with oxygen, it burns brightly producing blinding white light.

The white powdery substance left after burning is magnesium oxide (MgO).

⇒ \(2 \mathrm{Mg}+\mathrm{O}_2 \rightarrow 2 \mathrm{MgO}\)

When MgO reacts with water, it forms magnesium hydroxide which turns red litmus paper blue.

⇒ \(\mathrm{MgO}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Mg}(\mathrm{OH})_2\)

In a similar way, potassium, calcium, etc. react with oxygen to form basic oxides whose aqueous solutions are basic in nature.

⇒ \(2 \mathrm{Ca}+\mathrm{O}_2 \rightarrow 2 \mathrm{CaO} ; \quad \mathrm{CaO}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_2\)

⇒ \(4 \mathrm{~K}+\mathrm{O}_2 \rightarrow 2 \mathrm{~K}_2 \mathrm{O} ; \quad \mathrm{K}_2 \mathrm{O}+\mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{KOH}\)

3. Reaction of Oxygen with Metals and Formation of Amphoteric Oxides

Some metals such as Al, Zn, Pb, etc. react with oxygen to produce amphoteric oxide.

They are called amphoteric oxides because they have the properties of acidic and basic oxides. They undergo a neutralization reaction with both acids and bases.

They act as weakly basic oxides towards a strong acid and as weakly acidic oxides towards a strong base.

1. For example, the aluminium metal reacts with oxygen to form aluminium oxide (Al₂O₃).

⇒ \(4 \mathrm{Al}+3 \mathrm{O}_2 \rightarrow 2 \mathrm{Al}_2 \mathrm{O}_3\)

It reacts with an acid to form aluminium chloride (salt) and water.

⇒ \(\mathrm{Al}_2 \mathrm{O}_3+6 \mathrm{HCl} \rightarrow 2 \mathrm{AlCl}_3+3 \mathrm{H}_2 \mathrm{O}\)

It reacts with a base (such as sodium hydroxide) to produce sodium aluminate and water.

⇒ \(\mathrm{Al}_2 \mathrm{O}_3+2 \mathrm{NaOH} \rightarrow 2 \mathrm{NaAlO}_2+\mathrm{H}_2 \mathrm{O}\)

Practice Problems on Gas Preparation Experiments

2. Zinc oxide (ZnO) is an amphoteric oxide. Zinc metal reacts with oxygen to form zinc oxide (ZnO).

⇒ \(2 \mathrm{Zn}+\mathrm{O}_2 \rightarrow 2 \mathrm{ZnO}\)

It reacts with hydrochloric acid to form zinc chloride (salt) and water.

⇒ \(\mathrm{ZnO}+2 \mathrm{HCl} \rightarrow \mathrm{ZnCl}_2+\mathrm{H}_2 \mathrm{O}\)

It reacts with sodium hydroxide to produce sodium zincate and water.

⇒ \(2 \mathrm{ZnO}+4 \mathrm{NaOH} \rightarrow 2 \mathrm{Na}_2 \mathrm{ZnO}_2+2 \mathrm{H}_2 \mathrm{O}\)

3. As₂O₃ is a non-metallic amphoteric oxide. Arsenic reacts with oxygen to form arsenic oxide (As203).

⇒ \(4 \mathrm{As}+3 \mathrm{O}_2 \rightarrow 2 \mathrm{As}_2 \mathrm{O}_3\)

Reaction with acid: \(\mathrm{As}_2 \mathrm{O}_3+6 \mathrm{HCl} \rightarrow 2 \mathrm{AsCl}_3+3 \mathrm{H}_2 \mathrm{O}\)

Reaction with base: \(\mathrm{As}_2 \mathrm{O}_3+6 \mathrm{NaOH} \rightarrow 2 \mathrm{Na}_3 \mathrm{AsO}_3+3 \mathrm{H}_2 \mathrm{O}\)

4. Stannic oxide (SnO₂) is an amphoteric oxide with predominantly acidic properties, it is prepared by burning tin at white heat in the air.

It is insoluble in water and in all acids except concentrated H₂SO₄. It dissolves in concentrated H₂SO₄ to produce stannic sulphate (which is unstable). On fusion with sodium hydroxide, it forms sodium stannate, which is soluble in water.

⇒ \(\mathrm{Sn}+\mathrm{O}_2 \rightarrow \mathrm{SnO}_2\)

Reaction with acid: \(\mathrm{SnO}_2+2 \mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{Sn}\left(\mathrm{SO}_4\right)_2+2 \mathrm{H}_2 \mathrm{O}\)

Reaction with base: \(\mathrm{SnO}_2+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{SnO}_3+\mathrm{H}_2 \mathrm{O}\)

5. Lead monoxide is obtained by heating lead in the air. It reacts with HNO₃ forming lead nitrate. It dissolves in a hot sodium hydroxide solution forming sodium plumbite.

⇒ \(2 \mathrm{~Pb}+\mathrm{O}_2 \rightarrow 2 \mathrm{PbO}\)

Reaction with acid: \(\mathrm{PbO}+2 \mathrm{HNO}_3 \rightarrow \mathrm{Pb}\left(\mathrm{NO}_3\right)_2+\mathrm{H}_2 \mathrm{O}\)

Reaction with base: \(\mathrm{PbO}+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{PbO}_2+\mathrm{H}_2 \mathrm{O}\)

6. Formation of Peroxides

Oxygen reacts with excess sodium to form sodium peroxide.

⇒ \(4 \mathrm{Na}+\mathrm{O}_2 \rightarrow 2 \mathrm{Na}_2 \mathrm{O}_2\)

This contains peroxy linkage (- O – O -). When Na₂0₂ is treated with cold, dilute acids, such as HCI, it produces hydrogen peroxide.

⇒ \(\mathrm{Na}_2 \mathrm{O}_2+2 \mathrm{HCl} \rightarrow 2 \mathrm{NaCl}+\mathrm{H}_2 \mathrm{O}_2\)

7. Formation of Neutral Oxides

Some non-metallic oxides are neither acidic nor basic in character. These are called neutral oxides. Examples are carbon monoxide (CO), nitric oxide (NO), nitrous oxide (N20), etc.

4. Oxidation: Oxygen readily oxidizes ferrous, stannous and cuprous salts.

For example, if some solid ferrous sulphate is dissolved in dilute sulphuric acid, then after some time it is observed that the clear solution turns yellow.

This is due to the formation of ferric sulphate because of the oxidation of ferrous sulphate to ferric sulphate. [Note that “-ous” salt is oxidized to “-ic” salt].

⇒ \(4 \mathrm{FeSO}_4+2 \mathrm{H}_2 \mathrm{SO}_4+\mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2\left(\mathrm{SO}_4\right)_3+2 \mathrm{H}_2 \mathrm{O}\)

Similarly, on oxidation, stannous chloride is oxidized to stannic chloride and cuprous chloride is oxidized to cupric chloride.

⇒ \(2 \mathrm{SnCl}_2+4 \mathrm{HCl}+\mathrm{O}_2 \rightarrow 2 \mathrm{SnCl}_4+2 \mathrm{H}_2 \mathrm{O}\)

⇒ \(4 \mathrm{CuCl}+4 \mathrm{HCl}+\mathrm{O}_2 \rightarrow 4 \mathrm{CuCl}_2+2 \mathrm{H}_2 \mathrm{O}\)

Colourless nitric oxide is oxidized to brown-coloured nitrogen dioxide.

⇒ \(2 \mathrm{NO}+\mathrm{O}_2 \rightarrow 2 \mathrm{NO}_2\)

No can be prepared by heating copper turnings in 1:1 nitric acid.

⇒ \(3 \mathrm{Cu}+8 \mathrm{HNO}_3 \rightarrow 3 \mathrm{Cu}\left(\mathrm{NO}_3\right)_2+2 \mathrm{NO}+4 \mathrm{H}_2 \mathrm{O}\)

If this gas is exposed to air, then it is oxidized by the oxygen present in the air and brown nitrogen dioxide is produced. ‘

⇒ \(2 \mathrm{NO}+\mathrm{O}_2 \rightarrow 2 \mathrm{NO}_2\)

5. Oxidation of Hydrocarbons:

Hydrocarbons are compounds containing carbon and hydrogen only. The general formula of hydrocarbon is C₂H₂ (where x and y are integers).

When a hydrocarbon is burnt in oxygen, oxidation occurs. As a result, CO₂ and H₂O are produced along with heat energy. The general equation for the oxidation of hydrocarbon can be represented as follows:

⇒ \(\mathrm{C}_x \mathrm{H}_y+(x+y / 4) \mathrm{O}_2 \rightarrow x \mathrm{CO}_2+y / 2 \mathrm{H}_2 \mathrm{O}+\text { heat energy }\)

6. Respiration:

Respiration is a process by which complex foodstuff, such as glucose is oxidized in living organisms and heat energy is liberated.

Except for some lower animals and plants, most animals and plants consume the oxygen present in the air for the oxidation of complex foodstuffs.

We inhale air and use the oxygen present in it for respiration. As a result, carbonaceous foodstuff is converted into CO₂.

For example, during respiration in our body, glucose is oxidized to CO₂ and H₂O.

⇒ \(\mathrm{C}_6 \mathrm{H}_{12} \mathrm{O}_6+6 \mathrm{O}_2 \rightarrow 6 \mathrm{CO}_2+6 \mathrm{H}_2 \mathrm{O}+\text { heat energy }\)

7. Absorber of Oxygen :

1. Alkaline potassium pyrogallate solution absorbs oxygen and turns dark brown.

2. At normal or low pressure, certain metals like Au, Ag, Pt and Pd adsorb oxygen. This means oxygen molecules are weakly bound to the surface of these metals. On heating the metal, adsorbed oxygen is released (i.e. desorbed).

3. Ammonium cuprous chloride solution absorbs oxygen gas quickly and turns blue.

Tests Of Oxygen

1. The following tests can be performed to detect the evolution of oxygen gas.
2. The gas rekindles a glowing splinter.
3. It is neutral to litmus because oxygen is neither acidic nor basic in nature.
4. Colourless nitric oxide (NO) changes to reddish brown nitrogen dioxide (NO₂) gas.
5. It turns the alkaline Potassium pyrogallate solution to dark brown.

Uses of Oxygen

1. Oxygen is indispensable for respiration. Respiration is the process by which oxygen present in the inhaled air is utilized to oxidize the complex foodstuff in living beings.

2. Due to this, heat energy is released which is essential for performing various types of physiological processes.
(H) Oxidation of fuels (such as hydrocarbons) produces CO₂ and H₂O along with a significant amount of heat energy.

3. This energy is available for performing various kinds of mechanical work.

4. Oxygen is frequently used for medical purposes. For example, if patients suffer from asthma, pneumonia, etc. then they are supplied oxygen artificially (via gas cylinders).

5. For Patients who are rendered unconscious due to inhalation of poisonous gas, inhalation of carbogen (which is a mixture of 95% oxygen and 5% CO₂) is prescribed.

6. A mixture of oxygen and nitrous oxide is used for anaesthesia.

Understanding Laboratory Experiments for Oxygen and Hydrogen

7. Oxygen is a very important element in the chemical industry. It is used to remove impurities from crude iron and pure steel is produced. Impurities present in crude iron enhance rusting of iron.

During the preparation of H₂SO₄, oxygen is utilized. SO₂ is oxidized to SO₃ by a reaction with oxygen which is then converted to H₂S04 in subsequent steps.

H₂SO₄ is an essential component in car batteries, and storage cells and is used for making paints and fertilizers and for refining petroleum and metals like zinc and copper.

Oxygen is used during the industrial production of HNO₃ by the Ostwald process. HNO₃ is essential for producing fertilizers like ammonium nitrate (NH₄NO₃) and for preparing explosives.

Step-by-Step Guide to Preparing Oxygen in the Lab

In the Ostwald process for the preparation of nitric acid, initially, NH₃ and oxygen react at 900°C over a platinum wire gauge to form nitric oxide.

In the next stage, this NO reacts with oxygen and is oxidized to NO₂. NO₂ is dissolved in water and reacts with oxygen present in the air to form HNO₃

⇒ \(4 \mathrm{NH}_3+5 \mathrm{O}_2 \rightarrow 4 \mathrm{NO}+6 \mathrm{H}_2 \mathrm{O}\)

⇒ \(2 \mathrm{NO}+\mathrm{O}_2 \rightarrow 2 \mathrm{NO}_2\)

⇒ \(4 \mathrm{NO}_2+2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2 \rightarrow 4 \mathrm{HNO}_3\)

For welding and cutting of metals, oxy-hydrogen flame and oxy-acetylene flame are produced in which temperature as high as approximately 3000°C is achieved.

These flames are produced by the exothermic reactions between oxygen and hydrogen and between oxygen and acetylene. Liquid oxygen is used as rocket fuel.

Hydrogen

Hydrogen is the lightest element. It is diatomic. It is the most abundant element on earth. Nearly all hydrogen exists in the form of water.

In the combined state it is present in petroleum, coal, wood, oil and fats, etc. Organic substances also contain various compounds of hydrogen.

The interior of the sun contains a very large amount of hydrogen which is continuously converted into helium by a process called fusion, at a very high temperature (approximately 106 8C).

In this reaction, a huge amount of energy is released in the form of heat and light. In earth, free hydrogen occurs in the atmosphere in trace amounts (approximately one part in a million by volume).

Laboratory Preparation Of Hydrogen

Hydrogen is usually prepared in the laboratory by the reaction between dilute HCI or dilute H₂SO₄ with granulated zinc (or commercial zinc) in a Woulfe’s bottle (or a round bottom flask).

The hydrogen gas produced is collected by the downward displacement of water. The process is mentioned in detail below.

⇒ \(\mathrm{Zn}+\mathrm{H}_2 \mathrm{SO}_4\)

⇒ \(\mathrm{ZnSO}_4+\mathrm{H}_2\)

⇒ \(\mathrm{Zn}+2 \mathrm{HCl}\)

⇒ \(\mathrm{ZnCl}_2+\mathrm{H}_2\)

Theoretically, any metal which is more electropositive than hydrogen (i.e. situated higher than hydrogen in the electrochemical series) can replace hydrogen from dilute acids like dilute hydrochloric acid or dilute sulphuric acid.

The strongly electropositive metals like sodium and potassium are avoided because the reaction between them and dilute acid is very rapid and explosive in nature.

Chemicals and apparatus required A Woulfe’s bottle fitted with a two-holed cork, delivery tubes, thistle funnel, gas jar, water trough with a beehive shelf and a tripod stand, granulated commercial zinc, dilute sulphuric acid (or dilute hydrochloric acid).

[Commercial zinc is nothing but impure zinc available in markets].

Procedure: A few granules of commercial zinc are taken in a Woulfe bottle. The two-holed cork is tightly fitted. Through one of the holes, a thistle funnel is inserted in such a way that the end of which very nearly touches the bottom of the bottle.

Through the other hole of the cork passes a delivery tube, the end of which in the bottle is just below the cork. A small quantity of water is poured down the thistle funnel such that the end of the thistle funnel is well below the water level.

 

Step-by-Step Guide to Preparing Hydrogen in the Lab

The apparatus is made airtight, for hydrogen produces an explosive mixture with oxygen present in the air.

Making the apparatus airtight ensures that air from outside cannot enter the apparatus and any other gas can leak from inside the apparatus to the surrounding.

Moderately strong H₂SO₄ (or dilute HCI) is now poured down the thistle funnel.

A brisk reaction takes place, forming zinc sulphate (or zinc chloride, if dilute HCI is used) with the evolution of hydrogen gas.

⇒ \(\mathrm{Zn}+\mathrm{H}_2 \mathrm{SO}_4  \rightarrow \mathrm{ZnSO}_4+\mathrm{H}_2\)

⇒ \(\mathrm{Zn}+2 \mathrm{HCl} \rightarrow \mathrm{ZnCl}_2+\mathrm{H}_2\)

The gas is allowed to escape through the delivery tube for one to two minutes so as to carry away the air inside the bottle.

Then the hydrogen gas produced is allowed to collect in a gas jar (initially filled up with water, and kept inverted on a water-filled trough) by downward displacement of water.

1. Precautions

1. The end of the thistle funnel must remain below the liquid level in Woulfe’s bottle; otherwise, the gas will escape through it.

2. The apparatus must be air-tight. If air gets inside the bottle, oxygen present in the air will mix with the hydrogen gas produced. If this mixture somehow comes in contact with flame, it will explode.

3. There should be no air bubbles in the water used to collect the hydrogen gas. A mixture of oxygen and hydrogen will explode if it comes In contact with a flame.

4. Since hydrogen gas is inflammable, there must not be any kind of flame near the apparatus when it is being used for the preparation of hydrogen gas.

Preparation of Hydrogen

The metals which are situated on the left of hydrogen in the electrochemical series can liberate hydrogen from dilute, aqueous acid.

For example, Zn, Mg, Fe, Al, etc. all are situated on the left of hydrogen in the electrochemical series and can liberate hydrogen from dilute hydrochloric acid or dilute sulphuric acid.

⇒ \(\begin{aligned}
& \mathrm{Zn}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{ZnSO}_4+\mathrm{H}_2 \\
& \mathrm{Mg}+2 \mathrm{HCl} \rightarrow \mathrm{MgCl}_2+\mathrm{H}_2 \\
& \mathrm{Fe}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{FeSO}_4+\mathrm{H}_2 \\
& 2 \mathrm{Al}+6 \mathrm{HCl} \rightarrow 2 \mathrm{AlCl}_3+3 \mathrm{H}_2
\end{aligned}\)

2. Cold water is rapidly decomposed by metals such as lithium, sodium, potassium etc. and hydrogen is produced. Calcium, barium, etc. can slowly decompose in cold water and can produce hydrogen.

⇒ \(\begin{aligned}
2 \mathrm{Na}+2 \mathrm{H}_2 \mathrm{O} & \rightarrow 2 \mathrm{NaOH}+\mathrm{H}_2 \\
2 \mathrm{~K}+2 \mathrm{H}_2 \mathrm{O} & \rightarrow 2 \mathrm{KOH}+\mathrm{H}_2 \\
\mathrm{Ca}+2 \mathrm{H}_2 \mathrm{O} & \rightarrow \mathrm{Ca}(\mathrm{OH})_2+\mathrm{H}_2
\end{aligned}\)

3. Magnesium, aluminium, zinc, etc. decompose in boiling water and can liberate hydrogen.

⇒ \(\begin{aligned}
& 2 \mathrm{Al}+6 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{Al}(\mathrm{OH})_3+3 \mathrm{H}_2 \\
& \mathrm{Zn}+2 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Zn}(\mathrm{OH})_2+\mathrm{H}_2
\end{aligned}\)

Red hot magnesium, zinc, iron, etc. can decompose steam and produce hydrogen.

⇒ \(\begin{aligned}
& \mathrm{Zn}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{ZnO}+\mathrm{H}_2 \\
& \mathrm{Mg}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{MgO}+\mathrm{H}_2
\end{aligned}\)

4. Hydrogen is formed due to a reaction between some metals and alkali (viz. NaOH). A hot and concentrated NaOH.. solution dissolves zinc, aluminium, etc and hydrogen is evolved.

⇒ \(\begin{aligned}
& \mathrm{Zn}+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{ZnO}_2+\mathrm{H}_2 \\
& 2 \mathrm{Al}+2 \mathrm{NaOH}+2 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{NaAlO}_2+3 \mathrm{H}_2
\end{aligned}\)

Non-metals like silicon also react with a strong alkali like NaOH or KOH to produce hydrogen.

⇒ \(\mathrm{Si}+2 \mathrm{NaOH}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Na}_2 \mathrm{SiO}_3+2 \mathrm{H}_2\)

5. Electrolysis of water: Electrolysis of water acidified with dilute sulphuric acid can produce hydrogen at the cathode and oxygen gas at the anode. A platinum electrode is used as an anode and cathode.

If two gas jars filled with water are inverted carefully over the two electrodes, then the gases produced at each electrode can be collected within the gas jars by downward displacement of water.

6. Metallic hydrides like LiH, CaH₂, etc. react with water to form hydrogen.

⇒ \(\mathrm{LiH}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{LiOH}+\mathrm{H}_2\)

⇒ \(\mathrm{CaH}_2+2 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_2+2 \mathrm{H}_2\)

1. Physical Properties of Hydrogen
2. Hydrogen is a colourless, odourless and tasteless gas.
3. Hydrogen is almost insoluble in water.
4. Hydrogen is the lightest gas. Air is about. 14.4 times heavier than hydrogen.

Hydrogen may be condensed with difficulty to a colourless, transparent liquid. The melting point of hydrogen is -259.2°C and the boiling point of hydrogen is -252.6°C.

H It is the best conductor of heat among all gases. Hydrogen has three isotopes: Jh (ordinary hydrogen or protium),(heavy hydrogen or deuterium) and (tritium).

The natural abundance of the last two is negligibly small. Show with an experiment that hydrogen is lighter than air.

The density of hydrogen is 0.0899 grams per litre at 0°C and 1-atmosphere pressure.

Isotope	11H	21H	31H
Natural abundance	99.9844%	0.0156%	Negligibly small

 

Experiment-1

A balloon is filled with hydrogen gas. When inflated, its mouth is tied with a cord. If the balloon is now released, it goes up and touches the roof of the room where this experiment is being carried out.

This happens because the balloon filled up with hydrogen gas is lighter than the air displaced by it.

Experiment- 2

Two gas jars marked A and B are taken. One of the gas jars A is filled with hydrogen gas and is covered by a lid. The open mouth of another gas jar B is held upside down over gas jar A and the lid is then slowly removed.

Now a burning taper is introduced inside the gas jar B.

A “pop” sound is heard. The taper extinguishes but the gas burns with a bluish flame. This confirms that the gas in gas jar B is hydrogen.

Actually, hydrogen is lighter than air moves upwards and is collected in gas jar B by downward displacement of air.

Covered by a lid. The open mouth of another gas jar B is held upside down over gas jar A and the lid is then slowly removed.

Now a burning taper is introduced inside the gas jar B. A “pop” sound is heard. The taper extinguishes but the gas burns with a bluish flame. This confirms that the gas in gas jar B is hydrogen. Actually, hydrogen is lighter than air moves upwards and is collected in gas jar B by downward displacement of air.

Chemical Properties of Hydrogen

1. Hydrogen is a diatomic gas. When the gas is passed through an electric discharge (by using two tungsten electrodes) at a very high temperature (approximately 2000°C), it dissociates into two hydrogen atoms, called atomic hydrogen. This is an endothermic process. Atomic hydrogen is a very powerful reducing agent.

Hydrogen is not a supporter of combustion (i.e. does not allow substances to burn in it) but it is inflammable. It burns in air or oxygen with a very pale blue flame to form water.

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

Reaction with Non-metals: Some non-metals react with hydrogen under suitable conditions to form hydrides of the non-metal. [Binary compounds of elements with hydrogen are called hydrides.]

1. When a mixture of hydrogen and chlorine is exposed to diffused sunlight, Hydrogen chloride is produced. The reaction does not occur in absence of light.

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

2. At very high temperatures (approximately 550°C) and high pressure (approximately 200 atm) in presence of an iron catalyst, nitrogen combines with hydrogen to produce ammonia, which is a pungent-smelling gas. This is the industrial method (known as Haber’s process) to prepare ammonia.

⇒ \(\mathrm{N}_2+3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3\)

3. If hydrogen gas is passed over molten sulphur at a high temperature (approximately 600°C), H₂S gas is produced. The gas has a rotten egg-like smell.

⇒ \(\mathrm{H}_2+\mathrm{S} \rightarrow \mathrm{H}_2 \mathrm{~S}\)

4. Reaction with metals: Hydrogen gas reacts with metals like lithium, sodium, calcium, potassium, etc. to form metallic hydrides. These hydrides can produce hydrogen when they react with water. For example,

⇒ \(\begin{aligned}
& \begin{array}{l}
2 \mathrm{Li}+\mathrm{H}_2 \rightarrow 2 \mathrm{LiH}\left(680^{\circ} \mathrm{C}\right) \text {; } \\
\mathrm{LiH}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{LiOH}+\mathrm{H}_2 \\
2 \mathrm{Na}+\mathrm{H}_2 \rightarrow 2 \mathrm{NaH}
\end{array} \\
& \quad\left(300^{\circ}-400^{\circ} \mathrm{C}\right. \text { under pressure); } \\
& \mathrm{NaH}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{NaOH}+\mathrm{H}_2 \\
& 2 \mathrm{~K}+\mathrm{H}_2 \rightarrow 2 \mathrm{KH} \\
& \mathrm{KH}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{KOH}+\mathrm{H}_2 \\
& \mathrm{Ca}+\mathrm{H}_2 \rightarrow \mathrm{CaH}_2\left(150^{\circ} \mathrm{C}-300^{\circ}\right) ; \\
& {[\text { Calcium hydride is also known as }} \\
& \text { hydrolith.] } \\
& \mathrm{CaH}_2+2 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_2+2 \mathrm{H}_2
\end{aligned}\)

5. Reduction: Reaction with some metallic oxides:

Due to the great affinity of hydrogen towards oxygen, many heated metallic oxides react with hydrogen to produce corresponding metals. This occurs as hydrogen readily combines with the oxygen atoms present in the metallic oxides.

For example, when hydrogen gas is passed over hot cupric oxide (CuO), CuO is reduced to form metallic copper and hydrogen is oxidized to water.

⇒ \(\mathrm{CuO}+\mathrm{H}_2 \rightarrow \mathrm{Cu}+\mathrm{H}_2 \mathrm{O}\)

Similarly, PbO is reduced to metallic lead when PbO reacts with hydrogen.

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

6. Adsorption: Some metals like palladium, platinum, iron, nickel, etc. can adsorb hydrogen at normal temperatures. Palladium adsorbs the largest volume of hydrogen at 0°C.

This is known as occlusion. In this case, hydrogen is attached to the surface of these metals and is called adsorbed hydrogen. When heated, the adsorbed hydrogen is released.

Experimentally it has been found that adsorbed hydrogen is more reactive than normal hydrogen.

 

Tests For Hydrogen

Hydrogen burns in oxygen or in the air with a pale blue flame and produces water. Palladium absorbs hydrogen. The absorbed hydrogen evolves on heating the palladium used.

Nascent Hydrogen

Hydrogen, just at the moment of its birth or generation in a reaction, is called nascent hydrogen. It is chemically very active and a much more powerful reducing agent than ordinary molecular gaseous hydrogen (H₂).

Nascent hydrogen is generally obtained by generating the gas inside the reaction medium in presence of reactants. This is called in situ preparation. Hydrogen liberated at the cathode during electrolysis is endowed with properties similar to nascent hydrogen.

Uses of Hydrogen

Large quantity of hydrogen is used for the production of ammonia \(\left(\mathrm{N}_2+3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3\right)\).

This reaction is carried out at high pressure (approximately 200 atm) and high temperature (approximately 550°C) in presence of a suitable catalyst (such as iron).

Fertilizers like urea, ammonium nitrate and ammonium sulphate and nitric acid are prepared from ammonia. Hydrogen is used commercially to reduce metal oxides to metals.

Hydrogen is used for the hardening of fats and oils (where unsaturated fatty acids are converted to saturated compounds which have a higher melting point); hence can be used to prepare more useful products like margarine.

Many fuel gases (water gas, coal gas etc) contain hydrogen as a major constituent because of the high heat of combustion of hydrogen.