CBSE Class 7 Science Advanced Chemistry Air and Oxygen Notes Set 01

Air and Oxygen

Atmosphere

Of all the things the earth gives us, air matters most, because we cannot live without it even for a few minutes. Air is a homogeneous mixture of gases. The blanket of gases wrapped around the earth is the atmosphere. It is dense near the ground and becomes steadily thinner as we go up, until it fades into outer space, which has no air at all.

The atmosphere contains the air we breathe, water vapour and dust, which together bring about clouds and rain. It also works as a shield against harmful radiation from space. Depending on how composition, density and temperature change with height, the atmosphere is split into five layers.

Layers of the Atmosphere

1. Troposphere

1. It is the lowest layer, starting at the ground and reaching about 20 km up.
2. It contains the largest share of the mass of the whole atmosphere.
3. Air density and temperature both fall as we climb. The temperature drops from about 17°C at the bottom to about −51°C at the top. A fall of about 6°C for every kilometre of height is called the lapse rate.
4. Because temperature changes quickly here, nearly all weather events (clouds, rain, winds) happen in this layer.
5. Its thickness is not the same everywhere; it differs between the equator and the poles.

2. Stratosphere

1. It lies above the troposphere and reaches up to about 50 km from the ground.
2. It holds about 19% of the gases of the atmosphere.
3. Up to about 25 km the temperature stays nearly constant. Above that it rises with height, from about −60°C to −15°C, because the gases absorb high-energy ultraviolet (UV) radiation. This rise of temperature with height is called temperature inversion.
4. The part between about 25 km and 40 km is rich in ozone and is called the ozonosphere or ozone layer. It absorbs harmful UV rays and so protects living things.
5. Aeroplanes fly in this layer.

3. Mesosphere

1. It lies above the stratosphere and extends to about 85–90 km above the ground.
2. The temperature falls again, from about −15°C to −120°C.
3. Meteors (shooting stars) burn up in this layer.

4. Thermosphere (Ionosphere)

1. This is the upper atmosphere, reaching up to about 690–950 km.
2. The air is thinner than in the mesosphere. Its molecules absorb high-energy radiation and become ionised, so the layer is full of ions and free electrons.
3. The temperature rises with height because of this absorption and can touch about 2000°C at the top.
4. Radio waves bounce back from this layer, which makes long-distance radio communication possible. TV waves are not reflected, so artificial satellites are used to relay television signals.
5. The region beyond the ionosphere is called the valley of silence. Auroras form in this layer.

Auroras: These are glowing bands or curtains of light, seen near the northern and southern polar regions, that can stretch up to about 320 km in height. They are produced when fast-moving electrons from the sun meet the earth's magnetic field and collide with molecules such as nitrogen and oxygen. The electrons of these atoms jump to higher energy levels and give out coloured light as they fall back.

5. Exosphere

1. It is the outermost layer and extends to about 10,000 km from the ground.
2. It absorbs harmful ionising radiation.

The figure shows the five layers from the bottom: troposphere (with Mount Everest and an aeroplane), stratosphere (weather balloon), mesosphere (meteors), thermosphere (shuttle and aurora) and exosphere.

Composition of Air

Until the 15th century air was thought to be a single element. In 1674 Mayow suggested that air has two parts, one active and one inactive. In 1789 Lavoisier stated that oxygen and nitrogen are present in air in the volume ratio 1 : 4. Later experiments confirmed that these two are the main gases of air.

Nitrogen and oxygen together make up roughly 99% of dry air by volume. The rest includes water vapour, carbon dioxide, nitrous oxide, methane, ozone, hydrogen and traces of the inert gases neon, helium and argon. The make-up of air is not the same everywhere; it changes from place to place.

Atmospheric Pressure

Air has weight, so the layers of air press downwards on everything below. This pressure is called atmospheric pressure. Its SI unit is the pascal (Pa) or kilopascal (kPa), and it is measured with a barometer. As we go higher the air becomes less dense, so the pressure drops. It is greatest at sea level, where it equals 1 atm or 760 mm of mercury.

Importance of Nitrogen and Oxygen

Nitrogen is the inactive part of air; it weakens the action of oxygen so that burning and other reactions do not become too fierce. Oxygen is the active part. It is needed for respiration and combustion, and so supports life.

Importance of Inert Gases and Water Vapour

Example 1. What is the role of water vapour in the formation of fog?
Solution: Water vapour in the air affects the weather. On cold nights, especially in winter, it condenses on dust particles. Depending on the temperature it can form snow, fog or mist. When the condensed droplets stay hanging in the air around dust particles, we call it fog.

Example 2. Give any three ways in which the atmosphere is important.
Solution: 1. The air in it is what we breathe.
2. Its moisture and dust help in forming clouds.
3. It protects us from harmful radiation.

Example 3. Name the layers of the atmosphere.
Solution: Troposphere, stratosphere, mesosphere, thermosphere (ionosphere) and exosphere.

Oxygen

History

The Swedish chemist Carl Scheele obtained oxygen by heating nitre and mercuric oxide. Because the gas helped things burn, he called it "fire air". The English chemist Joseph Priestley made it by focusing sunlight with a convex lens on mercuric oxide in a test tube. Antoine Lavoisier showed that it is an element and, believing it to be a necessary part of every acid, named it oxygen, meaning "acid former".

Occurrence

Oxygen is found both free and combined, and it is the most plentiful element in the earth's crust.
1. Free state: Air contains about 21% oxygen by volume and 23% by mass. Oxygen is also dissolved in water, which lets fish and other water animals breathe.
2. Combined state: It is present in water, in plant and animal bodies, and in minerals such as silicates, carbonates and metal oxides, which hold about half of all the oxygen.

Preparation of Oxygen

General methods

Method 1 is not used in the laboratory because mercury vapour is poisonous, mercuric oxide is costly and only a little oxygen is obtained. Method 2 is not used because the gas comes off slowly and strong heating is needed.

Laboratory method: from hydrogen peroxide

Principle: Hydrogen peroxide breaks down into water and oxygen when a little manganese dioxide (\(\ce{MnO2}\)) is present. The manganese dioxide makes the reaction fast and lets it go on at room temperature.

\[ \ce{2H2O2 ->[MnO2] 2H2O + O2 ^} \]

A catalyst is a substance that changes the speed of a reaction without being used up in it.

Procedure: Put manganese dioxide in a clean, dry conical flask closed with a stopper. Add hydrogen peroxide solution drop by drop through a thistle funnel fitted into the stopper.
Observation: The oxygen leaves through the delivery tube and passes into a gas jar filled with water and inverted over a beehive shelf in a trough of water. Bubbles rise into the jar and push the water out. The gas relights a glowing splinter.
Collection: Oxygen is collected by downward displacement of water.
Precautions: 1. Do not heat the mixture. 2. Use pure manganese dioxide.

Physical Properties of Oxygen

Note: The vapour density of a gas is the mass of a given volume of the gas divided by the mass of the same volume of hydrogen, both measured at the same temperature and pressure. It is therefore a comparison, not an absolute value.

Chemical Properties of Oxygen

1. Reaction with metals

Metals combine with oxygen to give metal oxides, which are usually basic.

2. Reaction with non-metals

Non-metals combine with oxygen to give non-metal oxides, which are usually acidic.

3. Other chemical properties

1. Combustibility: Oxygen does not burn itself, so it is non-combustible, but it helps other substances burn, so it is a supporter of combustion. A glowing splinter bursts into flame in it.
2. Action on litmus: Oxygen is neutral; it changes the colour of neither blue nor red litmus.
3. Rusting: When iron is exposed to moist air it slowly changes into a hydrated oxide, called rust.

\[ \ce{4Fe + 3O2 -> 2Fe2O3} \qquad \ce{Fe2O3 + xH2O -> Fe2O3.xH2O} \]

Here x is a variable number of water molecules. Rust is a brown, flaky solid that crumbles off the metal and so weakens iron structures. Iron does not rust in dry air.

Prevention of Rusting

Iron is used for machines, tools, bridges, rails, furniture, cooker bodies, refrigerators, buckets and many other things, so it has to be protected. The methods are the following.

A. Covering the iron so that moist air cannot reach it
1. Red lead paint or tar: Bridges, electric poles and similar structures are painted with red lead oxide paint or coated with tar.
2. Enamel paint: Doors, windows and car and bus bodies get a first coat of red lead paint and then an enamel coat, which also improves the look.
3. Enamelling: Coating the iron surface with a mixture of silicates at high temperature. It is used for cooker bodies, refrigerators, plates, wash basins and ovens.
4. Oil or grease: Moving parts of machines and engines are smeared with oil or grease, which prevents rusting and also lubricates them. Paint would wear off these moving parts.

B. Galvanisation: Iron sheets are dipped in molten zinc and passed between heavy rollers, so they get a zinc coating. The product is galvanised iron, used for buckets, tubs and iron trunks.

C. Tinning: Coating iron sheets with a layer of tin, in a manner similar to galvanisation. Tinned iron is used for containers that store edible things such as oil, meat and vegetables.

D. Electroplating: Iron is coated with chromium or nickel using electrolysis. Examples are bicycle handles, rims and car bumpers.

E. Alloying: Melting a metal and mixing small amounts of other substances in it to improve its properties. Adding carbon and chromium to iron gives stainless steel, which resists corrosion. Typical composition: chromium 12% to 20%, carbon 0.07% to 0.1%. It is used for surgical instruments, tools, utensils and cutlery.

Laboratory Tests for Oxygen

1. It does not affect litmus paper because it is neutral.
2. It rekindles a glowing splinter.
3. It combines with colourless nitric oxide to give reddish-brown fumes of nitrogen dioxide.
4. It turns alkaline pyrogallol solution brown.

Uses of Oxygen

1. Respiration: The oxygen we breathe in oxidises glucose in the body to carbon dioxide and water and releases energy.

\[ \ce{C6H12O6 + 6O2 -> 6CO2 + 6H2O + energy} \]

2. It is needed to burn fuels.
3. Breathing aid: Divers, submarine crews, mountaineers, astronauts and pilots carry oxygen where the natural supply is poor.
4. Medical use: A mixture of 95% oxygen and 5% carbon dioxide, called carbogen, is given to patients with breathing trouble.
5. Cutting and welding: Hydrogen burning in oxygen gives the oxyhydrogen flame (about 2800°C) and acetylene burning in oxygen gives the oxyacetylene flame (about 3300°C). These flames melt metals, so they are used for cutting and welding.
6. Rockets and spacecraft: Liquid oxygen is used to burn the fuel.
7. Explosives: A mixture of coal, petroleum jelly and liquid oxygen, called a cartridge, is used for blasting rocks.
8. Industry: It is used to make sulphuric acid and nitric acid.

Carbon Dioxide

History

Van Helmont found that a gas forms when charcoal burns and named it "carbonium"; he also noticed it collects in coal mines and puts out a candle flame. Joseph Black obtained it by heating magnesium carbonate and called it "fixed air". Lavoisier showed it to be an oxide of carbon, and under the IUPAC system it is called carbon dioxide.

Occurrence

Carbon dioxide forms about 0.03% of the volume of air (Table above) and is the fifth most plentiful gas in the atmosphere.
1. Free state: in air, and dissolved in oceans, lakes and rivers.
2. Combined state: in carbonate minerals such as limestone, magnesite and dolomite.

Preparation of Carbon Dioxide

General methods

1. Burning carbon-containing fuels: coal, natural gas and similar fuels release carbon dioxide on burning.

\[ \ce{C + O2 -> CO2 ^ + heat} \qquad \ce{CH4 + 2O2 -> CO2 ^ + 2H2O + heat} \]

2. Heating metal carbonates: Most carbonates break down on strong heating into the metal oxide and carbon dioxide. Sodium carbonate and potassium carbonate are exceptions.

\[ \ce{CaCO3 ->[\Delta] CaO + CO2 ^} \qquad \ce{CuCO3 ->[\Delta] CuO + CO2 ^} \qquad \ce{ZnCO3 ->[\Delta] ZnO + CO2 ^} \]

3. Heating hydrogen carbonates: Sodium hydrogen carbonate (baking soda) on heating gives sodium carbonate, water and carbon dioxide.

\[ \ce{2NaHCO3 -> Na2CO3 + H2O + CO2 ^} \]

4. Action of dilute acids on carbonates and hydrogen carbonates: carbon dioxide is released along with a salt.

\[ \ce{Na2CO3 + H2SO4 -> Na2SO4 + H2O + CO2 ^} \qquad \ce{Ca(HCO3)2 + 2HNO3 -> Ca(NO3)2 + 2H2O + 2CO2 ^} \]

Laboratory method

Principle: Dilute hydrochloric acid acts on marble chips (calcium carbonate) and gives carbon dioxide.

\[ \ce{CaCO3 + 2HCl -> CaCl2 + H2O + CO2 ^} \]

Procedure: A conical flask is closed with a stopper having two holes. Dilute hydrochloric acid is added through a thistle funnel passing through one hole, onto the marble chips. The gas leaves through a delivery tube in the other hole and enters a gas jar.

Collection: Carbon dioxide is heavier than air, so it is collected by upward displacement of air.

Note: Dilute sulphuric acid is not used. The calcium sulphate formed is insoluble and coats the marble chips, which stops the acid from reaching them and halts the reaction.

\[ \ce{CaCO3 + H2SO4 -> CaSO4 v + H2O + CO2 ^} \]

Physical Properties of Carbon Dioxide

Chemical Properties of Carbon Dioxide

Uses of Carbon Dioxide

Carbon dioxide is used in industry, medicine and agriculture. It is the raw material plants use to make food by photosynthesis, and it helps keep the earth warm. Its uses are listed below.

1. Making chemicals: washing soda, baking soda (NaHCO₃) and white lead (Pb(OH)₂.PbCO₃).
2. Soft drinks: It is dissolved under pressure in aerated drinks to give a sharp taste.
3. Medicine: Carbogen (95% O₂ and 5% CO₂) stimulates breathing and is given to patients with asthma or pneumonia.
4. Refrigeration: Dry ice keeps meat, fish, fruit and vegetables fresh for a long time.
5. Preserving food: It protects stored food grains.
6. Photosynthesis: Plants use carbon dioxide and water, in sunlight, to make glucose and release oxygen back into the air.

\[ \ce{6CO2 + 6H2O ->[sunlight] C6H12O6 + 6O2 ^} \]

7. Fire extinguishers: Carbon dioxide does not burn, does not support burning and is heavier than air, so it covers a fire and puts it out.

Laboratory Tests for Carbon Dioxide

1. A burning candle goes out in it.
2. It turns lime water milky, and the milkiness disappears when more carbon dioxide is passed.

Nitrogen

Occurrence

1. Free nitrogen makes up about 78% of air by volume.
2. In the combined state it is found in salts such as potassium nitrate (KNO₃) and sodium nitrate (NaNO₃) and in proteins of living matter.

Uses

1. It dilutes oxygen and so prevents very fast burning.
2. It is used to make ammonia, whose compounds are fertilisers.
3. Nitrogen fixation: Plants and animals need nitrogen to make proteins, but plants cannot use free nitrogen of the air. Turning atmospheric nitrogen into compounds, such as nitrates, that plants can take up is called nitrogen fixation. It happens in two ways.

(a) By lightning: The spark of lightning makes nitrogen combine with oxygen to give nitric oxide, which takes up more oxygen to become nitrogen dioxide. This gas dissolves in cloud droplets to give nitric acid. The acid falls with rain and reacts with metal carbonates in the soil to form nitrates, which the plants absorb.

\[ \ce{N2 + O2 -> 2NO} \qquad \ce{2NO + O2 -> 2NO2} \qquad \ce{4NO2 + 2H2O + O2 -> 4HNO3} \qquad \ce{2HNO3 + MCO3 -> M(NO3)2 + CO2 + H2O} \]

(b) By bacteria: Leguminous plants have root nodules containing symbiotic bacteria (rhizobium). These bacteria take nitrogen from the air and change it into nitrogen compounds, which the plant can use and which also enrich the soil.

4. Because nitrogen is chemically unreactive (inert), it is used to preserve food.

Hydrogen

Hydrogen is the lightest of all elements and occurs only in traces in air. It was first made by the action of sulphuric acid on iron. In 1762 Henry Cavendish prepared it from zinc and sulphuric acid and showed it to be an element. Lavoisier named it hydrogen, meaning "water former" (hydro = water, gen = producing).

Occurrence

Free hydrogen is rare. It is present in traces in air, in volcanic gases, and up to 1.1% in the sun's atmosphere. Most hydrogen is combined: in proteins, nucleic acids, vitamins, hormones and enzymes of living things, and in wood, cotton, paper pulp and rubber. Water is its most important source, containing 11.1% hydrogen by weight.

Preparation of Hydrogen

General methods

Laboratory method

Principle: Any metal above hydrogen in the reactivity series releases hydrogen from dilute hydrochloric or sulphuric acid. Zinc is the preferred metal.

\[ \ce{Zn + 2HCl -> ZnCl2 + H2 ^} \qquad \ce{Zn + H2SO4 -> ZnSO4 + H2 ^} \]

Procedure: Granulated zinc is placed in a flat-bottomed flask closed by an air-tight cork with two holes, one for a thistle funnel and one for a delivery tube. Dilute sulphuric acid is added dropwise through the funnel.
Observation: Bubbles appear at first, and then a colourless, odourless gas is given off.
Collection: By downward displacement of water.
Purification: Commercial zinc is impure, so the hydrogen contains traces of H₂S, SO₂, PH₃, AsH₃ and CO₂. These are removed step by step.

The gas passes in turn through silver nitrate solution, lead nitrate solution, potassium hydroxide solution and a U-tube of anhydrous calcium chloride. Pure, dry hydrogen is then collected over mercury.

Precautions: 1. Keep the apparatus air-tight. 2. Keep flames away. 3. Start collecting only after all the air has been driven out. 4. The end of the thistle funnel should almost touch the bottom of the flask.

Although all active metals can give hydrogen, zinc is chosen. The reasons for not using others are in the table.

Industrial methods

1. Electrolysis of water: Passing electricity through acidified water gives oxygen at the anode and hydrogen at the cathode. This is mainly a method for oxygen, with hydrogen as a by-product. \(\ce{2H2O -> 2H2 + O2}\)
2. Bosch process: Steam passed over red-hot coke gives water gas, a mixture of carbon monoxide and hydrogen. \(\ce{C + H2O -> CO + H2}\). Water gas and more steam, passed over a catalyst of Fe₂O₃ at 450°C to 500°C, turn the carbon monoxide into carbon dioxide and leave hydrogen. \(\ce{CO + H2 + H2O ->[Fe2O3] CO2 + 2H2}\)

Steps of the process:
(a) Superheated steam at about 170°C is passed over red-hot coke to get water gas.
(b) Water gas and steam at about 500°C go over Fe₂O₃ or Cr₂O₃, which oxidises CO to CO₂.
(c) The CO₂ is removed by compressing the gas to 30 atm and passing it through water, which dissolves it. Hydrogen is insoluble and stays behind.
(d) Water vapour is removed by cooling the gas to −20°C.
(e) Traces of unconverted CO are removed with ammoniacal cuprous chloride solution. \(\ce{CuCl + CO + 2H2O -> CuCl.CO.2H2O}\)

The Bosch process can also start from natural gas, whose main constituent is methane. Methane mixed with superheated steam at high temperature and compressed to 30 atm gives water gas, which is then treated as above.

Physical Properties of Hydrogen

Adsorption (Occlusion)

Adsorption means the gathering of molecules of a substance on the surface of another. Hydrogen is adsorbed on the surface of metals such as palladium, platinum and nickel, and this property is called occlusion.

Chemical Properties of Hydrogen

1. With air: It burns in air with a blue flame and forms water. \(\ce{2H2 + O2 -> 2H2O}\)
2. With non-metals: It forms hydrides.
(a) With chlorine in sunlight (a photochemical reaction): \(\ce{H2 + Cl2 ->[light] 2HCl}\)
(b) With nitrogen, in the ratio 1 : 3, at high temperature and pressure: \(\ce{N2 + 3H2 <=> 2NH3 + heat}\)
(c) With molten sulphur: \(\ce{H2 + S -> H2S}\)
3. With metals: With highly electropositive metals it forms metal hydrides. \(\ce{2K + H2 ->[\Delta] 2KH}\), \(\ce{2Na + H2 ->[\Delta] 2NaH}\), \(\ce{Ca + H2 -> CaH2}\)
4. With metal oxides: The oxides of less reactive metals are reduced to the metal, and water is formed.

5. Hydrogenation: The addition of hydrogen to some carbon compounds in the presence of a catalyst. \(\ce{H2C=CH2 + H2 -> H3C-CH3}\) (ethylene to ethane). This is used to harden oils. Vegetable oils such as groundnut and sunflower oil have relatively little hydrogen and stay liquid. Animal fats like ghee have more hydrogen and set easily. When hydrogen is passed through vegetable oil at high temperature and pressure with nickel as catalyst, the oil turns into a solid fat, as in vanaspati.

Uses of Hydrogen

1. Making chemicals: ammonia by Haber's process \(\ce{N2 + 3H2 <=> 2NH3 + heat}\); hydrogen chloride \(\ce{H2 + Cl2 -> 2HCl}\); methyl alcohol from CO and H₂ with a catalyst; vanaspati and soap by hydrogenation; and petrol from coal.
2. Fuel: Fuel gases such as water gas and coal gas contain hydrogen, which gives a lot of heat on burning. Hydrogen is also used in fuel cells. The oxyhydrogen torch (about 2800°C) welds metals and melts platinum and quartz. Liquid hydrogen is a rocket fuel. In the atomic hydrogen torch an electric arc between tungsten electrodes splits hydrogen molecules into atoms; when the atoms recombine they release the energy as intense heat used for welding.
3. Metallurgy: It is a good reducing agent for getting metals such as tungsten and molybdenum from their oxides. \(\ce{WO3 + 3H2 -> W + 3H2O}\)

Nascent Hydrogen

Newly produced hydrogen, written [H], is called nascent hydrogen. It can reduce substances that ordinary hydrogen does not easily reduce. It is usually made by the action of metals on acids, such as zinc with sulphuric acid.
Applications: 1. It can be used as a rocket propellant; it would have to be stored in liquid helium to stop the atoms from combining into H₂ molecules. 2. It is said to reduce nitrites to ammonia and arsenic compounds to arsine even under ordinary conditions. 3. It has a temporary bleaching action.

Air Pollution

Human activities mainly change the natural composition of air by adding other gases and suspended particles, which harm living things. The change in the composition of air due to the release of harmful substances is air pollution. A substance that, when added to the atmosphere, harms people and surroundings is an air pollutant, for example CO, CFCs and ozone in the troposphere.

Main causes: burning fossil fuels, deforestation, volcanic eruptions, industrial and vehicle emissions, farming activities, lack of civic sense, pollen, and forest fires. Volcanoes, forest fires and pollen are natural causes.

Hazards of Air Pollution and Their Prevention

1. Acid rain: Nitrogen dioxide and sulphur dioxide from chimneys, fossil fuels and volcanoes dissolve in rain water to form nitric and sulphuric acids. This is acid rain.

\[ \ce{2SO2 + O2 -> 2SO3} \qquad \ce{SO3 + H2O -> H2SO4} \qquad \ce{2NO2 + H2O -> HNO3 + HNO2} \]

Effects: soil loses fertility; crops are damaged; rivers and seas are polluted; marble statues, monuments and buildings are eaten away.
Prevention: use less fossil fuel (CNG instead of petrol and diesel in some big cities); use non-conventional energy sources; dispose of industrial waste safely.

2. Greenhouse effect and global warming: The surface of the oceans and the land vegetation absorb sunlight and re-emit it as heat. Some gases in the air absorb this infrared radiation and send the heat back in all directions, which keeps the earth's surface warm. These gases are greenhouse gases and the process is the greenhouse effect; it supports life. Greenhouse gases are CO₂, water vapour, N₂O, CH₄, O₃ and CFCs. When their amount rises because of pollution, the earth's surface slowly gets hotter. This is global warming.
Effects: ice caps and glaciers melt; floods, soil erosion and unseasonal rain become more frequent.
Prevention: use less fossil fuel; stop deforestation; restrict CFCs.

3. Ozone depletion: Ozone in the troposphere adds to the greenhouse effect, but ozone in the stratosphere shields life by absorbing harmful UV rays. CFCs reach the stratosphere and destroy ozone over time, creating "holes". UV rays then reach the lower atmosphere and harm plants and animals.
Prevention: find substitutes for CFCs in refrigeration and aerosols.

4. Carbon monoxide hazard: Carbon monoxide combines with haemoglobin of blood about 200 times more strongly than oxygen does, forming stable carboxyhaemoglobin. This causes suffocation and can even be fatal.
Prevention: burn less fossil fuel and tobacco.

In general, pollutants harm plants and animals, cause breathing problems and damage monuments. Lead from vehicle smoke harms the nervous system.

Ways to Reduce Air Pollution

1. Dispose of waste safely.
2. Fit scrubbers and electrostatic precipitators in industries to stop harmful emissions.
3. Make eco-friendly chemicals.
4. Control the growth of population.
5. Use natural resources wisely.
6. Limit the use of plastic.
7. Stop deforestation and plant more trees.

Examples

Example 4. Write a short note on acid rain.
Solution: Gases like NO₂ and SO₂ enter the air from factory chimneys, burning fossil fuels and volcanic eruptions. They dissolve in rain water to give HNO₃ and H₂SO₄, causing acid rain. \(\ce{2SO2 + O2 -> 2SO3}\); \(\ce{SO3 + H2O -> H2SO4}\); \(\ce{2NO2 + H2O -> HNO3 + HNO2}\)

Example 5. Explain ozone depletion. How can it be prevented?
Solution: Stratospheric ozone absorbs harmful UV rays. CFCs that reach the stratosphere react with ozone and slowly thin the layer, so UV rays reach lower layers and cause skin cancer, cataract and damage to plants and crops. Using substitutes for CFCs in refrigeration and aerosols prevents the depletion.

Example 6. A combines with B to form C and D. A and D are major contributors to global warming and both are carbon compounds. B is a supporter of combustion. Magnesium burns in B with a dazzling white flame and forms a product X which is basic. X with C forms a compound that turns red litmus blue. Give the reactions.
Solution: A is methane, B is oxygen, C is water vapour (water) and D is carbon dioxide. Methane and CO₂ are carbon compounds that cause the greenhouse effect. \(\ce{CH4 + 2O2 -> CO2 + 2H2O}\). Magnesium burns in oxygen to form magnesium oxide, which is basic: \(\ce{2Mg + O2 -> 2MgO}\). Then \(\ce{MgO + H2O -> Mg(OH)2}\); magnesium hydroxide is a base and turns red litmus blue.

Points to Remember

1. Weather events such as clouds and rain occur in the troposphere; the stratosphere shields us from UV rays and is where aeroplanes fly.
2. Air is densest near the ground, in the troposphere.
3. Nitrogen is the inactive part of air and is used to preserve food; oxygen is the active part needed for burning and breathing.
4. Nitrogen dilutes oxygen, and its compounds are fertilisers.
5. In the laboratory oxygen is made from hydrogen peroxide and collected by downward displacement of water; carbon dioxide is made by dilute acid on marble chips and collected by upward displacement of air.
6. Carbon dioxide is needed for photosynthesis; being non-combustible, a non-supporter of combustion and heavier than air, it puts out fires. It is a greenhouse gas, and too much of it from deforestation and burning petrol and diesel leads to global warming.
7. Water vapour and dust form clouds, fog and mist.
8. Rust (Fe₂O₃.xH₂O) forms when iron meets moist air; it can be prevented by galvanisation, tinning, alloying, enamel paint and so on.
9. NO₂, SO₂, excess CO₂, N₂O, CO and dust are the major pollutants, and they cause acid rain, global warming, ozone depletion and breathing problems.
10. Planting trees, using CNG, fitting catalytic converters in vehicles, and using precipitators and scrubbers in industries reduce pollution.