Class 8 School Science

Chapter 4 Occurrence Of Carbon And Its Compounds In Nature Existence Of Carbon In Nature

Carbon in Life

The Latin word “carbo”, meaning coal, has been transformed into the English word carbon. Carbon is regarded as the building block of life.

Plant and animal bodies are made of carbon and its various compounds. For example, carbon is a major constituent of all biological molecules such as benzene, xylenes, etc. all as DNA (Deoxyribose Nucleic Acidjand RNA (Ribo Nucleic Acid), proteins, and carbohydrates, etc.

Small molecules like ATP, lipids, hormones, and enzymes – all are made up of carbon. Carbon comprises nearly 50% of the total mass of the body.

Source of Carbon

In a free state, carbon exists in crystalline states (such as diamond and graphite) as well as amorphous states (such as coal, charcoal, gas carbon, etc.).

In combined state carbon exists in various biomolecules such as carbohydrates, proteins, DNA, RNA, enzymes, etc. Oils and fats also contain carbon as one of their constituents.

In combination with hydrogen, it exists as hydrocarbons in the form of petroleum products, marsh gas, etc. Carbon is abundant in the earth’s crust.

As minerals, it exists in limestone (CaC03), dolomite (MgC03, CaC03), magnesite (MgC03), calamine (ZnC03), marble (CaC03), etc.

In the air, carbon exists as carbon dioxide and methane.

The shells of small marine animals such as mollusks are made of calcium carbonate. Even the materials used for making our garments (such as silk, wool, jute, etc.) are polymers of carbon.

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Uses Of Carbon

Several modern furniture, electronic equipment, sports accessories, etc. are being manufactured using polymers of carbon. Lifesaving drugs, solvents like alcohol, ketones, ether, and chloroform compounds of carbon.

Carbon compounds are abundant in nature, and a wide variety of man-made materials containing carbon are synthesized every day.

Carbon Cycle

The carbon cycle is the circulation and transformation of carbon back and forth between living things and the environment. The total amount of carbon present in the earth and the earth’s atmosphere is fixed. But the amount of carbon present in different compounds is always changing.

The carbon cycle is the bio-geo-chemical cycle by which carbon is exchanged or cycled among the earth’s oceans, atmosphere, ecosystem, and geosphere.

Occurrence Of Carbon And Its Compounds In Nature

The atmosphere acts as a reservoir of carbon in the form of carbon dioxide. Carbon is released to the atmosphere from what we call carbon sources and is stored in plants, animals, rocks and minerals, and water, which we call carbon sinks.
Now, let us discuss various processes of the carbon cycle.

Removal of carbon as CO₂ from

1. By fixation of atmospheric carbon dioxide: Carbon exists in the atmosphere mainly as CO₂ and methane (CH4). A large amount of CO₂ is utilized by green plants to produce glucose, using water and sunlight through a process called Photosynthesis. Thus carbon becomes part of their body.

Glucose is then converted into different complex organic molecules within the body. So atmospheric CO₂ is “fixed” by the green plants into different organic compounds in a living cell.

This carbon fixation step is also known as carbon assimilation. Some anaerobic bacteria can convert carbon monoxide into organic compounds.

 

WBBSE Class 8 Occurrence of Carbon notes

2. Through the formation of CaC03 by marine animals: Oceans contain very large amounts of carbon. Some amount of carbon dioxide present in the atmosphere dissolves in water and is converted into carbonates and bicarbonates. For example, dissolved CO₂ can be converted into carbonate rocks such as limestone.

Many marine animals also absorb CO₂, which is dissolved in water. Their shells contain carbon in the form of calcium carbonate (CaC03).

A part of the dissolved CO₂ is converted into organic carbon (i.e. carbon present in organic substances) by microorganisms through photosynthesis. In such a form, carbon can be exchanged throughout the food chain.

3. By conversion of CO₂ into carbonate minerals: Carbon dioxide in the atmosphere can be converted into carbonate rocks. Marble, limestone, dolomite, etc.

Are all carbonate minerals. Limestone caves have columns of stalactite and stalagmites which are made up of calcium carbonate. (Some limestone caves have been found in India also. Bora caves in Araku valley is one such example.)

By dissolving of atmospheric CO₂ by rainwater: The atmospheric CO₂ is dissolved in rainwater and is converted into carbonic acid. This is dissociated into bicarbonate ion (HC03)

1. Generation of CO₂ through respiration and oxidation of food: As we have just pointed] [out carbon in the earth’s atmosphere exists in two main forms carbon dioxide and methane.

Green plants absorb CO₂ from the atmosphere and are converted to glucose. When this glucose is oxidized within the plant body during respiration, CO₂]), which is ultimately converted into carbonate compounds such as calcium carbonates formed and released into the atmosphere.

Animals and humans inhale air (containing oxygen) and exhale CO₂ produced by the metabolism of food which is derived directly or indirectly from green plants.

2. By degradation of dead organisms: Living organisms are composed of organic compounds. After the organism dies, its body is broken down by bacterial and fungal action.

There are several different types of bacteria and fungi that can degrade complex organic compounds containing carbon into simple compounds.

Due to such processes, some carbon compounds remain stored in soil and water, and some return back to the atmosphere in the form of CO₂.

Methanogenic bacteria which are available in wetlands and rain forests can decompose dead plants into methane (CH4), which is released into the atmosphere.

Microbes present in the rumen of cattle and intestines of termites can produce methane from cellulose.

3. By human activities: The continuous and steady increase in the concentration of CO₂ in the atmosphere is mostly caused by human activities (i.e. anthropogenic).

These include the burning of carbonaceous fossil fuels (coal, oil, natural gas), clearing and burning of forests (mainly for agriculture and cattle(grazing), cement manufacture, etc.

4. By natural phenomena: Natural phenomena like volcanic eruptions and forest fires also contribute large amounts of carbon dioxide and carbon monoxide to the atmosphere.

5. Due to saturation of dissolved CO₂ in ocean water: Ocean water has huge quantities of CO₂ dissolved in it which is vital for marine plants.

Scientists have estimated that almost 50% of CO₂ having anthropogenic origin is absorbed in the oceans. But the steep increase in the production of CO₂ by human activities severely imbalances the capacity of the oceans to store additional CO₂ and as a result, a large amount of CO₂ remains unabsorbed.

6. By marine organisms: Carbon is stored in the shells of marine organisms as calcium carbonate (CaC03). When these shells are heated at high temperatures, CO₂ is released into the atmosphere.

Allotropy

Substances such as diamond, graphite, charcoal, coke, coal, etc. may physically appear different, but – they are all made of carbon.

If a fixed mass of all these substances is heated strongly in presence of pure oxygen, the same mass of carbon dioxide is formed in all the cases.

The phenomenon in which some elements like carbon, sulfur, phosphorus, etc. exist in more than one form, generally having the same chemical properties but some differences in physical properties in the same physical state is known as allotropy.

The different forms of an element having the same chemical properties but with different physical properties are called allotropes.

Causes of allotropy:

1. The difference in the mode of arrangement of atoms in molecules
2. Number of atoms present
3. Different methods of formation
4. Different amounts of internal energy are associated with the formation of each allotrope
5. Allotropes of Carbon
6. The element carbon has a number of allotropes and they may be classified as follows,

 

Carbon compounds in nature for Class 8

Generally, the allotropes of carbon can be seen in two forms.

1. Crystalline and
2. Amorphous.

1. Crystalline forms of carbon

1. Diamond

It is one of the hardest substances known to us. It is a transparent, hard solid with a high refractive index. This hardness of a diamond can be related to its internal structure.

Here, every carbon atom is surrounded by four other carbon atoms This arrangement is known as a tetrahedral arrangement.

 

 

Understanding carbon and its compounds Class 8

This is responsible for the rigidity, high density, and very high melting point of diamonds. Diamond is a very good conductor of heat. It has the highest thermal conductivity among the elements.

Diamond is a poor conductor of electricity. Chemically also it is quite inert. Its extreme hardness of it makes it suitable for making the tip of a boring drill that bores through rocks.

 

 

2. Graphite

The name is derived from the Greek word “Grapho” which means “I write”. It can put marks on paper. It is available in the free state of Sri Lanka, Mexico, and in the state of Rajasthan in India.

Graphite has a layered structure where two-dimensional sheets made of carbon are arranged parallel one over the other. The distance between two successive layers is relatively large and hence the force of attraction between the successive layers is weak. So when force is applied, one layer slides over the other.

 

WBBSE Chapter 4 summary on carbon

In each two-dimensional sheet, carbon atoms are arranged in a hexagonal planar fashion. Graphite is grayish in color. It is soft.

It is a good conductor of heat and electricity. That is why it is used as electrodes. A suspension of graphite in oil is used as a lubricant. Chemically, graphite is more reactive than diamond but less dense than diamond.

 

 

Comparison between diamond and graphite

Similarities between diamond and graphite:

 

Diamond	Graphite
1. Has crystalline structure.	1. Has crystalline structure.
2. Chemically very less active.	2. Chemically not very active.
3. Burns at high temperatures (800°C-850°C) in	3. Burns at high temperatures (700°C) in oxygen
oxygen to produce COr	to produce CO2.
4. Cannot absorb any gas.	4. Cannot absorb any gas.
5. Good conductor of heat	5. Good conductor of heat.

 

Dissimilarities between diamond and graphite:

 

Diamond	Graphite
1. Hardest natural element.	1. Soft and slippery element.
2. Colourless and transparent.	2. Blackish grey and opaque.
3. Non-conductor of electricity.	3. Good conductor of electricity.
4. Cannot put marks on the paper.	4. Can put a mark on the paper.

 

3. Fullerene

Fullerene – a new allotrope of carbon was first characterized in 1985 in the laboratory by Smalley and Kroto. This is a hollow, closed cage (polyhedral) cluster of 60 or 70 carbon atoms.

Its structure is based on polyhedra formed by fusing pentagons and hexagons-which is very much similar to geodesic domes used in architecture.

Fullerene is named after American architect R. Buckminster fullerene – the inventor of the geodesic dome. It has been subjected to extensive research since its discovery and is a promising candidate for use in electronics and medicines. It can be used to produce novel enclosure compounds by trapping metal ions within the C60cage.

 

Occurrence of carbon in the environment for Class 8

Amorphous allotropes of carbon

Various amorphous forms of carbon such as soot, carbon black, coke, coal, charcoal, etc. are all microcrystalline forms of graphite.

Coke can be used as fuel during metal extraction and as a reducing agent in redox reactions. Coal is primarily used as fuel.

Lamp black is used as a pigment and used for making printing ink for use in the printing press. Gas carbon is used to make electrodes for batteries, arc lamps, or electrolytic cells.

 

 

Charcoal has a remarkable property of adsorption. (Note that it is “adsorption” and not “absorption”). It can adsorb impurities from water and hence can purify water.

Specially prepared charcoal, known as activated charcoal, can adsorb large amounts of gas on its surface, so it is used to prepare gas masks.

Adsorption on charcoal: The adsorption property of charcoal can be easily shown by the following experiment. Let us take some amount of finely crushed charcoal powder.

Now dissolve some ink or some color in a small volume of water taken in a bottle with a lid. Now pour the charcoal powder in it and close the lid.

Then shake the bottle well for some time and then be allowed it to settle. Now the solution containing the charcoal dust is filtered.

It will be found that the intensity of the color of the solution after filtration has diminished considerably, indicating that a significant fraction of the molecules responsible for coloration has been “adsorbed” on the charcoal.

So, by filtration, when the solid charcoal powder is separated, the intensity of the color in the filtrate decreased.
Very recently another form of carbon- called graphene (consisting of planar sheets with atoms arranged in a honeycomb shape) has been discovered and looks promising for use in the field of nanotechnology.

 

All the allotropes of carbon contain the same element carbon

If samples of different allotropes of carbon are separately burnt in oxygen, carbon dioxide gas is produced in each case.

The gas can be tested by passing it into clear lime water that turns milky. This proves that each of the allotropes contains the same element carbon.

Again, burning equal weights of different allotropes of carbon in excess oxygen separately, the gas produced in each case is absorbed in a previously weighed tube containing caustic potash.

It will be observed that the increase in weight of each tube is the same, ie. an equal quantity of each allotrope produces the same amount of carbon dioxide gas.

 

Heating Value/Calorific Value of Fuel

Fuel includes all combustible substances that undergo oxidation or combine with oxygen from the atmosphere with the evolution of large amounts of heat capable of being economically applied to domestic or industrial purposes.

Examples of fuels are wood, coal, LPG, kerosene, diesel, petrol, etc. All these fuels contain carbon as the main constituent, so they are called carbonaceous fuels.

Fuels like hydrogen are not carbonaceous. In our everyday life, we use various types of fuels for different purposes.
The amount of heat released during the burning of different fuels is different.

 

 

Depending on the amount of heat released by a fixed mass of a particular fuel, it is used for specific purposes.

The amount of heat generated on the complete combustion of 1 kilogram of fuel with oxygen is called the heating value or calorific value of that particular fuel. It is generally expressed in units of kilocalorie per kg. or kilo Joule per kg.

Among all the fuels, hydrogen (150 kJ/g) has the highest heating value followed by LPG (liquefied petroleum gas). Wood has the lowest heating value among the above-mentioned fuels.

Classification of Fuels

Classification of fuels depends on several As per Physical state at normal pressure & temperature Primary Secondary Primary fuels are naturally occurring fuels in nature.

Examples: Wood, coal, petroleum, natural gas, etc. Secondary fuels are prepared or derived from primary fuels. Charcoal, semicoke, coke, kerosene, coke oven gas, etc are secondary fuels.

 

 

Fuels can be either vegetable (organic) or mineral (inorganic). Coal is an example of organic fuel while elemental sulfur, iron pyrites, etc are mineral or inorganic fuels.

Fuels can be classified into three classes on the basis of their physical state at normal temperature and pressure. They are solid, liquid, and gas. In the following table, different fuels are classified under these three categories and their uses are mentioned.

 

Physical State	Name of the Fuel	Calorific value	Uses
Solid	Wood	17 KJ/g	Cooking, for generating heat in colder countries
	Coal	(25-30)KJ/g	Cooking, generation of electricity, steam engines, preparation of bricks, etc.
Liquid	Kerosene	48KJ/g	Domestic purposes like cooking
	Petrol	50KJ/g	Automobile fuel
	Diesel	45KJ/g	Automobile fuel, diesel engines, industries
Gas	LPG (Liquefied Petroleum Gas)	50KJ/g	Domestic purposes such as cooking
	CNG (Compressed Natural Gas)	33-50 KJ/g	Automobile fuel

 

Conservation of Fuels

The advancement of modern civilization is dependent on the use of fuel. The fuels we mostly use nowadays are fossil – such as petroleum oil, coal, natural gas, etc.

Millions of years ago, prehistoric, plants and animals were buried under the earth probably due to severe earthquakes or another natural disaster of gigantic magnitude.

Their remains were decomposed under the high pressure of the earth in absence of air and were transformed into fossil fuels.
So, we can understand that the stock of fossil fuels available on the earth is limited.

But the irony is that the demand and use of fossil fuels are increasing every day. In the last century alone, we have burnt 90% of the total fossil fuel we have burnt during the total history of mankind.

If the demand for fossil fuels increases at such a rate, then all fossil fuels will be consumed probably by the end of the 21st century.

So we must have to think of ways to prolong the availability of fossil fuels and maintain a continuous and undisturbed supply Of energy for our various needs.

The strategies for this can be classified into two categories –

Judicious and optimized use of available fossil fuels Use of alternative sources of energy such as solar energy, wind energy, atomic energy, etc.

Judicious and optimized use of available fossil fuels: Since the stock of fossil fuels is limited and so far most of them have been consumed by human beings, hence, the rest of the fossil fuels still left have to be used rationally.

For example, coal is of two types – high-grade coal (having a higher calorific value) and low-grade co; (-having a lower calorific value).

The reserve of high-grade coal is smaller compared to the reserve of low-grade coal. High-grade coals should be used specifically for the generation of electricity, while low-grade coals should be used for domestic purposes such as cooking,

In coal mines areas, for example in Raniganj, Asansol, Jharia, etc. underground fires are slowly but steadily destroying our valuable coal reserves. This must be extinguished and care must be taken during mining so that no new fires break out.

Natural gas is available in various regions such as river basins etc. They have high calorific value and can be used for different purposes such as fuel for automobiles, generating electricity, etc.

But in many cases, natural gases are just burnt, thus a vital source of energy is simply wasted. The potential of natural gas as a natural source of energy has to be fully utilized.

The development of more efficient heat engines (which consumes less fuel and produces more work) can reduce the consumption of fossil fuel.

The development of energy-efficient electrical instruments and gadgets can ultimately reduce the consumption of electricity which in turn saves precious fossil fuel.

Examples of carbon compounds in nature Class 8

Use of public vehicles and carpooling can be encouraged instead of private vehicles.

The use of vehicles that do not require fossil fuel can be encouraged. For example, the use of bicycles on roads or non-mechanized boats in water can significantly reduce the consumption of fossil fuel.

In several countries, separate tracks for bicycles have been prepared to encourage the use of bicycles by the general public.

Use of alternative sources of energy: Fossil fuels are the conventional source of energy. But they can be substituted by alternative, non-conventional sorts of energy.

They include solar energy, wind energy, geothermal energy, tidal energy, bio-fuels, atomic energy, etc.

1. Solar energy

Solar energy is the most readily available source of energy. Solar energy has been used since prehistoric times for different purposes. But after 1970, when prices of petroleum soared, extensive research programs were initiated worldwide to exploit solar energy.

It can be used to generate electricity using solar panels on rooftops. Solar panels contain photovoltaic cells, which absorb the light energy of the sun and convert it to D.C. electricity directly.

This electricity can either be used immediately or can be stored in a battery for future use. This can be used for several applications such as domestic lighting, street lighting, rural electrification, water pumping, desalination of salty water, powering of remote telecommunication devices, etc.

India is one of the few countries with long daytime and plenty of sunshine. So in a country like India, the potential of solar energy can be best exploited.

2. Wind Energy

Wind energy is also an effective alternative to fossil fuels. This form of energy is free, plentiful, and renewable. Winds are caused by uneven heating of the atmosphere by the sun and the rotation of the earth.

This wind flow (or energy due to the motion of wind) can be converted to electricity by using wind turbines. A wind turbine is basically the opposite of a fan.

Wind turbines use the wind to make electricity. It has long blades. The wind turns the blades, which spin a shaft that is connected to a generator and electricity is produced.

During this operation, no greenhouse gas is emitted. So it is an “environmentally clean” alternative to fossil fuel.

 

3. Geothermal energy

Geothermal energy is the earth’s heat. Below the crust of the earth, the top layer of the mantle is a hot liquid rock called magma.

The crust of the earth floats on this liquid magma mantle. (When magma breaks through the surface of the earth in a volcano, it is called lava.)

Generally, for every 100 meters we go down, the temperature of the earth increases by about 3 degrees Celsius. So, if we go down to about 3000 meters below the ground, the temperature of the rocks there would be hot enough to boil water.

Deep under the surface, water sometimes makes its way close to hot rocks and turns into boiling hot water or into steam. When this hot water or steam comes upwards through a crack, we call it a hot spring.

There are many hot springs in India, but the potential of this energy source has not been exploited so far. The steam coming out from these hot springs can be directly used to rotate the turbine and thus produce electricity.

 

4. Biomass, Biogas, and Bio-fuel

1. Biomass is the term commonly used for the biological material derived from living or recently living organisms such as wood, waste materials, gases, and alcohol fuels.

In other words, it is dead material that was once living. It is a renewable energy resource derived from the carbonaceous waste of various human and natural activities.

It is derived from numerous sources, including by-products from the timber industry, agricultural crops, raw materials from the forest, major parts of household waste, and wood.

The composition of biomass is carbon, hydrogen, and oxygen. In addition, biomass energy is gaining significance as a source of clean heat for domestic heating and community heating applications.

 

2. Biogas:

Biomass can be converted to other usable forms of energy like biogas. This is generally prepared by the digestion of organic waste in absence of air.

Methane is produced from organic waste in this process. Organic waste includes municipal waste and drained water, agricultural waste like rotten vegetables, fruit pulp, feces of domestic animals, etc., human waste, left-over food products, etc.

In rural India, biogas technology is actively utilized. It is particularly useful for village households that have their own cattle.

The cattle dung is converted to biogas by a simple process and this biogas is used as fuel for cooking. So biogas plants are becoming quite popular in rural India.

 

3. Bio-fuel:

Biofuel is a fuel that is derived from biological materials, such as plants and animals. In more simple way we can say that bio-fuels are liquid and gaseous fuels produced from biomass.

So any hydrocarbon fuel that is produced from organic matter (living or once-living material) in a short period
of time (days, weeks, or even months) can be considered a biofuel.

Typical feed-stocks used for producing biofuel include sugarcane and sugar beet, starch-bearing grains like corn and wheat, oil crops like canola, soybean, and palm oil, and in some cases animal fats and used cooking oils.

Ethanol, bio-diesel, and methanol are the three most important examples of biofuels. Bioethanol is an alcohol made by fermentation, mostly from carbohydrates produced in sugar or starch-containing crops such as corn, sugarcane, etc.

Cellulosic biomass, derived from non-food sources, such as trees and grasses, is also being developed as a feedstock for ethanol production.

Ethanol can be used as a fuel for vehicles in its pure form, but it is usually used as a gasoline additive.
In India, bio-fuels are produced from oil obtained from the seeds of Jatropha plants.

Since Jatropha can be cultivated in less fertile and dry lands, so its cultivation is economically beneficial from the perspective of our country.

 

5. Tidal energy

Tidal energy is a form of hydropower that converts the energy of tides into more usable forms of energy, such as electricity. Tidal energy is produced by the surge of ocean waters during the rise and fall of tides. High tide and ebb always occur twice a day.

Unlike wind, tides are predictable and stable. Usually, turbines are placed in tidal streams (a fast-flowing body of water created by tides).

The kinetic energy of the tidal stream is utilized to rotate turbines which in turn produce electricity. Tidal energy is more powerful and effective than wind energy.

Since, our country has very long coastlines, so this form of renewable energy can be effectively utilized as an alternative form of energy.

 

6. Atomic energy

Neutrons and protons are bound within the nucleus of an atom by nuclear energy, if the nucleus is disintegrated, a huge amount of energy is released.

This energy can be utilized to generate electricity. Atomic energy can be generated by a process known as nuclear fission.

[Fission reaction was discovered by Otto Hahn and F. Strassmann in 1939. In this process nuclei of very heavy elements such as uranium, plutonium, or thorium are excited mainly by hitting them with a high-energy neutron.

As a result, the nucleus undergoes disintegration, producing two fragments along with a huge amount of energy. Otto Hahn was awarded Nobel Prize in 1944].

There are numerous electric power generating stations throughout the world, including India, which are based on nuclear fission technology.

In India, Dr. Homi J. Bhabha first envisioned a program for generating electricity utilizing atomic energy. This started with the setting up of Tarapur Atomic Power Station in 1969 in Maharashtra.

Later, several atomic power stations became operational, such as Rajasthan Atomic Power Station in Kota, Rajasthan (1973), at Kudankulam in Tamilnadu, Kalpakkam in Kerala (1985) and Kakrapar in Gujarat (1991), etc.

It is estimated that by the end of the year 2032, approximately 63,000 megawatts of electricity will be generated using atomic energy.

Carbon cycle and its compounds for Class 8

In the developed world, such as France nearly 70% of its electricity is generated by utilizing atomic energy. Nearly 40% of the total electricity generated is produced from atomic energy in countries like USA, UK, Germany and Japan.

The vast deposits of thorium available in India can be exploited to generate atomic energy.

But keeping in mind the incidents that happened in Chornobyl in Russia and Fukushima in Japan, proper care and round-the-clock maintenance is required in all atomic energy power plants.

Nuclear waste management is also a big issue that must be resolved since the fragments derived from the disintegration of the heavy nucleus are radioactive and so hazardous. Generally, the radioactive nuclear wastes are buried underground in sealed, thick-walled steel drums.

 

7. Solid waste

In many metropolitan cities, solid wastes are being utilized to produce electricity. Solid wastes are burnt and the heat generated is used to produce steam, which ultimately rotates a turbine to produce electricity.

But, solid wastes available in metropolitan cities contain a large variety of materials, and some of them can cause serious environmental pollution during their burning.

 

Hazardous effects on the environment due to the combustion of fossil fuel

1. Fossil fuel serves as the most important source of energy in today’s world. But the use of fossil fuels also is the cause of environmental pollution.

During incomplete combustion of carbonaceous fuels like wood, coal, and petroleum, unburnt carbon particles are produced which float in the air.

1. When inhaled, these suspended particles may deposit within the nose, throat, and respiratory tracts. This may cause respiratory problems.

2. The incomplete combustion of carbonaceous fossil fuel produces carbon monoxide which is very poisonous.

3. Complete combustion of carbonaceous fuel produces carbon dioxide (CO₂). The increasing concentration of CO₂ in the atmosphere is mainly responsible for global warming.

4. During the combustion of coal and diesel, sulfur dioxide (S02) is produced. During combustion at elevated temperatures, the nitrogen in the air is converted to oxides of nitrogen (commonly called NOx).

The oxides of carbon, sulfur, and nitrogen are acidic in nature and react with rainwater to produce acids such as sulphuric acid (H2S04), nitrous acid (HN02), nitric acid (HN03), and carbonic acid (H2C03).

These acids, dissolved in rainwater cause acid rain. Acid rain is harmful to crops, soil, and buildings. It can also affect aquatic ecology.

5. Excessive use of fossil fuels causes the release of greenhouse gases (eg. CO₂, N0x) that results in global warming.
Compared to these fossil fuels like coal, coke, petroleum products, etc.,

CNG and LPG are cleaner fuels. Complete combustion of LPG and CNG is relatively easy and their efficient combustion produces very low carbon particles (or soot).

 

Greenhouse Effect

Principle of action of greenhouse: A greenhouse is a house made of glass. It is used for raising plants of vegetables, fruits, and flowers in cold temperate regions where the growing of plants in the open fields becomes impossible due to severe cold.

During the daytime, sunlight enters the room through its glass walls and roof and warms the plants and air present inside the room.

Sunlight reaches earth through invisible radiation, called infrared radiation (IR radiation), which is primarily responsible for the sensation of heat.

But the heat generated by sunlight cannot be stored for an indefinite period and hence after some time the absorbed heat is radiated.

The radiated heat or emitted heat is less energetic than the original heat of the sunray which initially entered the glasshouse.

This emitted radiation cannot pass through the glass walls and roof. So some amount of heat is trapped inside the glass house and cannot escape. So during the daytime, the inside of the house gets warmer and stays quite warm during the night too.

Greenhouse gases and greenhouse effect: Gases present in the earth’s atmosphere such as carbon dioxide, water vapor, methane, nitrous oxide, ozone and some other man-made chemicals (such as chlorofluorocarbons) do exactly what the glass roof and walls do in a greenhouse.

They are known as greenhouse gases. The molecules of these gases cannot absorb solar radiation directly coming from the sun. Hence, during the daytime, the earth’s surface warms up.

At night it cools down by radiating the heat back to space. But a part of this low-energy infrared radiation, radiated by the earth, is absorbed by molecules of greenhouse gases.

Thus heat is trapped by the greenhouse gases in the atmosphere and is redirected again back to earth. This mechanism keeps the earth warm.

This is a natural phenomenon known as the greenhouse effect- So the earth’s surface remains warm and cozy which is essential to sustain life on earth.

 

Importance of carbon in nature for Class 8

But if the greenhouse effect is high, the earth’s surface gets warmer than what is required. This is known as the enhanced greenhouse effect or more commonly just Greenhouse Effect.

1. The presence of greenhouse gases in much greater concentration than the normal level is responsible for this enhanced greenhouse effect causing global warming.
2. The steps which are occurring during the enhanced greenhouse effect may be approximately summarized below:
3. Solar radiation reaches the earth
4. Sun’s energy is absorbed by the land and the oceans, and the earth’s surface becomes warm
5. During the night, heat is radiated by the earth back to space.
6. Some of the heat is trapped by the greenhouse gases present in the atmosphere, keeping the earth warm enough to sustain life.
7. Human activities such as burning fossil fuels are increasing the concentration of greenhouse gases in the atmosphere rapidly.
8. Trapping of the extra amount of heat causes the earth’s temperature to rise above normal and this leads to the phenomenon known as the enhanced greenhouse effect or more commonly greenhouse effect.

 

Sources Of Green House Gases:

Human activities, particularly the burning of fossil fuels (such as coal, oil, natural gas, etc.) and heating limestone in cement factories are increasing the concentration of greenhouse gases such as CO₂ rapidly.

Some bacteria known as denitrifying bacteria found in soil reduce the nitrate in several steps and during the process produce greenhouse gases such as nitrous oxide.

Methane (CH4) is another greenhouse gas that has about 20 times more greenhouse effect than CO₂. Methanogenic bacteria present in wetlands and rainforests produce methane.

Bacteria present in the rumen of cattle and in the intestine of termites also produce methane.

4. Use of chlorofluorocarbons (CFCs) in refrigeration and air conditioning systems and the use of CFCs and Halons in fire extinguishing systems contribute to rising in greenhouse gas concentrations.

The major concern is that the concentration of many greenhouse gases is increasing at an alarming level. The volume of CO₂ in the atmosphere has increased by over 35% in the last three hundred years.

Generally, the earth has its own mechanism for maintaining the level of CO₂ in the atmosphere. But if the level of such gases increases rapidly, this mechanism is bound to fail and this will increase the temperature of the earth. This will lead to Global Warming.

Result of global warming: As a direct consequence, the polar ice caps (in which more than 90% of the earth’s total drinking water is stored) will melt.

So the water level in the oceans will rise and several coastal cities (having a habitat of millions of people) will submerge. The ecosystem will be adversely affected.

Increasing temperature will enhance the growth of mosquitoes which in turn, will spread mosquito-borne diseases. Biodiversity will be hampered and several parts of the world may face drought-like conditions for a prolonged time. In short, the effect of the greenhouse effect (or enhanced greenhouse effect) will bring devastation.

Carbon-Containing Polymers and Their Uses

The modern age is often termed as plastic age. If we look around we can see a wide range of materials starting from rubbers, plastics, paints and surface coatings, resins, adhesives, etc.

Some of them are naturally obtained, and some of them are artificial (i.e. manufactured in a laboratory or industry). They all belong to a particular class of compounds known as polymers.

“Poly” means many and “meros” means parts. Polymers are high molecular weight compounds that are made up of a large number of (few hundred to few thousand) simple, repeating units called monomers.

In today’s world, we use a variety of carbon-containing polymers for a variety of purposes. Polythene, thermocol, nylon, etc. are all carbon-containing polymers.

The large size and their shapes are the two most important properties which are utilized for different purposes. Some of the polymers are flexible and some are hard.

Some are heat resistant and very good insulators of electricity. Some polymers have a high melting point and some are chemically quite inert.

Depending upon the need, one can choose a polymer that suits best a specific purpose. Plastics are sometimes used by common people to indicate all types of polymers.

But actually, polymer and plastics are not at all synonymous. Plastics are a special type of polymer.

 

 

Let us discuss briefly some common carbon-containing polymers, such as polyethylene, Teflon, PVC, nylon, and terylene.

1. Polythene (or Polyethylene)

Polythenes generally show excellent chemical resistance, meaning that it is not attacked by strong acids or strong bases.

It is also resistant to mild oxidizing and reducing agents. Polyethylene burns slowly with a blue flame. Crystalline samples are not soluble at room temperature but are usually soluble at higher temperatures in toluene, xylene, and chlorinated solvents. It is an insulator of electricity.

Uses of polythene: Polythene is widely used for food packaging, milk carton lining, and making shopping bags. It is used to make squeezable bottles, drums, and containers for different purposes.

High-density polyethylene is used to make dustbins, crates, buckets and bowls, food boxes, etc. It is also used to coat paper, which is then rendered a waterproof surface.

Pipes and hoses, flexible water pipes, etc. are made of polyethylene. Since it is an insulator of electricity it is used for coating electrical cables.

2. Teflon

Teflon is a high-molecular-weight compound consisting wholly of carbon and fluorine. It has a high melting point, very good insulator of electricity, and is chemically quite inert (even towards strong acids like aqua regia).

This is because of the very strong bonds between carbon and fluorine. It is a hydrophobic material (i.e. water cannot wet it). It experiences a very low frictional force against a solid surface.

Uses of Teflon: Most of us know Teflon for its non-stick properties in cookware applications. Teflon coatings are applied on a range of cookware as it is hydrophobic and possesses fairly high heat resistance.

Besides Teflon is widely used for wiring in aerospace and computer applications. It is used for making cables and connector assemblies.

In industrial applications, owing to its low friction, it is used for applications where the sliding action of parts is needed.

PVC (polyvinyl chloride)

PVC is the abbreviation of polyvinyl chloride. It consists of carbon and chlorine. PVC was accidentally synthesized in 1872 by German chemist Eugen Baumann.

The polymer appeared as a white solid inside a flask of vinyl chloride that had been left exposed to sunlight. Industrially, PVC is prepared from petroleum.

PVC is a relatively low-cost material. It has high hardness and mechanical properties. It is available both as transparent and opaque substances.

It is chemically and biologically quite inert and it is inert towards different solvents such as water, oil, petrol, and other chemicals. So, it is biocompatible.

The most important characteristic of PVC is its high fire resistance. As a result, it has found application for a wide variety of substances.

Uses of PVC: Since PVC is highly fire resistant, it is widely used in exterior construction materials such as window profiles, siding boards, or interior housing materials, such as wall-covering and flooring.

Fire-resistance property of PVC is used widely in a variety of applications such as electric cables for residential buildings, vehicles, household electrical appliances, cable coverings, insulating tapes, switch boxes, wire coverings, and protecting tubes for power and telecommunications cables.

It is used for sewerage pipes and other pipe applications where cost or vulnerability to corrosion limits the use of metal.

PVC is widely used in clothing, to either create a leather-like material or at times simply for the effect of PVC. PVC is used for preparing flexible containers and tubing – containers used for blood and blood components, for urine collection and tubing used for blood taking and blood giving sets, catheters, heart-lung bypass sets, hemodialysis sets, etc.

4. Nylon

Nylon is a polymer and is frequently referred to as polyamide. Nylon is a silky material. It is transparent. Strong, water-resistant fibers can be made from this material.

It can sustain high temperatures. Nylon was first used commercially in a nylon-bristled toothbrush (1938), and within the next couple of years, it was commercially introduced as a fabric.

In 1939, World War 2 started. During World War 2, when Asian silk and different products made of silk became scarce, extensive research was carried out to develop Nylon as a synthetic replacement for silk.

Uses of nylon: Soon it replaced silk in military applications such as parachutes and flak vests, and was used in many types of vehicle tires.

It was also used to make tents, ropes, and other military supplies. Nylon is being used for those purposes still today. Nylon fibers are used by the carpet manufacturing industry.

They are also used for making fishing lines and musical strings. Nylon resins are used for making food packaging films. Nylon has been used for meat and sausage wrappings.

5. Terylene

Polyester is a polymer where the individual units are held together by “ester linkage”. When it is used as a fiber to make clothes, it is then sometimes known as “terylene”.

When it is used to make bottles, it is usually called PET (polyethylene terephthalate). This material is soft and flexible, resistant to stains, and does not absorb water. It can produce long and strong fibers, which are resistant to folding and has a long life.

Uses of terylene: So it is widely used as fibers to make clothes. Terylene can be used together with cotton fibers to make a material known as tricot (terylene + cotton).

As a result, fibers made from tricot possess characteristic properties of both terylene and cotton.

Crisis Of Artificial Polymer

Man-made polymers or artificial polymers are widely used for various purposes. The main advantage of using them is that they are durable i.e. have a long life.

But the same durability properties which make plastics ideal for so many applications can lead to waste disposal problems because these materials are not readily biodegradable.

By the word non-biodegradable, it means that they do not undergo degradation by living organisms such as microbes, bacteria, fungi, etc. Because of their resistance to biodegradation, they accumulate in the environment.

In urban areas, it can lead to several problems such as blocking the drains and sewer lines. Poly bags, plastic pouches, etc.

made of such artificial, non-biodegradable polymers lying scattered here and there has become a very common scenario in almost every city and village of our country.

Efforts to tide over the crisis: Researchers are trying to develop biodegradable polymers to suit our needs.

Biodegradable polymers are a specific type of polymer that is decomposed naturally by environmental microorganisms after their intended use.

Cellulose is one such polymeric carbohydrate that is biodegradable. It is the main constituent of the fiber of cotton and straw.

These polymeric carbohydrates are decomposed in nature by microorganisms such as fungi and bacteria.

As a result, large polymer molecules are degraded to smaller molecules which are required for the growth of those micro-organisms which degrade the polymer.

Proteins are also polymers or more appropriately, macromolecules. They are present in pieces of meat and fish. They are degraded by some specific enzymes produced by some particular bacteria.

Cellulose-based polymers, chitosan-based polymers, and starch-based polymers are all examples of biodegradable polymers. Efforts are going on to prepare biodegradable polymers with desired chemical and mechanical properties, which are non-toxic and can be naturally decomposed.

Chapter 3 Some Common Gases Long Answer Questions

Question 1. Describe burette and pipettes. Why pipettes and burettes are used in the laboratory?
Answer:

Filter Paper

Filter paper is used to separate solid, insoluble particles from a liquid. It is a thick, porous, circular paper. Filter papers having different pore sizes are available.

The filter paper is chosen based on the size of the insoluble, solid particles which are to be separated from the liquid.

After getting some idea about some common laboratory equipment and apparatus, we can now discuss two common gases which are very come to know about their physical and chemical usage.

One of them is oxygen and the other is Properties, their sources, and their preparation of hydrogen. During this brief discussion, we will procedures and their uses.

Read And Learn More WBBSE Solutions For Class 8 School Science Long Answer Type Questions

Question 2. Discuss briefly the laboratory method of preparation of oxygen from potassium chlorate.
Answer:

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.

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

1. The gas is collected by the downward displacement of water, because,
2. The solubility of oxygen in water is low

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 small amount of MnO₂, KCIO₃ decomposes at about 250°C to produce oxygen.

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.

 

West Bengal Class 8 Science Gases Solutions

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 lower temperatures with the help of a catalyst.

Precaution

1. KCIO₃ 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.

Question 3. Discuss a laboratory method for the preparation of oxygen at room temperature.
Answer:

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.

Common Gases Chapter 3 WBBSE

Question 4. How oxygen can be produced by the electrolysis of water?
Answer:

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.

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.

WBBSE Solutions For Class 8 School Science Long Answer Type Questions	WBBSE Solutions For Class 8 School Science Short Answer Type Questions
WBBSE Solutions For Class 8 School Science Very Short Answer Type Questions	WBBSE Solutions For Class 8 School Science Review Questions
WBBSE Solutions For Class 8 School Science Solved Numerical Problems	WBBSE Solutions For Class 8 School Science Experiments Questions
WBBSE Solutions For Class 8 Maths	WBBSE Class 8 History Notes
WBBSE Class 8 History Multiple Choice Questions	WBBSE Solutions For Class 8 History
WBBSE Solutions For Class 8 Geography

 

Question 5. Discuss briefly the physical properties of oxygen.
Answer:

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. The freezing point of liquid oxygen is – 218°C and the boiling point of liquid oxygen is – 183°C.
4. 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.
5. 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	0.99763%	0.00037%	0.002%

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

Question 6. Discuss briefly the reaction of oxygen with non-metals.
Answer:

The 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.

When a piece of glowing charcoal 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 small quantity of sulphur is heated in a flame and 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.

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

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

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

WBBSE Class 8 Gases Practice Questions

Question 7. Oxides of non-metals are generally acidic—Why? Generally, most oxides of non-metals are acidic. Their aqueous solution produces acid.
Answer:

Oxides of non-metals are generally acidic:

For example, CO₂ SO₂ and P₂O₅ are three oxides of non-metal. When they are dissolved in water, they form acids (e.g. carbonic acid, sulphurous acid and phosphoric acid, respectively) and turn blue litmus paper red, indicating that their solution is acidic.

⇒ \(\begin{array}{ll}
\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2 ; & \mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{CO}_3 \\
\mathrm{~S}+\mathrm{O}_2 \rightarrow \mathrm{SO}_2 ; & \mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{SO}_3 \\
4 \mathrm{P}+5 \mathrm{O}_2 \rightarrow 2 \mathrm{P}_2 \mathrm{O}_5 ; & \mathrm{P}_2 \mathrm{O}_5+3 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{H}_3 \mathrm{PO}_4
\end{array}\)

Question 8. Discuss briefly the reaction of oxygen with metals.
Answer:

The reaction of oxygen with metals:

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

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).

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

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).

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

In a similar way, potassium, calcium, etc. react with oxygen to form basic oxides

\(\begin{gathered}
2 \mathrm{Ca}+\mathrm{O}_2 \rightarrow 2 \mathrm{CaO} \\
4 \mathrm{~K}+\mathrm{O}_2 \rightarrow 2 \mathrm{~K}_2 \mathrm{O}
\end{gathered}\)

Question 9. Oxides of metals are generally basic—Explain.
Answer:

Oxides of metals are generally basic:

Metal oxides are mostly basic oxides. Some basic oxides form hydroxides when they react with water. Hydroxides of some metals such as sodium, magnesium, calcium, potassium etc.

Are soluble in water and turn red litmus paper blue, indicating that the aqueous solution of those metal oxides is basic. Some examples are given below.

⇒ \(\begin{array}{ll}
\mathrm{C}+\mathrm{O}_2 \rightarrow \mathrm{CO}_2 ; & \mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{CO}_3 \\
\mathrm{~S}+\mathrm{O}_2 \rightarrow \mathrm{SO}_2 ; & \mathrm{SO}_2+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{H}_2 \mathrm{SO}_3 \\
4 \mathrm{P}+5 \mathrm{O}_2 \rightarrow 2 \mathrm{P}_2 \mathrm{O}_5 ; & \mathrm{P}_2 \mathrm{O}_5+3 \mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{H}_3 \mathrm{PO}_4
\end{array}\)

Question 10. Give two examples of amphoteric oxides. Why are they so-called? Explain.
Answer:

Examples of amphoteric oxides:

Aluminium oxide (Al₂O₃) and zinc oxide (ZnO) are two examples of 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.

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

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.

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

West Bengal Board Class 8 Science Revision

Question 11. Stannic oxide and lead monoxide are amphoteric oxides—Explain.
Answer:

Stannic oxide and lead monoxide are amphoteric oxides:

Stannic oxide and lead monoxide are called amphoteric oxide because they have properties of the acidic and basic oxide.

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.

Stannic oxide (SnO₂) is prepared by burning tin at white heat in the air. 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: \(\mathrm{SnO}_2+2 \mathrm{NaOH} \rightarrow \mathrm{Na}_2 \mathrm{SnO}_3+\mathrm{H}_2 \mathrm{O}\)

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

Question 12. Discuss briefly the uses of oxygen in industries.
Answer:

The uses of oxygen in industries:

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 reaction with oxygen which is then converted to HSO₂ 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.

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.

Question 13. Show with an experiment that hydrogen is lighter than air.
Answer:

An experiment that hydrogen is lighter than air:

With the help of an experiment, it can be shown that hydrogen is lighter than air.

Apparatus and chemicals required: Two gas jars with lids – one filled with hydrogen gas and the other filled with air, a taper.

Experiment: 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.

Observation: A “pop” sound Is heard. The taper extinguishes but the gas burns with a bluish flame.

Inference: This confirms that the gas in gas Jar B is hydrogen. Hydrogen was initially in the gas Jar A.

When the gas jar filled with air is Inverted over the gas jar filled with hydrogen, hydrogen is lighter than air, moves upwards and is collected in gas jar B by downward displacement of air.

Chapter 3 Common Gases Important Questions

Question 14. Discuss briefly the reaction of hydrogen with the following non-metals: chlorine, nitrogen and sulphur.
Answer:

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.

Class 8 Science Gases Study Material

Question 15. What precautions should be taken during the laboratory preparation of oxygen from potassium chlorate?
Answer:

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,

The solubility of oxygen in water is low

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.

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.

 

 

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 lower temperatures with the help of a catalyst.

Precaution

1. KCl03 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.

WBBSE Class 8 Environment and Science Notes

Question 16. A soft white metal A reacts with water to form a compound B and a colourless gas C. When C is passed over heated copper oxide, water and a red-brown coloured element D are formed. Identify A, B, C, and D and write the reactions.
Answer:

Given:

A soft white metal A reacts with water to form a compound B and a colourless gas C.

When C is passed over heated copper oxide, water and a red-brown coloured element D are formed.

A is sodium. Sodium reacts with water to form sodium oxide (Na₂O:B) and hydrogen (H₂:C) 2Na + H₂O = Na₂O + H₂O , When hydrogen is passed over heated cupric oxide (CuO), water and metallic copper (Cu :D) are formed.

⇒ \(\mathrm{CuO}+\mathrm{H}_2=\mathrm{Cu}+\mathrm{H}_2 \mathrm{O} \text {. }\)

WBBSE Chapter 3 Some Common Gases Short Answer Type Questions

Question 1. Why the use of LED is advantageous for laboratory use over common bulbs?
Answer:

LED is advantageous over common bulbs because:

1. it has a long life;

2. during repeated use they do not get fused easily and they can glow even when a very small current is flowing through the electric circuit. So it can detect very small currents flowing through an electrical circuit.

Question 2. How oxygen was formed on Earth?
Answer:

Oxygen was formed on Earth:

The earth was created approximately 4500 million years ago. The atmosphere of the earth is composed of main hydrogen, ammonia, nitrogen, carbon dioxide, hydrogen sulphide and methane.

The abundance of oxygen was very very low. About 2500 million years ago, a type of bacteria, called cyanobacteria emerged in the sea.

With the help of sunlight and special protein, they started splitting water to form oxygen. As a result the amount of oxygen in the atmosphere gradually started to rise and approximately 500 million years ago, the amount of oxygen attained the present level.

Question 3. Why the presence of oxygen in the atmosphere is a must for us?
Answer:

The presence of oxygen in the atmosphere is a must for us:

The presence of oxygen in the atmosphere is a must for us. Because it is essential for respiration. Respiration is a process by which energy is liberated from intracellular glucose in a very efficient way.

Read And Learn More WBBSE Solutions For Class 8 School Science Short Answer Type Questions

This liberated 20 energy is used for various biochemical and biophysical processes continuously occurring in our body so that we can remain alive.

Question 4. Do all organisms on the earth require oxygen? No, not all organisms require oxygen for their survival. Answer: There are some areas in the earth where oxygen cannot penetrate.
Answer:

For example, in the depth of marshes or into the sludge of urban sewers, oxygen is totally absent. Some bacteria still live there, which are known as obligate anaerobes.

They produce energy within their bodies by using a different mechanism, which does not require the involvement of oxygen.

Question 5. Is the presence of oxygen always beneficial for us?
Answer:

Energy is liberated from glucose in presence of oxygen in a very efficient way within our bodies.

But, along with the generation of energy vital for maintaining all our biophysical and biochemical processes, it forms some harmful chemicals such as hydrogen peroxide (H₂O₂) and superoxide ions (O₂).

They can cause great damage to cellular DNA even when they are present in small amounts.

WBBSE Class 8 Science Short Answer Questions

Question 6. How hydrogen peroxide and superoxide ions formed in our body are removed?
Answer:

Hydrogen peroxide (H₂O₂) and superoxide ions (O₂) are very harmful to our bodies.

They can cause great damage to cellular DNA even when they are present in small amounts. Certain enzymes present within the living cells of our body can destroy these species.

For example, the enzyme, catalase, can break down hydrogen peroxide into oxygen and water.

Question 7. Why manganese dioxide is used during the preparation of oxygen from potassium chlorate?
Answer:

During the preparation of oxygen from potassium chlorate, MnO₂ acts as a positive catalyst. If KCIO₃ is heated alone, oxygen is produced at a temperature higher than 610°C.

In presence , of little amount of \(\mathrm{MnO}_2, \mathrm{KClO}_3\) decomposes at about 250°C to produce oxygen.

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

Question 8. “Oxygen is chemically very reactive”—Briefly explain.
Answer:

“Oxygen is chemically very reactive”:

Oxygen is chemically very reactive and forms compounds with practically all elements. except for 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 presence of a suitable catalyst. The compound produced in this reaction is called oxide.

Question 9. Discuss briefly the use of oxygen for medical purposes.
Answer:

The use of oxygen for medical purposes:

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).

For patients who are rendered unconscious due to inhalation of poisonous gas, the inhalation of carbogen (which is a mixture of 95% oxygen and 5% COJ is prescribed. A mixture of oxygen and nitrous oxide is used for anaesthesia.

Class 8 Science Short Answer Format

Question 10. Briefly discuss the importance of oxygen in respiration.
Answer:

The importance of oxygen in respiration:

Respiration is a process by which complex foodstuffs, 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 food kinds of stuff are converted into CO₂. For example, during respiration in our body, glucose is oxidized to CO₂ and H₂0.

Energy is liberated during this process which is utilized in different physicochemical processes occurring in our bodies.

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

Question 11. When a balloon filled with hydrogen gas is released, it moves upwards—Why?
Answer:

When a balloon filled with hydrogen gas is released, it moves upwards:

Hydrogen is the lightest gas and air is nearly 14.4 times heavier than hydrogen. When a balloon filled with hydrogen gas is 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.

Question 12. Discuss briefly the “adsorption” of hydrogen.
Answer:

“Adsorption” of hydrogen:

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.

Question 13. Why pure zinc is not used in the preparation of hydrogen gas?
Answer:

This is because pure zinc reacts very slowly with dilute sulphuric acid or dilute hydrochloric acid. Thus the production of hydrogen gas will be extremely small in quantity.

[In fact, the reaction of zinc with dilute acids (such as HCI or H₂SO₄) is an electrochemical process. The impurities like iron, lead, etc.

which are commonly present in commercial zinc forms tiny, local electrochemical cells. The impurities act as the positive pole of a tiny electrochemical cell and zinc acts as a negative pole and the dilute acid acts as the electrolyte.

Due to the formation of so many tiny, local, electrochemical cells, the rate of production of hydrogen increases

WBBSE Science Chapter 3 Question Answers

Question 14. Why concentrated H₂S04 is not used in the preparation of hydrogen gas?
Answer:

This is because the acid produces sulphur dioxide gas during the reaction with zinc.

Also, when the acid is decomposed, it produces atomic oxygen (or nascent oxygen) \(\left[\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{H}_2 \mathrm{O}+\mathrm{SO}_2+\mathrm{O}\right]\) which combines with hydrogen (produced in the process) to form water.

Question 15. Why concentrated hydrochloric acid is not used in the preparation of hydrogen gas?
Answer:

This is because concentrated HCI is highly volatile and hence plenty of HCI vapour mixes with the hydrogen gas produced during the process.

Moreover, zinc reacting with HCI produces insoluble ZnCI2 which forms a coating on the metallic zinc and thus stops the further reaction.

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Question 16. Why strong or moderately strong nitric acid is not used in the preparation of hydrogen gas?
Answer:

The acid decomposes even at room temperature to produce oxygen [4HNO₃→4NO₂ + 2H_2O+ O₂. This oxygen will thus combine with the hydrogen gas produced in the process to form water.

Question 17. Show that occluded hydrogen is a more powerful reducing agent than ordinary hydrogen.
Answer:

Ferric chloride is not reduced when hydrogen is bubbled through an aqueous solution (yellow-coloured) of the salt.

But, if a small quantity of spongy palladium, containing occluded hydrogen, is dipped into a solution of ferric chloride and heated,

it is immediately reduced to ferrous chloride as is shown by the change of the yellow colour of the solution into almost colourless (or light green).

FeCl₃+H₂(molecularhydrogen)→NoreactionFeCl₃+[H](occludedhydrogen)→FeCl₂+HCl.(yellow)(Colurless)

 

Chapter 3 Some Common Gases VSAQs

Question 1. According to scientists, when the earth was created?
Answer:

According to scientists, the earth was created about 4500 million years ago.

Question 2. According to scientists, when life emerge on earth?
Answer:

According to scientists, life emerged on earth about 3500 million years ago.

Question 3. Which instrument is used in the laboratory to measure the temperature of an object? The thermometer is used in the Answer: laboratory to measure the temperature of an object.

Question 4. Where can you find cyanobacteria?
Answer:

Cyanobacteria can be found in wetlands, ponds and paddy fields.

Question 5. Name the polymer with which the copper wire is enveloped.
Answer:

Copper wire is enveloped by a polymer named PVC (polyvinyl chloride).

Read And Learn More WBBSE Solutions For Class 8 School Science Very Short Answer Type Questions

Question 6. What do you mean by obligate anaerobes?
Answer:

Obligate Anaerobes:

There are some bacteria which live in the depth of marshes or sludge in urban sewers. Here oxygen is not at all available. These bacteria can survive without oxygen and die in presence of oxygen. Such organisms are called obligate anaerobes.

Question 7. What is carbogen?
Answer:

Carbogen:

Carbogen is a mixture of 95% oxygen and 5% COr Inhalation of carbogen is prescribed for patients who are rendered unconscious due to inhalation of a

Question 8. What are the three naturally occurring isotopes of oxygen?
Answer:

Oxygen has three naturally occurring isotopes, 16eO, ”0 and. (But the natural abundance of the last two is very low).

WBBSE Class 8 Very Short Answer Questions

Question 9. Which gas of air is depleted due to combustion?
Answer:

Oxygen.

Question 10. Name a binary compound of oxygen which is not an oxide.
Answer:

FzO (fluoride of oxygen)

Question 11. What is an oxide?
Answer:

Oxide:

Actually oxide is a compound of two elements, one of which is oxygen.

Question 12. Give three examples of acidic oxide.
Answer:

Examples of acidic oxide:

Carbon dioxide (CO₂), sulphur dioxide (SO₂) and phosphorous pentoxide (P₂0₅) are examples of three acidic oxides.

Question 13. Name two amphoteric oxides.
Answer:

Aluminium oxide (Al203) and zinc oxide (ZnO) are two amphoteric oxides.

Question 14. Name one non-metallic peroxide and one metallic peroxide.
Answer:

Hydrogen peroxide (H₂O₂) is a non-metallic peroxide and sodium peroxide (Na₂O₂) is a metallic peroxide.

Very Short Answer Solutions for Class 8 Science

Question 15. Name a few metals which can adsorb oxygen.
Answer:

At normal or low pressure, certain metals like gold (Au), silver (Ag), platinum (Pt) and palladium (Pd) absorb oxygen.

Question 16. Which is the lightest element?
Answer:

Hydrogen is the lightest element.

Question 17. What are the melting point and boiling point of hydrogen?
Answer:

The melting point of hydrogen is – 259.2 °C and the boiling point of hydrogen is – 252.6°C.

Question 18. What are the three isotopes of hydrogen?
Answer:

Isotopes of hydrogen:

Hydrogen has three isotopes: and (The natural abundance of the last two is negligibly small)

WBBSE Very Short Answer Format Class 8

Question 19. Name a few metals which can adsorb hydrogen at normal temperature.
Answer:

Some metals like palladium (Pd), platinum (Pt), iron (Fe), nickel (Ni), etc. can adsorb hydrogen at normal temperature

Question 20. What do you mean by “hydride”?
Answer:

Hydride:

Binary compounds of elements with hydrogen are called hydrides.

Question 21. Name one oxide compound which can be heated to produce hydrogen.
Answer:

Mercuric oxide (HgO), when heated, produces hydrogen.

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Question 22. Name one compound which can be heated to produce oxygen.
Answer:

Oxygen is obtained from hydrogen peroxide (H₂O₂) at normal temperature.

Question 23. Name one metallic hydride compound.
Answer:

Lithium hydride (LiH)is an example of a metallic hydride compound.

Question 24. Which of the two gases hydrogen and oxygen is soluble in water
Answer:

Oxygen.

WBBSE Chapter 3 Some Common Gases MCQs

Question 1. The following is used as a catalyst during the preparation of oxygen from potassium chlorate

1. Phosphorous pentoxide
2. Manganese dioxide
3. Potassium permanganate
4. Vanadium pentoxide

Answer: 2. Manganese dioxide

Question 2. Hydrogen acts as an oxidising agent when it

1. Burns in presence of O₂
2. Forms ammonia when it reacts with nitrogen
3. Passes through boiling sulphur and forms [h₂s]
4. Reacts with metals to form metallic hydrides.

Answer: 2. Forms ammonia when it reacts with nitrogen

Read And Learn More WBBSE Solutions For Class 8 School Science Review Questions

Question 3. The gas which is used for the hardening of oil and fat is

1. Hydrogen
2. Oxygen
3. CO₂
4. Producer gas

Answer: 1. Hydrogen

Question 4. Na₂0₂ is

1. A basic oxide
2. A peroxide
3. An acidic oxide
4. An amphoteric oxide

Answer: 2. A peroxide

WBBSE Class 8 Science Chapter 3 Review Questions

Question 5. Hydrogen is

1. Completely insoluble in water
2. Sparingly soluble in water
3. Completely soluble in water
4. None of these

Answer: 2. Sparingly soluble in water

Question 6. Air is heavier than hydrogen by

1. 1.44 times
2. 14.4 times
3. 144 Times
4. Is of equal weight

Answer: 2. 14.4 times

Question 7. O₂ And h₂ prepared in the laboratory can be collected by

1. Downward displacement of air
2. Upward displacement of water
3. Upward displacement of air
4. Downward displacement of water

Answer: 4. Downward displacement of water

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Question 8. An example of an acidic oxide is

1. Co
2. MgO
3. SO₂
4. Zno

Answer: 3. SO₂

Question 9. Hydrogen is adsorbed on

1. Palladium
2. Sulphur
3. Iron
4. Zinc

Answer: 1. Palladium

WBBSE Science Chapter 3 Important Questions

Question 10. Oxygen is produced at room temperature due to a reaction between

1. Sodium peroxide and water
2. Potassium chlorate and MnO₂
3. Cupric oxide and hydrogen
4. Potassium chlorate with water

Answer: 1. Sodium peroxide and water

Question 11. Oxygen gas

1. Is very heavy compared to the air
2. Slightly heavier than air
3. Lighter than air
4. Has equal weight to air

Answer: 2. Slightly heavier than air

Question 12. Oxygen is produced during

1. Respiration
2. Photosynthesis
3. Combustion
4. Fermentation

Answer: 2. Photosynthesis

Question 13. A balloon is filled up with hydrogen gas.

1. It will lie on the floor
2. It will go upwards and will touch the roof of the room in which the experiment is carried out.
3. It will float in the air
4. It will catch fire

Answer: 2. It will go upwards and will touch the roof of the room in which the experiment is carried out.

Question 14. The colour of liquid oxygen is

1. Deep brown
2. Faint blue
3. Faint brown
4. Colourless

Answer: 2. Faint blue

Question 15. Oxygen is absorbed by

1. Potassium chlorate
2. Alkaline potassium pyrogallate solution
3. Potassium permanganate solution
4. Asbestos

Answer: 2. Alkaline potassium pyrogallate solution

Class 8 Science Gases Review Answers

Question 16. The gas which is used together with hydrogen to produce flame and produce a temperature as high as 2800°c is

1. Nitrogen
2. Oxygen
3. Acetylene
4. Carbon monoxide

Answer: 2. Oxygen

Question 17. \(\mathrm{X}+\mathrm{O}_2 \rightarrow \mathrm{Y}, \mathrm{Y}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Z}\) If Z turns red litmus to blue, then X may be

1. Ca
2. S
3. C
4. P

Answer: 1. Ca

Chapter 3 Some Common Gases Fill In The Blanks

Question 1. About 2500 million years ago a bacteria, called __________ began to split water to form oxygen.
Answer: Cyanobacteria

Question 2. The enzyme __________ found in our body breaks down hydrogen peroxide into water and oxygen.
Answer: Catalase

Question 3. About __________ million years ago, the amount of oxygen reached the present level.
Answer: 500

Question 4. The density of oxygen at normal pressure and temperature is __________ gram per litre.
Answer: 1.428

Question 5. The solubility of oxygen at O°C and 1 atmospheric pressure is __________ mg per litre.
Answer: 14.8

Question 6. The freezing point of liquid oxygen is __________ °C.
Answer: -218

Question 7. The boiling point of liquid oxygen is __________ °C.
Answer: -183

Question 8. At – 183°C, oxygen condenses to form a pale,__________ liquid.
Answer: Bluish

Question 9. When cooled with liquid hydrogen, liquid oxygen gives a __________ coloured solid.
Answer: Blue

Question 10. The oxygen molecule is __________.
Answer: Diatomic

Question 11. Oxygen is non-combustible i.e. it itself does not burn. But it is a__________ of combustion.
Answer: Supporter

Question 12. Oxygen gas is slightly __________ than air.
Answer: Heavier

Question 13. Alkaline __________ solution absorbs oxygen and turns dark brown.
Answer: Potassium pyrogallate

WBBSE Class 8 Science Review Exercises

Question 14. __________ solution absorbs oxygen gas quickly and turns blue.
Answer: Ammonium cuprous chloride

Question 15. During the preparation of oxygen by heating potassium chlorate,__________ is used as a catalyst.
Answer: Manganese dioxide(MnO₂)

Question 16. __________ is the lightest element.
Answer: Hydrogen

Question 17. Hydrogen is __________ in water.
Answer: Insoluble

Question 18. Air is about __________ times heavier than hydrogen.
Answer: 14.4

Question 19. __________ gas is the best conductor of heat among all gases.
Answer: Hydrogen

Question 20. __________ adsorbs the largest volume of hydrogen at 0°C.
Answer: Palladium

Question 21. __________ is not a supporter of combustion but it is inflammable.
Answer: Hydrogen

Question 22. __________ is also known as hydrolith.
Answer: Calcium Hydride (CaH₂)

Question 23. A gas balloon is filled up with __________gas.
Answer: Hydrogen

Question 24. The density of hydrogen is __________ gram per litre at 0°C and 1-atmosphere pressure.
Answer: 0.0899

Question 25. __________ is an allotrope of oxygen.
Answer: Ozone

Question 26. __________ is a neutral oxide.
Answer: Nitric Oxide (NO)

Question 27. If only potassium chlorate is heated, oxygen gas is produced at a temperature of __________ °C.
Answer: 650

Chapter 3 Some Common Gases Identify as True or False

Question 1. Oxygen is inflammable and not a supporter of combustion.
Answer: False

Question 2. Hydrogen is inflammable and not a supporter of combustion.
Answer: True

Question 3. LED is used as a bulb in a torch because it has a long lifetime.
Answer: True

Question 4. Sulphur dioxide is a basic oxide.
Answer: False

Question 5. Carbon dioxide is an acidic oxide.
Answer: True

Question 6. Water is an amphoteric oxide.
Answer: False

Question 7. Rusting is basically the oxidation of iron metal.
Answer: False

Question 8. Atomic hydrogen is a very good reducing agent.
Answer: True

Question 9. Oxygen is collected in a gas jar by upward displacement of water.
Answer: False

Question 10. A gas balloon is filled up with hydrogen because it is much lighter than air.
Answer: True

Question 11. Nickel can adsorb a maximum volume of hydrogen at 0°C.
Answer: False

Question 12. Oxygen is produced easily when water is added to solid sodium peroxide at room temperature.
Answer: True

Question 13. The enzyme catalase destroys hydrogen peroxide and superoxide ions in our bodies.
Answer: True

Question 14. Round bottom flasks are round bottom, narrow neck glass containers, most suitable when heating of reaction mixture is required during any chemical reaction.
Answer: True

Question 15. Watch glass is like a small, circular, glass plate used for carrying/ testing a small amount of solid or liquid substances.
Answer: True

Question 16. A funnel is used for transferring liquids or solids from one container to the other.
Answer: True

Question 17. A measuring cylinder is used to measure the volume of liquid.
Answer: True

Question 18. Burettes and pipettes are used for transferring a definite volume of liquid from one container to another.
Answer: True

Question 19. Hydrogen gas is highly soluble in water.
Answer: False

Question 20. The respiration of aquatic animals is dependent on the dissolved oxygen in the water.
Answer: True

Question 21. Hydrogen is used for the hardening of fats and oils (where unsaturated fatty acids are converted to saturated compounds which have higher melting points).
Answer: True

Class 8 Science Chapter 3 Question Bank

Question 22. Magnesium oxide is an acidic oxide.
Answer: False

Question 23. Hydrogen is usually prepared in the laboratory by the reaction between dilute HCI or dilute H₂S04 with granulated zinc in a Woulfe’s bottle (or a round bottom flask).
Answer: True

Question 24. 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.
Answer: True

Question 25. For welding and cutting of metals, oxy-hydrogen flame and oxy-acetylene flame are used.
Answer: True

Question 26. Amphoteric oxide undergoes a neutralization reaction with both acids and bases.
Answer: True

Chapter 3 Some Common Gases Match The Columns

1.

Column – A	Column – B
A. acidic oxide	1. MgO
B. basic oxide	2. SO2
C. amphoteric oxide	3. CO
D. neutral oxide	4. ZnO

Answer: A-2,B-1,C-4,D-3

2.

Column – A	Column – B
A. liquid oxygen	1. CaH2
B. odourless	2. Mixture of CO2 and
C. hydrolith	3. faint blue coloured
D. carbogen	4. hydrogen

Answer: A-3,B-4,C-1D-2

3.

Column – A	Column – B
A. cyanogen	1. oxygen
B.  MnO2	2. hydrogen
C. supporter of combustion	3. production of oxygen by splitting water
D. available in trace amount in the atmosphere in a free state	4. catalyst during the production of oxygen from KCIO

Answer: A-3,B-4,C-1,D-2

 

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.

Read And Learn More WBBSE Solutions For Class 8 School Science Experiments Questions

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).

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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.

Chapter 3 Some Common Gases Some Equipment Used In the Laboratory

Experiments are the most important part of the study of science. Any theory must be verified by experiments and theories are developed on the basis of experiments.

Experiments are probably the only tool available in our hands by which we can understand the underlying logic of various naturally observed phenomena.

Experiments are commonly carried out in laboratories, which need to be equipped properly to carry out various hands-on experiments.

On school premises also, you will find a separate room meant for doing different experiments. In any laboratory, you will find some basic instruments and apparatus, with which experiments related to chemistry can be carried out. Below we will briefly describe these basic instruments and apparatus.

Thermometer

A common mercury thermometer is used in laboratories to measure the temperature of any object under investigation. The scale of the thermometer may be in Centigrade (in which the freezing point of water at 1 atmospheric pressure is 0°C and the boiling point of water at 1 atmospheric pressure is 100°C)

Or in Fahrenheit (in which the freezing point of water at 1 atmospheric pressure is 32°F and the boiling point of water at 1 atmospheric pressure is 212°F). The range of temperature is usually between 0°C to 100°C or may be from 0°C to 200°C.

For temperatures higher and lower than this, thermometers with a suitable range are used. In sophisticated laboratories, electronic apparatus is used to measure the temperature of an object.

Read And Learn More WBBSE Notes For Class 8 School Science

Electrical Cell

Electrical cells are used as a source of electricity. Generally, dry cells, which are better known as batteries, are used.

More than one battery can be connected to meet the requirement for higher energy. During the combination of batteries, the “+” end of one battery must be in contact with the end of another battery, or vice versa.

Otherwise no current will flow. When the batteries are connected to any electrical circuit, the proper ends of the battery must be connected.

WBBSE Class 8 Common Gases notes

Switch

Switches are an important part of an electrical circuit. With this the flow of electricity through an electrical circuit can be started or stopped, as and when required.

(We can find switches in our home also, with which we can switch on or off any electrical equipment such as light, fan, television, etc.).

Wires

Different electrical connections are done using copper wires insulated with polymers like PVC (polyvinyl chloride).

PVCs have good insulating properties, so live wires are covered with them for safety. Long wires are cut to pieces to get wires of suitable lengths required for a particular circuit.

Bulb

To check whether electricity is flowing through a circuit or not, the easiest way is to connect a bulb to the circuit. If the bulb glows,

when an electric current is switched on, then it confirms that electricity is flowing through the circuit.

Otherwise, it will not glow. Generally, small bulbs such as LEDs, are preferably used nowadays. They have a long life and they do not get fused during repeated use.

Understanding common gases for Class 8

Chemical Balance

This is an instrument by which we can weigh any object very accurately. It is also very sensitive (i.e. it can detect even a very small difference in mass as low as 0.1 mg or lower).

These balances are accurate and give the same result in successive weighing. The upper limit of a chemical balance is generally 100 g or 200 g.

Generally, there are two pans in it – on the left pan the sample to be weighed is placed and on the right pan required weights are placed. Forceps are used to place weights on the pan or to remove from there.

This is because oils, dust or dirt present in our hands may alter the exact masses of the weights. Nowadays in advanced laboratories, electronic balance is used which can detect very small weights with utmost accuracy.

(You must have seen similar balances in some grocery shops, but those are not equipped to measure very small weights, but they can weigh up to a few kilograms of substance.)

 

 

Clamp and stand

Iron stands with a heavy base and clamps are used to hold something, particularly glass apparatus such as a burette, separating funnel, condensers etc. during an experiment.

Nowadays, Teflon-coated stands and clamps are available in modern laboratories. The Teflon coatings prevent rusting and ensure longer life.

 

Bunsen burner and spirit lamps

Bunsen burner is used widely in most laboratories for attaining the moderately high temperatures required during any chemical reaction. Here, LPG (liquefied petroleum gas) is ignited to produce flames.

A regulator is there near the base which is used to regulate the flow of air to increase (or decrease) the temperature. Generally, the maximum temperature is attained by adjusting the regulator so as to allow more air than is required to produce a non-luminous flame.

(But too much air produces a “noisy” flame, not suitable for the experimental purpose). Spirit lamps are also frequently used for carrying out small and simple experiments.

 

WBBSE Chapter 3 gases summary

Various Glass Apparatus Used In Laboratory

All the above apparatus are made of glass. (Nowadays “plastic” apparatus made of polymer are being used in modern laboratories).

In order to avoid the introduction of impurities during the analysis or synthesis of any substance, borosilicate glass (or Pyrex glass) is preferred over ordinary soda glass.

They are least affected by acidic or alkaline solutions. Generally, a glass apparatus should not be heated in a naked flame.

 

 

Wire gauze with an asbestos centre should be placed between the glass apparatus and the flame. For special purposes, heat-resistant high silica glasses can be used.

Glass containers of different sizes and shapes are used as per requirement. Test tubes are the most common glass apparatus used in laboratories. It is a glass tube with one open end.

Its wall may be thin (thin glass test tube) or thick (hard glass test tube). Hard glass test tubes are sometimes held directly over the flame during some experiments.

Beakers are cylindrical glass apparatuses with a spout (a squeezed side at its mouth). The presence of a spout makes pouring liquid from it convenient.

 

Also, when the beaker is covered with an ordinary watch glass, it acts as an outlet for any gas evolving from the beaker during a chemical reaction.

Round bottom flasks are round bottom, narrow neck glass containers, most suitable when heating of reaction mixture is required during any chemical reaction.

The round-shaped bottom provides more surface area through which heat can enter and the solution within it is efficiently heated.

Properties of common gases Class 8

Conical flasks (or Erlenmeyer’s flasks) are cone-shaped glass vessels with small, narrow necks.

Conical flasks of the capacity of 100 mL, 250 mL and 500 mL are frequently used in laboratories.

Woulfe’s bottle is a glass bottle with double mouths. It is usually used for the preparation of gases (such as hydrogen).

Through one mouth, chemical substances (reactants) are introduced into the bottle and the gas produced during the chemical reaction comes out through the other mouth.

 

This mouth is connected by a bent glass tube the other end of which is introduced to another larger, cylindrical jar, called a gas jar, in which the liberated gas is collected and stored for future use.

Watch glass is like a small, circular, glass plate used for carrying/testing a small amount of solid or liquid substances.

 

Funnel

It is a cone-shaped glass apparatus fitted with a stem (narrow glass tube) at its bottom.

A funnel is used for transferring liquids or solids from one container to the other (for example a solution of acid is transferred from a glass bottle to a measuring cylinder).

 

Test Tube Holder And Test Tube Rack

A test tube holder is used to hold a test tube (containing any solution) with thin metallic tongs during heating. This way direct contact between the hand and the test tube is avoided.

Test tube racks are meant for placing the test tubes containing solids or liquids vertically. The bottom of a test tube is round, so it cannot stand on its own. Test tube racks are made up of wood or plastic.

 

Chemical properties of common gases Class 8

Measuring Cylinder

A measuring cylinder is used to measure the volume of liquid. It is a cylindrical tube, and on its wall markings are there, from which we can estimate the volume of any liquid poured into it.

Measuring cylinders of different capacities (say 50 mL, 100 mL, etc.) are frequently used. Tripod stand and wire gauze

For heating, a solution is placed in a glass apparatus (say a beaker or a conical flask), and a tripod stand made of pig iron is used.

A wire gauge (preferably asbestos centred) is placed on the tripod stand and the glass apparatus is placed on it.

The Bunsen burner is placed below the wire gauze. Wire gauze evenly distributes the heat coming from the flame throughout its surface.

 

Pipette And Burette

Burettes and pipettes are used for transferring a definite volume of liquid from one container to other. They are frequently used during titration experiments.

Burettes are long, cylindrical tubes of uniform bore throughout the graduated length. A burette is closed at the bottom by means of a glass stopcock.

Burettes are usually of 50 mL capacity and graduate to the tenths of a millilitre. Burettes of another capacity (say 25 mL, 10 mL, etc.) are also used. It is used to deliver variable volumes of liquid.

Pipette consists of a cylindrical bulb joined at both ends to narrow tubes. This type of pipette is commonly known as a transfer pipette.

It has only one mark and it is used to deliver a small, definite volume of liquid. (Another type of pipette known as graduated pipette is also available where the stems are graduated and are used to deliver various small volumes of liquid by using a single pipette.)

Filter Paper

Filter paper is used to separate solid, insoluble particles from a liquid. It is a thick, porous, circular paper. Filter papers having different pore sizes are available.

The filter paper is chosen on the basis of the size of the insoluble, solid particles which are to be separated from the liquid.

After getting some idea about some common laboratory equipment and apparatus, we can now discuss two common gases which are very come to know about their physical and chemical usage.

One of them is oxygen and the other is Properties, their sources, and their preparation of hydrogen. During this brief discussion, we will procedures and about their uses.

Common gases and their uses for Class 8

Oxygen

Nitrogen and oxygen are the main components of the atmosphere by volume.

Oxygen is exchanged between the atmosphere and living species through photosynthesis and respiration.

During photosynthesis, oxygen is produced along with the formation of glucose from carbon dioxide and water in presence of sunlight.

Respiration is the opposite process of photosynthesis.

Here, oxygen is combined with glucose producing water and carbon dioxide and energy is released which sustains all the physicochemical processes occurring within the living bodies.

Scientists believe that the earth was created approximately 4500 million years ago, and life first emerged on earth about 3500 million years ago.

For the first two billion years after the earth was formed, its atmosphere was anoxic (i.e. free from oxygen).

About 2.5 billion years ago oxygen in the atmosphere increased from a trace amount to approximately 1% of the earth’s atmosphere.

This was probably due to a type of bacteria, called cyanobacteria, the first organism capable of producing oxygen through photosynthesis.

In presence of sunlight and with the help of a special protein, they started producing oxygen by splitting water. The level of oxygen in the atmosphere slowly started rising.

Cyanobacteria are found even today in wetlands, ponds and paddy fields. Then came the green algae and other plants.

They too could produce oxygen. The level of oxygen in the atmosphere started rising very slowly.

Another jump in the level of oxygen in the earth’s atmosphere took place approximately 500 million years ago.

The level of oxygen increased to levels closer to the composition of the atmosphere which is found today.

(However, scientists are still not sure what exactly caused this second jump in oxygen level at that time).

Oxygen is essential for almost all living organisms found on earth. It is utilized during respiration.

During respiration, energy is liberated from food kinds of stuff like glucose within the living cells.

Examples of common gases for Class 8

The liberated energy is utilized to maintain all types of physicochemical processes occurring within the body of a living organism, necessary to remain alive.

Respiration is a very efficient process for producing energy from intracellular glucose.

But we should point out that not all living beings are dependent on oxygen. There are places on earth where oxygen is not available.

Into the depth of the marshes and in the sludge of urban sewers, where oxygen is not available, some life forms are available.

These are called obligate anaerobes.

The mechanism of production of energy intracellularly is different from ours. (In feet they will die if they come in contact with oxygen).

But oxygen can also produce harmful chemicals like hydrogen peroxide and superoxides, which can cause irreparable damage to the living cells even when present in minute amounts.

These chemicals can damage DNA. Some scientists have linked them with the process of ageing in living organisms.

They are of the opinion that free radicals like superoxides trigger and increase cell death mechanisms within the body.

Fortunately, there are certain enzymes present in our body which can specifically destroy such harmful species.

For example, the enzyme, catalase, can destroy hydrogen peroxide and form oxygen and water.

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

Oxygen has also found major applications in several industrial processes. We will discuss them later.
Let us now discuss the various methods for the preparation of oxygen in the laboratory.

 

WBBSE Chapter 2 Element Compound And Chemical Reaction Chemical Effects Of Electricity Long Answer Questions

Question 1. Describe briefly the process of electrolysis of molten sodium chloride, using suitable electrodes.
Answer:

The process of electrolysis of molten sodium chloride, using suitable electrodes:

When electricity is passed through the molten or fused sodium chloride through graphite anode and iron cathode, NaCI is electrolyzed.

Na⁺ ions move towards the cathode and negatively charged Cl^– ions move towards the anode. At the cathode, Na⁺ ions take up electrons and are converted to sodium metal.

At the anode, Cl^– ions give up electrons and are ultimately converted to chlorine gas (Cl₂) Electrons are taken up by Na⁺ ions at the cathode.

So reduction occurs at the cathode. Oxidation occurs at the anode because the electron is given up by the Cl₂ ion. So sodium metal is produced at the cathode and chlorine gas is produced at the anode.

Electrode reaction: At cathode (Reduction) : \(\mathrm{Na}^{+}+e \rightarrow \mathrm{Na}\)

At anode (Oxidation): \(2 \mathrm{Cl}^{-}-2 e \rightarrow \mathrm{Cl}_2 (gas)\)

Read And Learn More WBBSE Solutions For Class 8 School Science Long Answer Type Questions

WBBSE Class 8 Chemical effects of electricity long answer questions

Question 2. Describe briefly the process of electrolysis of acidified water, using suitable electrodes.
Answer:

Electrolysis Of Sodium Chloride:

Let us now discuss what happens when fused sodium chloride is electrolyzed using the above setup,

We know that Aien NaCI is dissolved in water, it is completely dissociated into Na⁺ and Cl^– ions and they move freely within the solution. In molten or fused states also Na⁺ and ions are produced which move freely within the molten state.

⇒ \(\mathrm{NaCl} \rightleftharpoons \mathrm{Na}^{+}+\mathrm{Cl}^{-}\)

When electricity is passed through the molten or fused sodium chloride through graphite anode and iron cathode,

Na⁺ ions (which are positively charged) move towards the cathode (connected to the negative terminal of the battery) and negatively charged Cl^– ions move towards the anode (connected to the positive terminal of the battery).

At the cathode, Na⁺ ions take up electrons and are converted to sodium metal. At the anode, Cl^– ions give up their “extra” electrons and are ultimately converted to chlorine gas (Cl₂).

Electrons are taken up by Na⁺ ions at the cathode. So reduction occurs at the cathode while oxidation occurs at the anode because the electron is given up by the Cl^– ion. Hence sodium metal is produced at the cathode and chlorine gas is produced at the anode.

The electrolysis of molten or fused sodium chloride can be summarized in the following table.

 

Electrolysis of molten or fused sodium chloride
Electrolyte	Molten or fused sodium chloride (NaCI)
Electrode	Anode: graphite Cathode: Iron
Electrode reaction :
At the cathode (Reduction):	\(\mathrm{Na}^{+}+\mathrm{e} \rightarrow \mathrm{Na}\)
At anode (Oxidation):	\(2 \mathrm{Cl}^{-}-2 \mathrm{e} \rightarrow \mathrm{Cl}_2 \text { (gas }\)

 

Long answer type questions on chemical reactions for Class 8

Question 3. The phenomenon of electrolysis is actually a phenomenon of oxidation reduction-explain with a suitable example. Or “During electrolysis, reduction occurs at the cathode and oxidation occurs at anode”. Explain with a suitable example.
Answer:

Let us describe it with the help of electrolysis of molten sodium chloride.

When electricity is passed through the molten or fused sodium chloride through graphite anode and iron cathode, positively Na⁺ ions solution of the chloride salts of No, Mg or Ca is electrolyzed, and the metals cannot be extracted.

moves towards the cathode and negatively charged Cl^– ions move towards the anode. At the cathode, Na⁺ ions take up electrons and are converted to sodium metal.

At the anode, Cl^– ions give up the electron and are ultimately converted to chlorine gas (Cl₂). So sodium metal is produced at the cathode and chlorine gas is produced at the anode.

Electrons are taken up by Na⁺ ions at the cathode. So reduction occurs at the cathode. Oxidation occurs at the anode because the electron is given up by the Cl^– ion.

Electrode reaction:At cathode (Reduction) :  \(\mathrm{Na}^{+}+e \rightarrow \mathrm{Na}\)

At anode (Oxidation) : \(2 \mathrm{Cl}-2 e \rightarrow \mathrm{Cl}, (gas)\)

Question 4. Discuss the electro-refining of metal with a suitable example.
Answer:

Purification or electro-refining of metals:

Electrolysis is used in the purification of impure metals that are extracted from their ores. In this process:

A thick rod of impure metal is made the anode. It is connected to the positive terminal of the battery.

A thin strip of pure metal is made of the cathode. It is connected to the negative terminal of the battery.

A water-soluble salt of the metal to be purified is taken as an electrolyte.

On passing an electric current, the metal dissolves from the impure anode and goes into the electrolyte solution. The metal present in dissolved form in the electrolyte gets deposited on the cathode in the pure form.

The impurities are left behind in the electrolyte solution. Metals like copper, zinc and aluminium etc are purified by electrolysis.

Purification of copper may be taken up as a typical example to discuss electro-refining. Copper is obtained from copper ores which contain various impurities.

Hence when copper is extracted from the ores, it contains various impurities. The metallic copper with impurities is not suitable for various uses such as in electrical appliances such as copper wire.

So, impure copper is to be purified by electrolysis. The impure metal in the form of thick blocks is used as an anode and the pure metal in the form of a thin sheet is used as a cathode.

An acidified copper sulphate solution is taken as an electrolyte. On carrying out electrolysis, copper atoms dissolve out as Cu²⁺ ions from the anode and are deposited at the cathode as metallic copper.

So cathode now consists of pure copper. As the process continues more and more copper dissolves from the anode and is deposited as pure copper on the anode. This process is known as the electro-refining of metals.

Electro-refining of copper
Electrolyte	Acidified copper sulphate solution (CuSO4)\(\left[15 \% \mathrm{CuSO}_4+5 \% \mathrm{H}_2 \mathrm{SO}_4\right]\)
Electrode	Anode: copper bar containing impurities Cathode: a thin sheet of pure copper
Electrode reaction :
At anode (oxidation) :	\(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2+}+2 e (pure copper comes into) solution as \left(\mathrm{Cu}_2+\right. ion )\)
At the cathode (reduction):	\(\mathrm{Cu}^{2+}+2 e \rightarrow \mathrm{Cu} (copper metal is deposited on the cathode)\)

 

WBBSE Chapter 2 electricity and chemical reactions detailed answers

Question 5. Describe with a suitable example, the extraction of a metal from its halide salt.
Answer:

The extraction of a metal from its halide salt:

This can be illustrated by the extraction of calcium metal from calcium chloride. In this process, molten calcium chloride is taken in a graphite container.

The walls of the graphite container act as an anode while an iron electrode is immersed partially in the molten salt and acts as a cathode. Calcium metal is formed at the cathode and chlorine gas is liberated at the anode.

Extraction of Calcium by Electrolysis of molten or fused calcium chloride
Electrolyte	Molten or fused magnesium chloride (CaCl2)
Electrode	Anode: graphite Cathode: iron
Electrode reaction:
At the cathode (Reduction) :	\(\mathrm{Ca}^{2+}+2 e \rightarrow \mathrm{Ca}\)
At anode (Oxidation) :	\(2 \mathrm{Cl}-2 e \rightarrow \mathrm{Cl}_2 \text { (gas) }\)

 

(In the industrial process, calcium chloride is mixed with calcium fluoride. This lowers the melting point of calcium chloride. The molten electrolyte is heated at 700°C and at this temperature electrolysis is carried out.)

Question 6. Describe briefly the process of electroplating copper on iron nails.
Answer:

The process of electroplating copper on iron nails:

During the electroplating of copper on iron nails, the pure copper metal is used as anode and the iron nails (i.e., the objects which are to be electroplated) are used as cathode.

An acidified copper sulphate solution is taken as an electrolyte. On carrying out electrolysis, copper atoms gradually dissolve out as Cu²⁺ ions from the anode and are deposited at the cathode as metallic copper.

As the process continues more and more copper dissolves from the anode and is deposited as pure copper on the iron nails (which are used as the anode). So after some time, the iron nails are coated with a thin layer of copper.

Electro-refining of copper on iron nails
Electrolyte	Acidified copper sulphate solution (CuS04)
Electrode	Anode: pure copper metal Cathode: iron nails
Electrode reaction:
At anode (oxidation) :	\(\mathrm{Cu} \rightarrow \mathrm{Cu}^{2+}+2 e (pure copper comes into) solution as \left(\mathrm{Cu}_2+\right. ion )\)
At the cathode (reduction):	\(\mathrm{Cu}^{2+}+2 e \rightarrow \mathrm{Cu} (copper metal is deposited on the cathode)\)

 

In-depth explanations of chemical effects of electricity for Class 8

Question 7. Describe briefly the electroplating of different metals on various objects familiar to us.
Answer:

The electroplating of different metals on various objects familiar to us:

The various objects familiar to us can be electroplated by various metals for various purposes such as to protect from rusting, to improve the external appearance, to enhance commercial value, etc.

In the case of electroplating, the metal which is to be electroplated is used as an anode and the article on which the coating is given is used as the cathode, and a suitable salt solution is used as an electrolyte.

During electroplating, the pure metal from the anode is gradually dissolved as metal ions into the solution and the metal ions are deposited on the cathode material in the form of a thin film.

Various instances have been presented in tabular form below:

Objects which is to electroplated	The metal which has to be electroplated on the objects	The material used as cathode	The material used as anode	Electrolyte
Iron pipe	Zinc	Iron pipe	zinc	An aqueous solution of zinc chloride
The handle of the bicycle made of iron	Chromium	The handle of the bicycle made of iron	Pure sheet of chromium	An aqueous solution of chromic sulphate and chromic acid
Tap made of brass	Chromium	Tap made of brass	Pure sheet of chromium	The aqueous solution of chromic sulphate and chromic acid

 

WBBSE Class 8 Science practice long answer questions on chemical reactions

Question 8. What are the chemical changes associated with electrolysis?
Answer:

The chemical changes associated with electrolysis:

The products obtained at electrodes vary as the electrodes used in electrolysis support. The chemical changes brought about by electrolysis are-

1. A bubble of gases may be formed on the electrodes
2. deposits of metals may occur on electrodes

changes in the colour of the solution may occur Electrolysis of the same electrolyte with different electrodes produce different products.

When an aqueous solution of CuSO₄ is electrolysed with Cu electrodes, SO^2- ions, in preference to OH^– ions accumulate at the anode and Cu²⁺ ions will be discharged at the cathode.

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But, if platinum electrodes are used in this electrolysis, OH^– ions, in preference to SO₄ ions will be discharged at the anode and Cu²⁺ ions will be discharged at the cathode.

Similarly, if an aqueous solution of NaCI is electrolysed with platinum electrodes, H⁺ ions, in preference to Na⁺ ions are discharged at the cathode.

But, if a mercury cathode is used in this electrolysis, Na⁺ ions, in preference to H⁺ ions will be discharged at the cathode.

Chapter 2 Element, Compound, And Chemical Chemical Effects Of Electricity SAQs

Question 1. What do you mean by electrolyte?
Answer:

Electrolyte:

This is the chemical compound which when dissolved in water produces ions and the solution becomes the conductor of electricity.

It undergoes dissociation when electricity is passed through It. The same word Is also commonly used to refer to a solution of all electrolytes in water. The term electrolyte solution is then used.

Question 2. What do you mean by strong electrolyte and weak electrolyte?
Answer:

Strong Electrolyte And Weak Electrolyte:

When dissolved in water or in a molten state, the electrolyte which produces a large number of ions is called the strong electrolyte.

For example, sodium chloride, sodium hydroxide, dilute sulphuric acid etc., are all strong electrolytes.

When dissolved in water or in a molten state, the electrolyte which produces a smaller number of ions is called a weak electrolyte. For example, acetic acid, ammonium hydroxide, etc., are all weak electrolytes.

Chemical effects of electricity concepts with short answers for Class 8

Question 3. During electrolysis of water 10 cc oxygen collects at the anode, what and how much other substance collects at the cathode at that time?
Answer:

Electrolysis of water produces hydrogen and oxygen In the volumetric ratio of 2:1 as shown:

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

Hence when 10 cc of oxygen collects at the anode; O₂ ccs of hydrogen shall collect at the cathode at that time.

Question 4. An aqueous solution of ammonia is a weak electrolyte whereas sodium chloride is a strong electrolyte. Why?
Answer:

An aqueous solution of ammonia is a weak electrolyte whereas sodium chloride is a strong electrolyte:

Ammonia dissolves in water to produce ammonium hydroxide (NH₄OH). In an aqueous medium, NH₄OH dissociates to a very small extent to produce small numbers of NH⁴⁺ ions and OH^– ions.

So it behaves as a weak electrolyte. An aqueous solution of sodium chloride dissociates completely to produce a large number of Na⁺ ions and Cl^– ions. So it behaves as a strong electrolyte.

Read And Learn More WBBSE Solutions For Class 8 School Science Short Answer Type Questions

Question 5. Sugar, wax, butter, etc., are not electrolytes. Why?
Answer:

Sugar, wax, butter, etc., are not electrolytes:

To behave as an electrolyte, a substance should produce ions. But when sugar is dissolved in water, it does not produce any ions. Wax and butter are not soluble in water.

But even in their molten state also, they do not produce any ions. So they are not electrolytes.

Question 6. Why one should not touch an electric switch with a wet hand?
Answer:

Pure water does not conduct electricity. But usually, the water we use contains a significant concentration of salt, which make the water conductor of electricity. So, when one touches an electric switch with wet hands, one may experience an electrical shock.

WBBSE Class 8 Chemical effects of electricity short answer questions

Question 7. When electricians are working with a live wire, they usually stand on wooden tools. Why?
Answer:

Our body is a good conductor of electricity because our blood contains significant concentrations of different salts. So if one touches a live wire, then electricity passes through our body and touches the ground. It may cause even death to the concerned person.

But wood is a bad conductor of electricity. So when someone is standing on a wooden tool, the electric current cannot flow through the human body to touch the ground. Hence, no harm is done.

Question 8. How one can detect the flow of a very small amount of electricity flowing through an electrical circuit?
Answer:

To detect whether any small amount of electricity flowing through an electrical circuit or not, one has to place a magnetic needle pointing north-south, very near the electrical wire.

If even a very small amount of electricity flows through the wire, the magnetic needle is deflected, indicating the flow of electricity through the circuit.

Question 9. What do you mean by electroplating? Why it is done?
Answer:

Electroplating:

The process by which a thin layer of one metal is coated on the surface of the other to protect the latter from rusting or to improve its appearance is called electroplating.

For example, a thin layer of zinc is coated on the surface of the materials made up of iron to prevent it from rusting. Gold is electroplated on silver ornaments to give them a gold-like appearance. Metal fittings are plated with chromium to preserve their polish.

Question 10. What is the main, principle of electroplating?
Answer:

Main Principle Of Electroplating:

In the case of electroplating, the metal which is to be electroplated is used as the anode and the article on which the coating is given is used as the cathode, and a suitable salt solution is used as the electrolyte.

During electroplating, the pure metal from the anode is gradually dissolved as metal ions into the solution and the metal ions are deposited on the cathode material in the form of a thin film.

Short answer questions on chemical reactions for Class 8

Question 11. What do you mean by galvanization? Why it is done?
Answer:

Galvanization:

The process of coating the surface of a material with a thin film of zinc is known as galvanization. Generally surface of the materials made up of iron is coated with zinc to prevent it from rusting.

Question 12. How can you differentiate a substance made of pure stainless steel from an iron substance electroplated with nickel?
Answer:

When an iron substance is coated with nickel, it appears like stainless steel.

So, it is difficult to visibly differentiate a substance made of pure stainless steel from an iron substance electroplated with nickel.

But stainless steel is not attracted by a magnet, but iron is. So when a magnet is brought closer to stainless steel objects it is not attracted, but the iron substances electroplated with a thin layer of nickel will be attracted towards the magnet.

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Question 13. When molten sodium chloride or magnesium chloride or calcium chloride is electrolyzed using suitable electrodes, the respective metal is deposited at the cathode. But when an aqueous solution of the chloride salts of Na, Mg or Ca is electrolyzed, the metals cannot be extracted. Why?
Answer:

When molten sodium chloride or magnesium chloride or calcium chloride is electrolyzed using suitable electrodes, the respective metal is deposited at the cathode. But when aqueous

This is because, when an aqueous solution of metal halides is electrolyzed, both H⁺ ions and Metal ions move towards the cathode. But in this case, H⁺ ions are preferentially discharged at the cathode. As a result, hydrogen gas is formed at the cathode.

At the cathode (Reduction): \(2 \mathrm{H}^{+}+2 e \rightarrow \mathrm{H}_2\)

Chapter 2 Element Compound And Chemical Reaction Chemical Reaction SAQs

Question 1. What do you mean by a chemical reaction? What are reactants and products? Describe a suitable example.
Answer:

Chemical Reaction:

A chemical reaction means a permanent rearrangement or regrouping between the atoms or radicals of the combining substances (i.e., the reactants) to form one or more substances of different properties.

Chemicals participating in a reaction are called reactants.

Chemical substances that are produced due to a permanent rearrangement or regrouping between the reactants are called products.

For example, \( \mathrm{N}_2+3 \mathrm{H}_2 \rightarrow 2 \mathrm{NH}_3\) . In this reaction, nitrogen and hydrogen are reactants, and due to the reaction between them, ammonia is produced. The chemical properties of ammonia are different from nitrogen and hydrogen.

Question 2. What are the factors we need to induce, influence and regulate chemical changes?
Answer:

The factors we heed to induce, influence and regulate chemical changes are physical contact, heat, light, pressure, the presence of ‘solvent, electricity and catalyst.

Question 3. Why heat is needed to increase the rate of a chemical reaction?
Answer:

Heat is a form of energy. So when heated, the reactants gain energy. The kinetic energy of the reactant molecules is increased. These energized substances then react at a much faster rate.

WBBSE Class 8 Elements and Compounds short answer questions

Question 4. Give one example to show that a chemical reaction is initiated when the reactants are heated.
Answer:

Heat is one of the factors which is needed to initiate a chemical reaction. For example, when blue crystals of hydrated cupric nitrate are heated, the evolution of brown fumes of nitrogen dioxide (NO₂)takes place and a black residue of cupric oxide (CuO) is left. In absence of heating no such reaction occurs,

⇒ \(\begin{gathered}
\text { heat } \\
2 \mathrm{Cu}\left(\mathrm{NO}_3\right)_2 \rightarrow 2 \mathrm{CuO}+4 \mathrm{NO}_2+\mathrm{O}_2
\end{gathered}\)

Question 5. Give one example to show a chemical reaction which is initiated by pressure.
Answer:

Industrially ammonia gas is produced from nitrogen and hydrogen gas at a pressure as high as 200 times the normal atmospheric pressure.

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

Question 6. Give one example to show that a chemical reaction which is initiated by sound.
Answer:

Sound is sometimes a factor which initiates a chemical reaction. For example, when a loud sound is made near a gas jar filled with acetylene gas (C₂H₂), the gas dissociates into carbon and hydrogen.

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

Question 7. Give one example to show a chemical reaction which is initiated by electricity.
Answer:

If electricity is passed through water (through two platinum electrodes), containing a little amount of sulphuric acid, then hydrogen and oxygen gas are separately evolved at two electrodes. In this case, the water undergoes electrolysis. The dissociation of water is achieved by passing electricity through it.

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

Short answer questions on chemical reactions for Class 8

Question 8. In a photography shop, there is a dark room. Why?
Answer:

In a photography shop, there is a dark room

In film cameras, photographic films are used. Films are coated with various substances which remain unchanged when stored in dark.

When exposed to light, the photo-sensitive substances undergo photochemical reactions. To avoid exposure to light, a dark room is necessary.

Question 9. Give a suitable example to indicate the significance of the presence of a solvent during a chemical reaction.
Answer:

The presence of solvent is very important for many of the reactions.

For example, when dry baking soda (NaHCO₃) and dry crystals of tartaric acid are mixed, no visible change occurs, indicating no reaction is taking place.

But when an aqueous solution of baking soda and an aqueous solution of tartaric acid are mixed with each other, bubbles of carbon dioxide evolved from the reaction mixture.

NaHCO₃ and tartaric acid are both soluble in water. When they are dissolved in water, then sufficient contact between the reactants is achieved and the reaction takes place.

Question 10. If ether or benzene is used as a solvent during the reaction between baking soda and tartaric acid, no reaction takes place. Why?
Answer:

If ether or benzene is used as a solvent during the reaction between baking soda and tartaric acid, no reaction takes place.

Baking soda is an “ionic compound”, and it produces ions when dissolved in water. Ionic substances are soluble in water but not in organic solvents like benzene, ether, etc. So, when organic solvents like benzene are used as the solvent, baking soda remains insoluble, and hence no reaction takes place.

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Question 11. What is meant by a catalyst and a catalytic reaction?
Answer:

Catalyst And a catalytic reaction:

The substance which can influence (i.e., increase or decrease) the rate of a chemical reaction is called a catalyst There are certain catalysts that slow down the rate of a particular chemical reaction. They are termed negative catalysts. The chemical reaction which involves the use of a catalyst is called a catalytic reaction.

Question 12. What is an enzyme?
Answer:

Enzyme:

Enzymes are biological molecules that act as catalysts and help complex reactions occur. The enzymes are basically proteins, but they may be associated with non-protein substances (known as coenzymes or prosthetic groups) that are essential for the action of the enzyme.

Question 13. Why do we get a smell of ammonia in the urinal?
Answer:

Urea (N H₂CON H₂) is present in the urine of mammals. Microbes in urinals decompose urea present in urine to produce ammonia and carbon dioxide, and that is why we get the smell of ammonia in the urinal.

⇒ \(\mathrm{NH}_2 \mathrm{CONH}_2+\mathrm{H}_2 \mathrm{O} \rightarrow 2 \mathrm{NH}_3+2 \mathrm{CO}_2\)

Question 14. Why is the use of protective goggles strongly recommended during welding?
Answer:

Burning of acetylene gas (C₂H₂) produces so much heat that temperature may rise above 2000°C and at this temperature, even iron melts. So this exothermic reaction is utilized for welding.

⇒ \(\mathrm{C}_2 \mathrm{H}_2+5 \mathrm{O}_2 \rightarrow 4 \mathrm{CO}_2+2 \mathrm{H}_2 \mathrm{O}+\text { heat }\)

At this high temperature, light is emitted which contains a significant amount of ultraviolet (UV) radiation which is very harmful to our eyes. So use of a proper spectacle or protective goggles which can shield our eyes from UV radiation is strongly recommended during welding.

Question 15. When water is added to calcium oxide, then steam is evolved. Why?
Answer:

When water is added to calcium oxide, then steam is evolved.

When water is added to calcium oxide (CaO), calcium hydroxide is produced. This is an exothermic reaction. A large amount of heat is liberated and this heat is sufficient to vaporize water. That is why steam evolved from the reaction mixture.

⇒ \(\mathrm{CaO}+\mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_{2-}+\text { heat }(\mathrm{Q})\)

Question 16. How can you recognize a process as a chemical reaction?
Answer:

Primarily if:

1. something is precipitated;
2. some change of colour of the reaction mixture occurs and
3. heat is absorbed or evolved, then we may conclude that a chemical change has taken place.

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Question 17. Name three industrial processes where catalysis finds an application. Also, mention the names of the catalysts used.
Answer:

 

Nature of matter concepts with short answers for Class 8

Question 18. Why is the layer of rust on an iron article unable to protect the underlying surface?
Answer:

Layer of rust on an iron article unable to protect the underlying surface:

The main constituent of rust is hydrated Ferric oxide (Fe₂O₃nH₂O). Moreover, CO₂ is produced in the process of rusting which itself accelerates more rusting in the presence of O₂ and H₂O

Thus a layer of rust cannot prevent further rusting of the underlying surface the brown coating is thickened and iron becomes brittle.

Question 19. Give one example of a redox reaction which is also
Answer: 

Example of a redox reaction which is also

1. A combination reaction
2. A displacement reaction

1. A combination reaction:

 

 

2. A displacement reaction:

WBBSE Class 8 Science practice short answer questions on chemical reactions

Question 20. The same substance may act both as an oxidising and a reducing agent-give an example to establish the phenomenon.
Answer:

The same substance may act both as an oxidising and a reducing agent

SO₂ is commonly regarded as a good reducing agent.

It reduces a purple (acidified) solution of potassium permanganate to an almost colourless manganous salt and is itself oxidised to sulphuric acid.

 

 

But during its reaction with hydrogen sulphide, it acts as an oxidising agent and oxidises H₂S to S. In this reaction, SO₂ itself is reduced to sulphur.

Hence, in different reactions, the same substance may sometimes play the role of an oxidant and sometimes that of a reductant.