Class 10 Physical Science

Chapter 3 Chemical Calculations Synopsis

1. The chemical substances participating in a chemical reaction are known as reactants and the substances produced in the reaction are called products.

2. A chemical equation is the representation of a chemical reaction with the help of symbols and formulas. The reactants are written on the left side while the products are written on the right side of the equation.

3. A chemical equation gives us qualitative as well as quantitative information about the reaction. We can calculate the mass, number of moles and volumes of the involved reactants and products from the chemical equation.

4. Law of conservation of mass: According to the law of conservation of mass proposed by Lavoisier, in any physical or chemical change total mass of all the substances before the change is equal to the total mass of all the substances formed after the change.

5. Law of conservation of mass and energy: The law states that the total amount of mass and energy remains constant before and after any chemical or physical transformation.

Read and Learn Also WBBSE Solutions for Class 10 Physical Science and Environment

6. Equivalence of mass and energy: According to Einstein’s theory of relativity, mass and energy are different aspects of the same entity and are interconvertible.

If a substance ‘A’ of mass ‘m’ is converted into another substance ‘B’ to produce an equivalent amount of energy T, then according to Einstein’s equation of mass-energy equivalence, E=mc², where c = velocity of light in vacuum.

7. In normal chemical reactions very minute change of mass takes place, which is too small to be measured. Hence, the law of conservation of mass and energy is not applicable to a general chemical reaction. However, the law is applicable to nuclear reactions where a change in energy is detectable.

8. The mass (measured in gram units) of 1 L of a gas at STP is known as the standard density of the gas.

9. The vapour density or relative density of a gas is defined as the ratio of the mass of a given volume of the gas at a given temperature and pressure to the mass of an equal volume of hydrogen gas at the same temperature and pressure.

10. Vapour density is a ratio and hence, it has no unit.

11. Vapour density is independent of temperature and pressure. So, its value is constant at any temperature or pressure.

12. Molecular mass of a gas is twice its vapour density.

13. At STP, the volume of 1 mol of any gas is 22.4 L.

14. At STP, the mass of 1 L of any gas = vapour density of the gas x 0.089

15. Calculation of relative quantities of reactants and products in a chemical reaction is known as stoichiometry.

Chapter 3 Chemical Calculations Short And Long Answer Type Questions

Question 1. Discuss the significance of a balanced chemical equation in chemical calculation.
Answer:

The significance of a balanced chemical equation in chemical calculation

A chemical equation provides us with qualitative as well as quantitative information about a chemical reaction.

Qualitative information: From the chemical equation we can easily identify the reactants and products in a chemical reaction.

1. Quantitative information:
   
   It gives us an idea about the number of atoms and molecules of reactants and products involved in the reaction.
2. It gives us an idea about the mass of the reactants that reacts with each other and the mass of the products formed.
3. If the reactants or products are gases, then at a given temperature and pressure, their volume can be determined.

Question 2. What information can be obtained from the equation \(2 \mathrm{H}_2+\mathrm{O}_2 \rightarrow 2 \mathrm{H}_2 \mathrm{O}\)?

Answer:

Information can be obtained from the equation \(2 \mathrm{H}_2+\mathrm{O}_2 \rightarrow 2 \mathrm{H}_2 \mathrm{O}\)

The equation provides us with the following information:

Qualitative information: Hydrogen and oxygen react with each other to form water.

Quantitative information:

1. 2 mol of hydrogen reacts with 1 mol of oxygen to produce 2 mol of water.
2. Here, 2 X (2 x l)g = 4 g hydrogen reacts with (2 X 16)g = 32g oxygen to produce 2 x (2 x 1 + 16) = 36g water.
3. At the same temperature and pressure, 2 volumes of hydrogen combine with 1 volume of oxygen to produce 2 volumes of water vapour.
4. At STP, 2 x 22.4L = 44.8L hydrogen combines with 22.4 L oxygen to produce 2 x 22.4 L = 44.8 L steam.

Question 3. State the limitations of a chemical equation.
Answer:

A chemical equation cannot give us information about:

1. Concentration of reactants and products.
2. The rate of the reaction or the time required for the completion of the reaction.
3. Whether the reaction has completed or not.

Question 4. In a chemical reaction, mass is conserved explain.
Answer:

In a chemical reaction, mass is conserved as

In a chemical reaction, reactants are transformed into products. The total mass of the products is found to be equal to the total mass of the reactants.

Let us consider, A and B react with each other to produce C and D in a reaction. Hence, the mass of A + mass of B = mass of C + mass of D Thus, the total mass of A and B = total mass of C and D.

Question 5. Show with the help of an experiment that mass is conserved in a chemical reaction.
Answer:

With The help of an experiment that mass is conserved in a chemical reaction:

The law of conservation of mass can be proved with the help of rusting of iron.

Materials required: A hard glass test tube, some new iron nails, water, rubber cork, and common balance.

Experiment: A small amount of water is taken in a hard glass test tube and some glossy iron nails are dipped into the water. The mouth of the test tube is covered with the help of a rubber cork. The mass of the test tube is measured with a common balance and is left undisturbed for some days.

Observation: After some days it will be observed that brown-coloured rust is formed on the iron nails. In this condition, the mass of the test tube is measured again with a common balance and it is observed that the mass of the test tube before and after the experiment is equal.

Conclusion: Iron nails react with oxygen and water present in the test tube to form rust.

∴ \(\mathrm{Fe}+\mathrm{O}_2 \text { + water vapour } \rightarrow \mathrm{Fe}_2 \mathrm{O}_3 \cdot x \mathrm{H}_2 \mathrm{O}\)

x = a number of water molecules.

In this case, the total mass of the iron piece along with unreacted oxygen and water vapour is equal to the mass of the rusted iron. Thus, we can say that mass is conserved during chemical reactions.

Question 6. A piece of iron gets heavier when it is kept in moist air for a long time state whether mass is conserved in this case?
Answer:

When iron is kept in moist air for a long time rust is formed over the surface of the iron piece. Oxygen and water vapour present in the air react with iron to form hydrated ferric oxide (Fe₂O₃ . xH₂O).

As a result, the mass increases. So, the total mass of the iron piece along with oxygen and water vapour is equal to the mass of rusted iron. Hence, mass is conserved in this process.

Question 7. On burning, the mass of a candle decreases state whether mass is conserved in this process.

Answer: On burning, the wax of the candle reacts with aerial oxygen to form carbon dioxide (CO₂), other gases and water vapour. These products escape into atmosphere. If all the products are collected together, then it would be observed that the total mass of the products is equal to the total mass of wax and used oxygen.

Thus, it can be said mass of wax + mass of used oxygen before burning = mass of remaining candle + mass of produced CO₂ + mass of other gases produced + mass of water vapour.
Hence, mass is conserved in this process.

Question 8. Explain with reasons whether law of conservation of mass is followed during electrolysis of acidified water.

Answer: During electrolysis of acidified water, water decomposes to form hydrogen and oxygen.

If electrolysis is carried out in a closed vessel it will be observed that the total mass remains same before and after the electrolysis, i.e., total mass of acidified water = mass of remaining acidified water + mass of produced hydrogen + mass of produced oxygen.

So, it can be said that law of conservation of mass is obeyed during the electrolysis of acidified water.

Question 9. In the normal chemical reactions, mass does not undergo measurable change-explain.

Answer: In the normal chemical reactions, heat is either evolved or absorbed. According to mass-energy equivalence, if heat is evolved in a reaction, an equivalent amount of mass of the reactant will decrease. On the other hand, if heat is absorbed, an equivalent amount of mass of the reactant will increase.

But, in normal chemical reactions, the change in mass observed due to the absorption or evolution of heat is too small to be measured with a balance. Hence, we can say that no measurable change of mass takes place in a chemical reaction.

Question 10. What is the significance of the equation E=mc²? In normal chemical reactions heat is either evolved or absorbed, yet, the change of mass is not observed. Why?

Answer: When a substance of mass ‘m‘ is converted into another substance, an equivalent amount of energy E is produced. The equation, E=mc², represents the conversion of mass into energy and vice versa, where c = velocity of light in vacuum.

The change of mass involved in the normal chemical reactions is too small to be measured. Thus, a change of mass is not generally observed in normal chemical reactions.

Question 11. The law of conservation of mass is not applicable in high energy transformations— explain.
Answer:

During high-energy transformations such as in nuclear reactions, the total mass is converted into an equivalent amount of energy.

As a result, the large amount of energy produced can be determined from Einstein’s equation, E=mc². Hence, it can be said that in high-energy transformations law of conservation of mass is not valid. Instead, in such cases, law of conservation of mass and energy is applicable.

Question 12. 1g of a substance is completely converted into energy. Using the equation E=mc² find the amount of energy obtained due to the conversion.

Answer: The eqaution, E=mc², represents mass-energy equivalence where,

E = energy, m = mass and c = speed of light in vacuum = 3 x 10¹⁰ cm . s^-1

∴ E = 1 x (3 x 10¹⁰)²erg

= 9 x 10²⁰ erg = 2.14 x 10¹³ calorie. [∴1 cal = 4.2 x 107 erg]

Question 13. What do you mean by standard density of a gas? What is vapour density or relative density of a gas?

Answer:

1. The mass (measured in gram units) of 1 litre of a gas at STP is known as standard density of the gas.

2. The vapour density or relative density of a gas is defined as the ratio of mass of a given volume of the gas at a given temperature and pressure to the mass of an equal volume of hydrogen gas at the same temperature and pressure.

Question 14. Deduce the relationship between the standard density and vapour density of a gas.

Answer: Vapour density of a gas = mass of a given volume of the gas at a given temperature and pressure/ mass of an equal volume of hydrogen gas at the same temperature and pressure

or, Vapour density of a gas = \(\frac{\text { mass of } V \text { volume of the gas at STP }}{\text { mass of } V \text { volume of hydrogen gas at STP }}\) = \(\frac{\text { mass of } 1 \mathrm{~L} \text { of the gas at STP }}{\text { mass of } 1 \mathrm{~L} \text { of hydrogen gas at STP }}\)

or, Vapour density of the gas =\(\frac{\text { standard density of the gas }}{\text { standard density of hydrogen gas }}\)

or, Vapour density of the gas =\(=\frac{\text { standard density of the gas }}{0.089 \mathrm{~g} \cdot \mathrm{L}^{-1}}\)

Question 15. Deduce the relationship between molecular mass (M) and vapour density (D) of a gas.
Answer:

The relationship between molecular mass (M) and vapour density (D) of a gas;

Vapour density of a gas =  mass of a given volume of the gas at a given temperature and pressure/ mass of equal volume of hydrogen gas at the same temperature and pressure

According to Avogadro’s law, if ‘n’ number of molecules are present in 1/ L of a gas at a given temperature and pressure, ‘n’ number of H₂ molecules will also be present in V L of hydrogen gas at the same temperature and pressure.

Hence, vapour density of the gas (D) = \(\frac{\text { mass of } n \text { molecules of the gas }}{\text { mass of } n \text { molecules of hydrogen }}\)

= \(\frac{\text { mass of } 1 \text { molecule of the gas }}{\text { mass of } 1 \text { molecule of hydrogen }}\)

= \(\frac{\text { mass of } 1 \text { molecule of the gas }}{\text { mass of } 2 \text { atoms of hydrogen }}\) (as hydrogen molecule is diatomic)

= \(\frac{1}{2} \times \frac{\text { mass of } 1 \text { molecule of the gas }}{\text { mass of } 1 \text { atom of hydrogen }}\)

= \(\frac{1}{2} \times \text { molecular mass of the gas }(M)\)

∴ \(D=\frac{M}{2}\)  ∴ M = 2D

Thus, the molecular mass of gas = 2 Χ vapour density of the gas.

Question 16. How will you understand whether a gas is heavier than air or not? The vapour density of chlorine is 35.5. How will it displace air from a gas jar?
Answer:
1. lf the vapour density of a gas is found to be greater than that of air (14.4), it can be assumed that the gas is heavier than air.

2. The vapour density of chlorine is greater than air. Hence, it is heavier than air. So, there will be an upward displacement of air when chlorine gas is collected into a gas jar filled with air.

Question 17. Vapour density of a gas does not change with temperature justified.

Answer: The vapour density or relative density of a gas is defined as the ratio of the mass of a given volume of the gas at a given temperature and pressure to the mass of an equal volume of hydrogen gas at the same temperature and pressure.

With an increase or decrease in temperature the volumes of hydrogen as well as the concerned gas increase or decrease in the same proportion. Hence, the vapour density of a gas does not change, with temperature.

Question 18. Air contains 74% nitrogen, 24% oxygen and 2% CO₂ by volume. Which of the following gases are heavier and which ones are lighter than air CO₂, NH₃, CH₄, and Cl₂?

Answer: Average molecular mass of air = \(\frac{28 \times 74+32 \times 24+2 \times 44}{100}\) = 29.28.

So, Vapour density of air = \(\frac{29.28}{2}\) = 14.64.

Now, the vapour density of CO₂ = \(\frac{44}{2}\) = 22

Hence, it is heavier than air.

Vapour density of NH₃ = \(\frac{17}{2}\) = 8.5.

Hence, it is lighter than air.

Vapour density of CH₄ =\( \frac{16}{2}\) = 8.

Hence, it is lighter than air.

Vapour density of Cl₂ = \(\frac{71}{2}\) = 35.5.

Hence, it is heavier than air.

Question 19. Dry air is heavier than moist air. Explain with reason.

Answer:

Dry air is heavier than moist air:

The average vapour density of air is 14.4 while that of water vapour is 9. Hence, a given volume of dry air is heavier than an equal volume of water vapour. Moist air contains more amount of water vapour as compared to dry air. That is why dry air is heavier than moist air.

Chapter 3 Chemical Calculations Very Short Answer Type Questions Choose The Correct Answer:

Question 1. Law of conservation of mass was proposed
1. Dalton
2. Lavoisier
3. Arrhenius
4. Proust

Answer: 2. Lavoisier

Question 2. The difference in mass between reactants and products is
1. Observed in low-energy transformations
2. Observed in a general chemical reaction
3. Observed in high-energy transformations
4. Never observed

Answer: 3. Observed in high-energy transformations

Question 3. Two substances, A and B react with each other to produce C and D. Which of the following statements is true?
1. Total mass of A and B > total mass of C and D
2. Total mass of A and B < total mass of C
3. Total mass of A and B = total mass of C and D
4. Mass of A = mass of 8 and mass of C = mass of D

Answer: 3. Total mass of A and B = total mass of C and D

Question 4. The mass-energy equivalence is represented by the equation
1. E²=mc
2. E=mc²
3. E= m²c
4. E = mc

Answer: 2. E=mc²

Question 5. Which of the following cannot be determined from a chemical equation?
1. Number of moles of reactants & products
2. Volume of reactants and products at STP
3. Concentration of reactants and products
4. Mass of reactants and products.

Answer: 3. Concentration of reactants and products

Question 6. The equation E=mc² is related to
1. Newton
2. Planck
3. Avogadro
4. Einstein

Answer: 4. Einstein

Question 7. Measurable change in mass takes place in
1. General chemical reactions
2. Nuclear reactions
3. Exothermic reactions
4. Endothermic reactions

Answer: 2. Nuclear reactions

Question 8. In the equation E=mc², the terms E, m and c respectively represents
1. Mass, energy and speed of light in vacuum
2. Speed of light in vacuum, mass and energy
3. Energy, mass and speed of light in vacuum
4. Energy, speed of light in vacuum and mass

Answer: 3. Energy, mass and speed of light in vacuum

Question 9. According to mass-energy equivalence, in an exothermic reaction the mass of products
1. Will decrease in an equivalent amount
2. Will increase in an equivalent amount
3. Will increase or decrease in an equivalent amount
4. Will remain unchanged

Answer: 1. Will decrease in an equivalent amount

Question 10. According to mass-energy equivalence, in an endothermic reaction the mass of products
1. Will decrease in equivalent amount
2. Will increase in equivalent amount
3. Will increase or decrease in equivalent amount
4. Will remain unchanged

Answer: 2. Will increase in equivalent amount

Question 11. Total amount of mass and energy after any physical or chemical transformation
1. Decreases
2. Increases
3. Remains constant
4. Norie of these

Answer: 3. Remains constant

Question 12. The relationship between mass and number of moles is given by

1. \(\text { number of moles }=\frac{\text { molar mass }}{\text { given mass }}\)

2. \(\text { molar mass }=\frac{\text { number of moles }}{\text { given mass }}\)

3. \(\text { number of moles }=\frac{\text { given mass }}{\text { molar mass }}\)

4. number of moles = given mass x molar mass

Answer: 3. \(\text { number of moles }=\frac{\text { given mass }}{\text { molar mass }}\)

Question 13. The amount of N₂ produced from 64g NH₄NO₂ is
1. 14g
2. 28g
3. 32g
4. 40g

Answer: 2. 28g

Question 14. The volume of NO₂ produced when 2 mol of NO reacts with O₂ at STP is
1. 11.2 L
2. 44.8 L
3. 22.4 L
4. 5.6 L

Answer: 3. 22.4 L

Question 15. Which of the following relations is correct?
1. At STP, mass of 1L of a gas = vapour density of the gas x 0.89
2. At STP, mass of 1L of a gas = vapour density of the gas x 0.089
3. Vapour density of a gas = 0.089 x mass of 1L of a gas at STP
4. Vapour density of a gas = 2.24 x molecular masss of the gas

Answer: 2. At STP, mass of 1L of a gas = vapour density of the gas x 0.089

Question 16. 1 mol of nitrogen combines with 3 mol of hydrogen to produce.
1. 1 mol NH₃
2. 2 mol NH₃
3. 3 mol N₃  
4. 4 mol NH₃

Answer: 2. 2 mol NH₃

Question 17. The relationship between molecular mass and vapour density is given by
1. 2M = D
2. M = D²
3. M = 3D
4. M = 2D

Answer: 4. M = 2D

Question 18. The vapour density of SO₂ is
1. 16
2. 32
3. 64
4. 48

Answer: 2. 32

Question 19. The molecular mass of a gas is 46. Its vapour density is
1. 23u
2. 23 amu
3. 23 g
4. 23

Answer:  4. 23

Question 20. The vapour density of a gas is 22. The volume of 44g of the gas at STP is
1. H.2L
2. 5.6L
3. 22.4L
4. 44.8L

Answer: 3. 22.4L

Question 21. The average vapour density of air is
1. 14.4
2. 16.5
3. 17.2
4. 20.1

Answer: 1. 14.4

Question 22. The molecular mass of chlorine is 71. Which of the following statements is correct?
1. Chlorine gas is lighter than air
2. Chlorine gas is heavier than air
3. Chlorine gas has the almost same mass as that of air
4. Unpredictable

Answer: 2. Chlorine gas is heavier than air

Question 23. Mass of 112 mL O₂ at NTP is
1. 0.64 g
2. 0.96 g
3. 0.32 g
4. 0.16g

Answer: 4. 0.16g

Question 24. Mass of 6.022 x10²² number of molecules of water is
1. 18 g
2. 1.8 g
3. 9 g
4. 0.18 g

Answer: 2. 1.8 g

Question 25. Volume of CO₂ gas at STP, obtained by heating 1 mol of CaCO₃ is
1. 22.4 L
2. 11.2 L
3. 5.6 L
4. 44.8L

Answer: 1. 22.4 L

Question 26. The percentage increase in mass of an Mg wire after its combustion in pure oxygen is
1. 100%
2. 66.7%
3. 24%
4. 80%

Answer: 2. 66.7%

Question 27. The volume of hydrogen at STP obtained from the reaction of 4 mol Na with water is
1. 22.4 L
2. 44.8 L
3. 11.2 L
4. 5.6L

Answer: 2. 44.8 L

Question 28. The number of moles of NH₃ gas formed by the reaction of 2 mol N₂  with 6 mol H₂ is
1. 6 mol
2. 2 mol
3. 8 mol
4. 4 mol

Answer: 4. 4 mol

Question 29. Vapour density of sulphur trioxide is
1. 40
2. 32
3. 36
4. 80

Answer: 1. 40

Question 30. Vapour density of a gas is 32. Which of the following is the molecular mass of the gas?
1. 8
2. 16
3. 32
4. 64

Answer: 4. 64

Question 31. The volume of CO₂ at STP, obtained by the complete combustion of 1.2 g of C is
1. 1.2 L
2. 2.24 L
3. 4.8 L
4. 4.4 L

Answer: 2. 2.24 L

Question 32. Molecular mass of methane is 16. Which of the following is its vapour density?
1. 22.4
2. 8
3. 16
4. 32

Answer: 2. 8

Question 33. The volume of 2 mol CO₂ at STP is
1. 11.2 L
2. 22.4 L
3. 44.8 L
4. 44.8 mL

Answer: 3. 44.8 L

Question 34. The mass of 1 L of H₂ gas at STP is
1. 1g
2. 2g
3. 0.089 g
4. 22.4 g

Answer: 3. 0.089 g

Question 35. Amount of nitrogen required to produce 4 mol NH₃ is
1. 14 g
2. 28 g
3. 42 g
4. 56 g

Answer: 4. 56 g

Question 36. The volume of 1.7 g of ammonia gas, at STP is
1. 22.4 L
2. 2.24 L
3. 11.2 L
4. 1.12 L

Answer: 2. 2.24 L

Question 37. The volume of 2.8 g ethylene, at STP is
1. 224 L
2. 22.4 L
3. 2.24 L
4. 0.224 L

Answer: 3. 2.24 L

Question 38. Vapour density of a gas is 16. The mass of 11.2 L of the gas at STP is
1. 22.4 g
2. 8 g
3. 16 g
4. 11.2 g

Answer: 3. 16 g

Question 39. Ratio of the standard density of a gas at NTP to its vapour density is
1. 0.89
2. 0.98
3. 0.089
4. 8.9

Answer: 3. 0.089

Question 40. The volume of 17 g of a gas at STP is 22.4 L. Vapour density of the gas will be
1. 17
2. 0.85
3. 8.5
4. 1.7

Answer: 3. 8.5

Question 41. The ratio of molecular mass to the vapour density of a gas with respect to hydrogen is
1. 2:1
2. 1:2
3. 1:1
4. 1:3

Answer: 1. 2:1

Question 42. Number of molecules formed by the complete reaction of 1 mol of C with 1 mol of O₂ is
1. 6.022 X 10²³
2. 1.806 X 10²⁴
3. 6.022 X 10²²
4. 6.022 x 10²⁴

Answer: 1. 6.022 X 10²³

Question 43. Value of c in the equation, E=mc² is
1. 3 x 10¹⁰ m/s
2. 3 x 10⁹ cm/s
3. 3 x 10⁹ m/s
4. 3 x 10¹⁰ cm/s

Answer: 4. 3 x 10¹⁰ cm/s

Question 44. If the population of the earth be 7 x 109, then the number of moles of human is
1. 1
2. 6.022 X 10~14
3. 1.16 X 10~23
4. 1.16 x 10~14

Answer:

Question 45. Mass of 11.2 L CO₂ at STP is
1. 88 g
2. 44 g
3. 32 g
4. 22 g

Answer: 4. 22 g

Question 46. Molar mass of water (in g .  mol^-1 unit) is
1. 16
2. 18
3. 20
4. 22

Answer: 2. 18

Question 47. Mass of 12.046 X 10²³ number of molecules of N₂ is
1. 28 g
2. 56 g
3. 14 g
4. 7g

Answer: 2. 56 g

Question 48. If the volume of 44 g of CO₂ at STP be 22.4L, the mass of 5.6 L of CO₂ will be
1. 11g
2. 22g
3. 33 g
4. 44 g

Answer: 1. 11g

Question 49. According to E=mc² equation, the unit of E will be
1. meV
2. MeV
3. mev
4. mV

Answer: 2. MeV

Question 50. What is the amount of chalk to be reacted with HCI to produce 5.6 L of CO₂ at STP?
1. 25 g
2. 30g
3. 15 g
4. 20 g

Answer: 1. 25 g

Question 51. If the vapour density of a compound be 13, its molecular formula may be
1. CO₂
2. C₂H₄
3. C₂H₆
4. C₂H₂

Answer: 4. C₂H₂

Question 52. The gas whose vapour density is equal to that of CO gas is
1. O₂
2. N₂
3. H₂
4. Cl₂

Answer: 2. N₂

Question 53. Amount of carbon required to produce 2 mol CO₂is
1. 12 g
2. 6g
3. 24 g
4. 44 g

Answer: 3. 24 g

Question 54. The percentage of Ca in CaCO₃ is
1. 25
2. 40
3. 60
4. 80

Answer: 2. 40

Question 55. Amount of hydrogen formed by electrolysis of 5 mol of water is
1. 4 g
2. 5 g
3. 20 g
4. 10g

Answer: 4. 10g

Question 56. Vapour density of carbon dioxide is
1. 22
2. 11
3. 44
4. 88

Answer: 1. 22

Question 57. 22 g CO₂ is formed by the reaction of 6 g carbon with oxygen, the required amount of oxygen is
1. 6g
2. 18g
3. 16g
4. 28g

Answer: 3. 16g

Chapter 3 Chemical Calculations Answer In Brief

Question 1. What do you mean by ‘conservation of mass’ in a chemical reaction?
Answer: In a chemical reaction, total mass of all the reactants is equal to total mass of all the products. This is known as the ‘conservation of mass’ in a chemical reaction.

Question 2. State the law of conservation of mass and energy.
Answer: The law of conservation of mass and energy states that the total mass and energy before and after a transformation is always constant.

Wbbse class 10 physical science chapter 3 question answers

Question 3. When is the ‘law of conservation of mass and energy’ applicable instead of ‘law of conservation of mass’?
Answer: In high energy transformations like nuclear reactions, ‘law of conservation of mass and energy’ is applied instead of ‘law of conservation of mass’.

Question 4. Define ‘reactants’ and ‘products’.
Answer: The chemical substances participating in a chemical reaction are known as ‘reactants’. The substances produced due to chemical reactions among the reactants are known as ‘products’.

Question 5. Is the change in mass detectable in normal chemical reactions?
Answer: No, the change in mass is not detectable in normal chemical reactions.

Question 6. Write the relationship between molecular mass (M) and vapour density (D) of a gas.
Answer: Molecular mass (M)=2x vapour density (D).

Question 7. Vapour density of SO₂ gas at 25°C and normal pressure is 32.-What will be its vapour density at 50°C?
Answer: Since, vapour density is independent of temperature and pressure, hence, the vapour density of SO₂ at 50°C will remain 32.

Question 8. Define ‘stoichiometry’.
Answer: Stoichiometry is defined as the calculation of relative quantities of reactants and products in a chemical reaction.

Wbbse class 10 physical science chapter 3 question answers

Question 9. What is the unit of vapour density?
Answer: Since vapour density is a ratio, thus it has no unit.

Question 10. Mass of 5.6 L of a gas at STP is 11g. Find the vapour density of the gas.
Answer: Mass of 5.6 x 4 = 22.4L of the gas at STP is 11 x 4 = 44g. Hence, vapour density of the gas = 44/2 = 22.

Question 11. Name the physical quantity whose unit is ‘mole’.
Answer:’Mole’ is the SI unit of the amount of substance.

Question 12. What is the number of molecules in 22.4 L of a gas at STP?
Answer: 6.022 x 10²³.

Question 13. What is the number of moles in m g of a gas having molar mass of M g/mol?
Answer: Number of moles of the said gas = M/N.

Question 14. What is the number of molecules in 1.8 g water?
Answer: Number of molecules present in 1.8 g water = 6.022 X 10²³ X 1.8/18 = 6.022 x 10²²

Question 15. How much ammonia is equivalent to 4 mol ammonia?
Answer: 4 mol ammonia is equivalent to 17 x 4 = 68 g of ammonia.

Question 16. What is the mass of 12.046 x 10²³ number of nitrogen molecules?
Answer: Mass of 6.022 x 10²³ number of nitrogen molecules = 28 g

∴ Mass of 12.046 x 10²³ number of nitrogen molecules

∴ \(\frac{28 \times 12.046 \times 10^{23}}{6.022 \times 10^{23}} \mathrm{~g}\)

Question 17. Calculate the molecular mass of a gas if the volume of 50 g of the gas is 44.8 L at STP.
Answer: Gram-molecular mass of the gas = mass of 22.4L of the gas at STP = \(\frac{50 \times 22.4}{44.8}\) = 25g
∴ Molecular mass of the gas is 25.

Question 18. What is the relationship between standard density and vapour density?

Wbbse class 10 physical science chapter 3 question answers

Answer: Standard density = 0.089 x vapour density.

Question 19. If the vapour density of a bimolecular gas be 14, calculate its atomic mass.
Answer: Molecular mass of the gas = 2 x 14 = 28

Since the gas is bimolecular, hence its atomic mass = 28/2 =14.

Chapter 3 Chemical Calculations Fill In The Blanks:

Question 1. A chemical equation. provides us with _________ and ____________ information.
Answer: Quantitative, qualitative

Question 2. The velocity of light in a vacuum is _________.
Answer: 3Χ10⁸m.s^-1.

Question 3. The change in mass is __________ in normal chemical reactions.
Answer: Undetectable

Question 4. Mass is ____________ in case of rusting of iron.
Answer: Conserved

Question 5. The volume of H₂S produced from 88g FeS is __________.
Answer: 22.4L

Question 6. When hydrogen reacts with oxygen to form water, the number of moles ___________.
Answer: Decreases

Question 7. Vapour density is __________ of temperature and pressure.
Answer: Independent

Question 8. CO₂ is heavier than air as the__________ of CO₂ is greater than air.
Answer: Vapour density

Question 9. If the vapour density of a gas is 20, the number of moles of 80g of the gas is ________.
Answer: 2

Question 10. The mass of 1L H₂ at STP is ________.
Answer: 0.089g

Question 11. In a chemical reaction, the total mass of reactants is = the total mass of the _________.
Answer: products

Wbbse class 10 physical science chapter 3 question answers

Question 12. The proposer of E=mc² equation is ____________
Answer: Albert Einsein

Question 13. At STP/ 22.4 L of H₂ reacts with 22.4 L of Cl₂ to produce ________ L of HCl.
Answer: 44.8

Question 14. Vapour density is a ______ less quantity.
Answer: Unit

Chapter 3 Chemical Calculations State Whether True Or False:

Question 1. The total mass remains constant before and after a chemical reaction.
Answer: True

Question 2. The vapour density of a gas is twice its molecular mass.
Answer: False

Question 3. ³⁵Cl atom is taken as the standard for the determination of atomic mass.
Answer: False

Wbbse class 10 physical science chapter 3 question answers

Question 4. 8g oxygen contains 3.011 Χ 10²³ molecules.
Answer: False

Question 5. Vapour density is a unitless quantity.
Answer: True

Question 6. 16g oxygen is equal to 0.5 mole oxygen.
Answer: True

Question 7. 6.022 x 10²³ molcules of water weighs 18g.
Answer: True

Question 8. The substances produced in a chemical reaction are called reactants.
Answer: False

Question 9. The volume of lg gas at STP is 22.4L.
Answer: True

Wbbse class 10 chemical calculation

Question 10. Scientist Lavoisier proved the ‘law of conservation of mass’.
Answer: True

Question 11. Interconversion of mass and energy is possible.
Answer: True

Question 12. At STP, the mass of 22.4 L of any substance refers to its molecular mass.
Answer: False

Question 13. The amount of O₂ formed by heating 4 mol of KCIO₃ is 2 mol.
Answer: False

Question 14. Vapour density of a gas x 0.089 =standard density of the gas.
Answer: True

Question 15. If the molecular mass of a gas is 44, its vapour density will be 24.
Answer: False

Chapter 3 Chemical Calculations Numerical Examples

Question 1. 3 g of carbon reacts completely with 8 g oxygen to produce carbon dioxide. How much carbon dioxide will be produced? Justify your answer.
Answer: During chemical reactions, law of conservation of mass is obeyed and hence, according to the law, total mass of the reactants must be equal to the total mass of the products.

Here, mass of carbon = 3g; mass of oxygen = 8g
Hence, mass of produced CO₂ = (3 + 8)g = 11 g

Question 2.  15.9 g copper sulphate and 10.6 g sodium carbonate react with each other to produce 14.2 g sodium sulphate and 12.3 g copper carbonate. Show that the results are in accordance with the law of conservation of mass.

Answer: According to the given data,

Mass of copper sulphate \(=15.9 \mathrm{~g}\)

Mass of sodium carbonate \(=10.6 \mathrm{~g}\)

So, total mass of reactants =26.5 g similarly for the products,

Wbbse class 10 chemical calculation

Mass of sodium sulphate 14.2 g

Mass of copper carbonate =12.3 g

Total mass of products =26.5 g

Hence, in the given reaction, total mass of the reactants = total mass of the products. So, we can say that the given data is in accordance with the law of conservation of mass.

Question 3. When 3.0g ethane (C₂H₆) is burnt in air 8.8g CO₂ and 5.4g H₂O are produced. Show that the result is in accordance with law of conservation of mass.

Answer: When ethane is burnt in air, following reaction takes place

2\(\mathrm{C}_2 \mathrm{H}_6+7 \mathrm{O}_2 \rightarrow 4 \mathrm{CO}_2+6 \mathrm{H}_2 \mathrm{O}\)

Gram-molecular mass of ethane = 30g.

Hence, from the equation it can be deduced that (2 x 30) g ethane reacts with (7 x 32)g oxygen.

∴3g ethane reacts with \(\frac{7 \times 32 \times 3}{2 \times 30}\) = 11.2 g oxygen.

Thus, total mass of reactants = (3 + 11.2 g) =14.2g

Total mass of products = (8.8 + 5.4)g = 14.2g.

Thus, the result agrees with the law of conservation of mass.

Question 4. 30g potassium chlorate on heating produces 14.9g potassium chloride and 9.6g oxygen. How much potassium chlorate remains unreacted?

Answer: Let us consider, the mass of unreacted potassium chlorate is xg.

In a chemical reaction mass is conserved. Hence,mass of potassium chlorate taken = mass of unreacted potassium chlorate + mass of produced potassium chloride + mass of produced oxygen

∴30 = x+ 14.9 + 9.6

X = 5.5

Hence, 5.5 g potassium chlorate will remain unreacted.

Wbbse class 10 chemical calculation

Question 5. When 12.25g of potassium chlorate is strongly heated, it produces 7.45g potassium chloride and gives out 3.36 L oxygen at standard temperature and pressure. Show that the reaction obeys the law of conservation of mass.

Answer: The molecular mass of oxygen = 32.

Hence, at SIP 22.41 oxygen weighs 32g.

∴ At STP 3.36 L oxygen weights \(\frac{32 \times 3.36}{22.4}\) g= 4.8g.

Mass of potassium chlorate = 12.25 g.

Total mass of produced oxygen and KCI = (4.8 + 7.45)g = 12.25g.

Thus, we can say that the reaction obeys the law of conservation of mass.

Question 6. What amount of sodium chloride will react with 3.41g silver nitrate to produce 1.70g sodium nitrate and 2.87g silver chloride? Consider that the mass is conserved completely in the reaction.

Answer:

1. As the mass is conserved in a chemical reaction, we can say, mass of sodium chloride + mass of silver nitrate = mass of sodium nitrate + mass of silver chloride or, mass of sodium chloride + 3.41g

Therefore, mass of sodium chloride = (1.70 + 2.87 -3.41)g = 1.16 g

Question 7. 6.3g NaHCO₃ when added to 15g acetic acid solution produces 18g of residue. Find the mass of evolved CO₂.

Answer: According to the law of conservation of mass in a chemical reaction, total mass of reactants = total mass of products

Here, total mass of reactants = (6.3 + 15) g = 21.3 g

Hence, 21.3g = mass of residue + mass of evolved CO₂ or, 21.3 g = 18 g + mass of evolved CO₂

Therefore, mass of evolved CO₂ = (21.3- 18) g = 33 g.

Wbbse class 10 chemical calculation

Question 8. At STP, mass of 250 ml of a gas is 0.7317g. If the density of hydrogen at STP is 0.08987g •L^-1, then find the vapour density of the gas. What will be its molecular mass?

Answer: At STP, mass of 250 mL of hydrogen gas = \(\frac{0.08987 \times 250}{1000} \mathrm{~g}\) = 0.0224g

Hence, vapour density of the gas \(\frac{\text { mass of } 250 \mathrm{~mL} \text { of the gas at STP }}{\text { mass of } 250 \mathrm{~mL} \text { hydrogen at STP }}\)   = \(\frac{0.7317 \mathrm{~g}}{0.0224 \mathrm{~g}}\) = 32.66.

Molecular mass of the gas = 2 x vapour density = 2 x 32.66 = 65.32 .

Question 9. At 750mm pressure and 25° C 44.8 temperature, 0.15 g of a gas occupies 45.95mL. Find the standard density,vapour density and molecular mass of the gas.

Answer: Let, the volume of the gas at STP = V₁ mL

p₁ = 760mm, T₁ = 273 K V₂ = 45.95mL, p₂ = 750 mm, T₂= (25 + 273) K = 298K

By combining of Boyle’s law and Charles’ law \(\frac{P_1 V_1}{T_1}=\frac{P_2 V_2}{T_2}\)

∴ \(V_1=\frac{750 \times 45.95 \times 273}{760 \times 298} \mathrm{~mL}\) = 41.54mL

Thus, at STP, 41.54mL of the gas weighs 0.15g.

Hence, at STP, 1000 mL of the gas weights \(\frac{0.15 \times 1000}{41.54} \mathrm{~g}\) = 3.61g.

Therefore, the standard density of the gas = is 361 g. L^-1.

Vapour density = \(\frac{\text { standard density }}{0.089}\) = \(\frac{3.61}{0.089}\) = 40.56.

Molecular mass = 2 x vapour density ^ 2 x 40.56 = 81.12.

Question 10. 44.8 L of oxygen is required to completely convert 24g of a solid element into its oxide at STP. At STP, the gaseous oxide occupies 44.81. Find the mass of the gaseous oxide produced and its vapour density.

Answer: At STP, 22.4 L oxygen weighs 32 g.

Hence, 44.8 L oxygen at STP weights \(\frac{32 \times 44.8}{22.4} \mathrm{~g}\) = 64 g.

Mass of gaseous oxide formed = (24 + 64)g = 88g.

At STP, mass of 44.8 L of the gaseous oxide is 88g.

∴ At STP, 22.4 L of the gaseous oxide weighs \(\frac{88 \times 22.4}{44.8} \mathrm{~g}\) = 44g.

Thus, molecular mass of the gaseous oxide = 44g.

Vapour density =44/2 = 22 .

Wb Class 10 Physical Science Solutions

Question 11. Vapour density of a gas is 8.62. What volume will be occupied by 8.62 g of the gas gat STP?
Answer: Vapour density of the gas = 8.62.

The molecular mass of the gas = 2 x 8.62 = 17.24 g.

Thus, at STP, 17.24 g of the gas occupies 22.4L

Hence, at STP, 8.62g of the gas occupies

= \(\frac{22.4 \times 8.62}{17.4} \mathrm{~L}\) = 11.1L.

Question 12. Vapour density of a gaseous element is 5 times than that of oxygen. If the molecule of the gas Is triatomic, find its atomic mass.

Answer: Vapour density of oxygen (D)= 32/2 = 16

∴ Hence, vapour density of the gas = 5×16 = 80.

Molecular mass of the gas = 2 x 80 = 160.

The gas is triatomic, hence its atomic mass = 160/3= 53.33.

Question 13. At standard temperature and pressure, vapour density of sulphur vapour is 9.143 times the vapourdensity of nitrogen gas. Find the molecular formula of sulphur vapour.

Answer: Vapour density of sulphur vapour with respect to nitrogen gas =\(\frac{\text { molecular mass of sulphur vapour }}{\text { molecular mass of nitrogen }}\)

Hence, molecular mass of sulphur vapour = 9.143 Χ 28 = 256.004.

So, atomicity of sulphur vapour = \(\frac{256.004}{32}\) = 8.

∴ Molecular formula of sulphur vapour = S₈.

Question 14. Vapour density of a gas with respect to air is 1.528. At 27°C and 750 mmHg pressure, find the mass of 2 L of the gas. (vapour density of air with respect to hydrogen is 14.4)

Answer: Vapour density of the gas with respect to air = \(\frac{\text { molecular mass of the gas }}{\text { molecular mass of air }}\)

Hence, molecular mass of the gas = 1.528 x 28.8 = 44g.

If the volume of the gas at STP be V, then \(\frac{750 \times 2}{(273+27)}\) = \(\frac{760 \times V}{273}\)

or, \(V=\frac{750 \times 2 \times 273}{300 \times 760} L\) = 1.796L.

At STP, mass of 22.4 L of the gas = 44 g.
∴ At STP, mass of 1.796 L of the gas = \(\frac{44 \times 1.796}{22.4} \mathrm{~g}\) = 3.528g

Wb Class 10 Physical Science Solutions

Question 15. log calcium carbonate reacts with excess dilute hydrochloric acid to produce CO₂ gas.
Answer:

\(\begin{array}{cc}
\mathrm{CaCO}_3+\mathrm{HCl} & \rightarrow \mathrm{CaCl}_2+\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \\
(40+12+16 \times 3) \mathrm{g} & (12+16 \times 2) \mathrm{g} \\
=100 \mathrm{~g} & =44 \mathrm{~g}
\end{array}\)

1. From the above equation, it is observed that 44 g CO₂ gas is produced from 100 g CaCO_3.

Again, 44g CO₂ = 1 mol CO₂

Thus, we can say , 100g CaCO₃ produces 1 mol CO₂ 10 g CaCO₃ produces = 10/100 Χ 1 mol CO₂  = 0.1 mol CO₂ .

2. At STP, 1 mol CO₂ occupies 22.4 L.

∴ At STP, 0.1 mol CO₂ occupies = 0.1 X 22.4 L = 2.24 L.

Question 16. What amount of calcium carbonate will react with excess dilute HCL to produce 66g CO₂?

Answer: CaCO₃ reacts with dilute HCI to form calcium chloride and CO₂. The reaction can be represented by the following equation,

Wb Class 10 Physical Science Question and Answers

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

⇒ \(\begin{array}{cc}
(40+12+16 \times 3) \mathrm{g} & (12+16 \times 2) \mathrm{g} \\
=100 \mathrm{~g} & =44 \mathrm{~g}
\end{array}\)

Wb Class 10 Physical Science Solutions

Thus, 44g CO₂ is produced from 100 g CaCO₃ . 66g CO₂ is produced from 66 g = 150g CaCO₃ .

Question 17. How much potassium chlorate is required to produce 9.6 g oxygen? [K = 39, Cl = 35.5]

Answer: Oxygen is produced on thermal decomposition of potassium chlorate. The reaction is represented by the following balanced equation,

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

⇒ \(\begin{array}{cc} (40+12+16 \times 3) \mathrm{g} & (12+16 \times 2) \mathrm{g} \\
=100 \mathrm{~g} & =44 \mathrm{~g} \end{array}\)

From the above equation it can be said that, to produce 96g oxygen, the amount of potassium chlorate required is 245g.

∴ To produce 9.6 g oxygen, the amount of potassium chlorate required is= \(\frac{245 \times 9.6}{96} \mathrm{~g}\) = 245 x g

Question 18. What amount of Cu turnings and conc. H₂SO₄ are required to mproduce 64g of SO₂ gas?
Answer: The reaction between copper turnings and sulphuric acid can be represented by the following balanced equation,

Cu\(+2 \mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{CuSO}_4+\mathrm{SO}_2+2 \mathrm{H}_2 \mathrm{O}\)

∴ \(\begin{array}{ccc}
63.5 \mathrm{~g} \quad 2(2+32+16 \times 4) \mathrm{g} & (32+16 \times 2) \mathrm{g} \\
=196 \mathrm{~g} & =64 \mathrm{~g}
\end{array}\)

Wb Class 10 Physical Science Question and Answers

From the above equation we can conclude that, 63.5g copper and 196 g sulphuric acid is required to produce 64g SO₂ gas.

Wb Class 10 Physical Science Solutions

Question 19. What amount of ferrous sulphide is required to produce 1.12L H₂S at STP on reaction with excess dilute H₂SO₄?

Answer: The reaction is represented by,

⇒ \(\begin{aligned}
& \mathrm{FeS}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{FeSO}_4+\mathrm{H}_2 \mathrm{~S} \\
& (56+32) \mathrm{g}=88 \mathrm{~g} \quad 22.4 \mathrm{~L} \text { at STP } \\
&
\end{aligned}\)

Thus, at STP, 22.41 H₂S is produced from 88g FeS

∴ At STP 1.12 L H₂S is produced from \(\frac{1.12 \times 88}{22.4} \mathrm{~g}\)= 4.4g FeS

Question 20. When steam is passed over red hot iron ferrosoferric oxide is formed and hydrogen gas is evolved. To produce 8 mol of hydrogen how many moles of iron is needed?

Answer: The above reaction can be represented as

⇒ \(\underset{3 \mathrm{~mol}}{3 \mathrm{Fe}}+4 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Fe}_3 \mathrm{O}_4+\underset{4 \mathrm{~mol}}{4 \mathrm{H}_2}\)

To produce 4 mol of hydrogen number of moles of iron required is 3.

∴ To produce 8 mol of hydrogen number of moles of iron required is =\(\frac{3 \times 8}{4}\) = 6.

Hence, 6 mol of iron will be required.

Wb Class 10 Physical Science Question and Answers

Question 21. What amount of Zn-dust is required to produce 7g hydrogen? Zn-dust contains 88% zinc [Zn = 65.5].

Answer: The balanced equation of the reaction is

⇒ \(\underset{65.5 \mathrm{~g}}{\mathrm{Zn}}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{ZnSO}_4+\underset{1 \times 2 \mathrm{~g}=2 \mathrm{~g}}{\mathrm{H}_2}\)

2g hydrogen is produced from 65.5g Zn dust.

∴ 7g hydrogen is produced from \(\frac{65.5 \times 7}{2} \mathrm{~g} \mathrm{Zn}\) dust = 229.25g Zn dust.

Now, 88g zinc is present in l00g of Zn-dust.

∴ 229.25g zinc is present in \(\frac{100 \times 229.25}{88} \mathrm{~g}\) = 260.51 g of Zn-dust.

Hence, 260.51g of Zn-dust will be required to produce 7g of hydrogen.

Question 22. 15.25g of a mixture containing KClO₃ and MnO₂ gives off 4.8g oxygen on heating. What amount of MnO₂ was present in the mixture as a catalyst? [K = 3S, Ci = 35.5, O = 16]

Answer: Oxygen is produced on the thermal decomposition of potassium chlorate. The reaction is represented by the following balanced equation,

⇒ \(2 \mathrm{KClO}_3+2 \mathrm{KCl} \rightarrow 3 \mathrm{O}_2\)

⇒ \(\begin{array}{cc}
2(39+35.5+16 \times 3) g & 3(16 \times 2) \mathrm{g} \\
=245 \mathrm{~g} & =96 \mathrm{~g}
\end{array}\)

From the equation it can be said that, to produce 96g of oxygen, the amount of potassium chlorate required is 245g.

∴ To produce 4.8g oxygen, the amount of potassium chlorate required is 4.8g oxygen is produced from 12.25g KCIO₃.

The total mass of the mixture = 15.25 g.

Hence, the amount of MnO₂ present in the mixture = (15.25 -12.25)g = 3.00g.

Wb Class 10 Physical Science Question and Answers

Question 23. When 174g pure MnO₂ is added to excess amount of cone. HCI, a gas is evolved. What amount of l₂ will be released from Kl by the gas produced in this reaction? [Mn = 55,1 = 127]

Answer: MnO₂ reacts with HCI according to the following balanced equation,

4\(\mathrm{HCl}+\mathrm{MnO}_2 \rightarrow \mathrm{MnCl}_2+\mathrm{Cl}_2+2 \mathrm{H}_2 \mathrm{O}\)

55\(+2 \times 16=87 \mathrm{~g} \quad 2 \times 35.5=71 \mathrm{~g}\)

87 g MnO₂ produces 71 g chlorine gas.

∴ 174 g MnO₂ produces \(\frac{174 \times 71}{87} \mathrm{~g}\) = 142g chlorine gas.

Again, \(\begin{aligned}
& \mathrm{Cl}_2+2 \mathrm{Kl} \rightarrow 2 \mathrm{KCl}+\mathrm{I}_2 \\
& 2 \times 35.5=71 \mathrm{~g} \quad 2 \times 127=254 \mathrm{~g} \\
&
\end{aligned}\)

Thus, 71 g Cl₂ gas releases 254g iodine.

Hence, 142g Cl₂ gas releases \(\frac{254 \times 142}{71} g\) = 508g iodine.

Question 24. When excess AgNO₃ solution is added to 1,873 g of a mixture containing KCI and NaCl, 3.731g AgCI is precipitated. Find the amount of KCI in the mixture. [K = 39,Ag = 108, Cl = 35.5 ]

Answer: Let us consider that the mixture contains xg NaCI. Hence, the amount of KCI = (1.873 – x)g.

∴ \(\mathrm{NaCl}+\mathrm{AgNO}_3 \rightarrow \mathrm{NaNO}_3+\mathrm{AgCl}\)

23 \(+35.5=58.5 \mathrm{~g} \quad 108+35.5=143.5 \mathrm{~g}\)

58.5g NaCI precipitates 143.5 g AgCI.

∴ xg NaCI precipitates \(\frac{143.5 \times x}{58.5} \mathrm{~g}\) = 2.453 x g AgCI.

Again, \(\mathrm{KCl}+\mathrm{AgNO}_3 \rightarrow \mathrm{NaNO}_3+\mathrm{AgCl}\)

74.5g KCI precipitates 143.5g AgCI

∴ (1.873 -x)g KCI precipitates \(\frac{143.5(1.873-X)}{74.5} \mathrm{~g}\) = (3.61-1.926x)g AgCI.

According to the problem, 2.453x + 3.61 – 1.926x = 3.731 Or, x = 0.2296 g.

Amount of KCI = (1.873 – 0.2296)g = 1.6434g.

Question 25. When 6.23g of a mixture of magnesium oxide and magnesium carbonate is strongly heated, the mass of the mixture decreases by 2.2g. Find the percent composition of the mixture.

Answer: On heating MgCO₃, it decomposes to MgO and CO₂ , but MgO has no effect on the mass. The produced CO₂ gas is released from the system and hence the mass decreases.

⇒ \(\underset{24+12+3 \times 16=84 \mathrm{~g}}{\mathrm{MgCO}_3} \stackrel{\Delta}{\longrightarrow} \mathrm{MgO}+\mathrm{CO}_2\)

44 g CO₂ is produced from 84g MgCO₃.

∴ 2.2g CO₂ is produced from \(\frac{84 \times 2.2}{44} g\) = 4.2g MgCO₃ .

Thus, the mixture contains 4.2g MgCO₃ .

∴Mass percent of MgCO₃ in the mixture is = \(\frac{4.2}{6.23} \times 100\) = 67.42%.

Hence, the mass per cent of MgO in the mixture =(100-67.42)% = 32.58 %.

Wb Class 10 Physical Science Question and Answers

Question 26. 112g pure iron is converted to ferric chloride and dissolved in water. What volume of H₂S gas at STP is required for the complete conversion of ferric chloride to ferrous chloride? [Fe = 56, Cl = 35.5]

Answer: The reactions involved can be represented by the following equations:

1. \(\underset{2 \times 56=112 \mathrm{~g}}{2 \mathrm{Fe}+3 \mathrm{Cl}_2} \rightarrow \rightarrow_{2(56+3 \times 35.5)=325 \mathrm{~g}}^{2 \mathrm{FeCl}_3}\)

2. \(\begin{aligned}
& 2 \mathrm{FeCl}_3+\mathrm{H}_2 \mathrm{~S} \rightarrow 2 \mathrm{FeCl}_2+2 \mathrm{HCl}+\mathrm{S} \\
& 2 \times(56+35.5 \times 3) \quad 22.4 \mathrm{~L} \\
& =325 \mathrm{~g} \quad \text { (at STP) } \\
&
\end{aligned}\)

Thus, according to the first equation, 112g of pure iron produces 325g of FeCl₃ .

Again, from the second equation it can be concluded that 325g FeCl₃ is completely reduced to  FeCl₂ by 22.4L H₂S gas at STP.

Question 27. The mass of 2.24 L of a gas at NtP is 3.17 g. Calculate its molecular mass, vapour density and standard density.

Answer: At NTP, 2.24 L of the gas has the mass of 3.17 g.

∴ At NTP, 22.4 L of the gas has the mass of \(\frac{3.17 \times 22.4}{2.24} \mathrm{~g}\) = 31.7g

∴ Molecular mass of the gas = 31.7

Vapour density of the gas = 31.7/2 = 15.85

∴ Standard density of the gas = vapour density x 0.089 g. L^-1 = 15.85 x 0.089 g. L^-1 = 1.41 g. L^-1 respectively.

Molecular mass, vapour density and standard density of the gas is 31.7, 15.85 and 1.41 g • L_1 respectively.

Wb Class 10 Physical Science Question and Answers

Question 28. What amount of NaOH will be formed by reacting 23g Na with water at room temperature.

Answer: Balanced equation of the corresponding reaction,

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

⇒ \((2 \times 23) g=46 g\)      \(2(23+16+1) g=80 g\)

According to the equation, 46 g Na produces 80g of NaOH

∴ 23 g Na will produce = \(\frac{80 \times 23}{46}\) = 40g of NaOH.

Question 29. What amount of sulphuric acid is required to neutralise 4 g of NaOH? Or, What amount of sulphuric acid will be needed to neutralise 40g NaOH?

Answer: Balanced chemical equation:

⇒ \(\begin{aligned}
& 2 \mathrm{NaOH}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{Na}_2 \mathrm{SO}_4+2 \mathrm{H}_2 \mathrm{O} \\
& 2(23+16+1) g(2 \times 1+32+4 \times 16) g \\
& =80 \mathrm{~g} \quad=98 \mathrm{~g} \\
&
\end{aligned}\)

Mass of H₂SO₄ required to neutralise 4g NaOH = \(\frac{98 \times 4}{80} \mathrm{~g}\) = 4.9 g

∴ 4.9 H₂SO₄ is needed to neutralise 4g NaOH.

Question 30. CaCl₂, CO₂ and H₂O are formed by the reaction of CaCO₃ with diL HCI. To form 55.5g of CaCl₂, 22.0g CO₂ and 9.0g of H₂O from 50.0 g CaCO₃, what amount (in g) of HCI will be needed? Calculate the number of moles of the required amount of HCI.

Answer: Balanced chemical eqaution:

⇒ \(\begin{gathered}
\mathrm{CaCO}_3+2 \mathrm{HCl} \rightarrow \mathrm{CaCl}_2+\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O} \\
(40+12+3 \times 16) \mathrm{g} 2(1+35.5) \mathrm{g}(40+2 \times 35.5) \mathrm{g}(12+2 \times 16) \mathrm{g}(2 \times 1+16) \mathrm{g} \\
=100 \mathrm{~g}=73 \mathrm{~g} \quad=111 \mathrm{~g}=44 \mathrm{~g} \quad=18 \mathrm{~g}
\end{gathered}\)

According to the equation, 100/2 g or 50g of will form 111/2 or 55.5g of CaCl₂, 44/2 or 22g of CO₂ and 18/2 or 9g of H₂O by reacting with 73/2 g or 36.5g HCL.

Wbbse Class 10 Physical Science Solutions

Question 31. What amount of ammonium sulphate will be required to produce 6.8g ammonia by reaction of caustic soda with ammonium sulphate solution?

Answer: Balanced chemical equation:

(NH₄)₂SO₄ + 2NaOH →Na₂SO₄ + 2NH₃ + 2H₂O
2 x (14+1 x 4)+32+(16 x 4)g    2 x (14 +1 x3)g
= 132g                                          =34g

Now, to form 34g NH₃ amount of (NH₄)₂SO₄ needed = 132g

∴ To form 6.8 g of NH3, amount of (NH₄)₂SO₄ needed = \(\frac{132 \times 6.8}{34} \mathrm{~g}\) = 26.4g

Question 32. If the purity of limestone is 95%, then calculate the amount of CO₂ formed by heating 200 g of the said sample. (Assume, no impurity present in the sample produces CO₂ on heating)

Answer: Balanced chemical equation:

⇒ \(\begin{array}{cc}
\mathrm{CaCO}_3 \stackrel{\Delta}{\longrightarrow} & \mathrm{CaO}+\mathrm{CO}_2 \\
(40+12+16 \times 3) \mathrm{g} & (12+16 \times 2) \mathrm{g} \\
=100 \mathrm{~g} & =44 \mathrm{~g}
\end{array}\)

Now amount of pure CaCO₃ in 200 g of 95% pure lime stone sample = 200 x 95/100g = 190g

∴ According to the equation, the amount of CO₂ formed from 100g CaCO₃ is 44 g.

Amount of CO₂ formed from 190g CaCO₃ is \(\frac{44 \times 190}{100} \mathrm{~g}\)

∴ 83.6g CO₂ will be formed from 200g of the 95% pure limestone sample.

Question 33. What amount of ferrous sulphide will be needed to react with dilute sulphuric acid to form 1.7g hydrogen sulphide gas.
Answer: Balanced chemical equation:

⇒ \(\mathrm{FeS}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{FeSO}_4+\mathrm{H}_2 \mathrm{~S}\)

(56 + 32)g = 88g   (1 Χ 2 +32)g = 34g

Now, to form 34g H₂S, FeS needed = 88g

∴ To form 1.7g H₂S, Fes needed = \(\frac{88 \times 1.7}{34} \mathrm{~g}\) = 4.4g.

∴ 1.7 g of H₂S will be obtained by the reaction between dil.  H₂SO₄ and 44g FeS.

Question 34. When zinc oxide is heated with carbon, metallic zinc and carbon monoxide are formed. What amount of carbon will be needed to form 31.785g Zn and 14.000g CO from 40.685g zinc oxide? How many moles of carbon monoxide will be formed?

Answer: Balanced chemical equation:

∴ \(\mathrm{ZnO}+\mathrm{C} \rightarrow \mathrm{Zn}+\mathrm{CO}\)

Mass of ZnO = 40.685g

Mass of Zn = 14.000g

Mass of CO = 14.000g

∴ Required amount of carbon = (31.785 +14 – 40.685)g + 5.1g

Number of moles of CO = 14/28 = 0.5

Wbbse Class 10 Physical Science Solutions

Question 35. What amount of H₂ will be formed when steam is passed over 21g of red hot Iron? What will be its volume at STP?
Answer: Balanced chemical equation:

3\(\mathrm{Fe}+4 \mathrm{H}_2 \mathrm{O} \rightarrow \mathrm{Fe}_3 \mathrm{O}_4+4 \mathrm{H}_2\)

(3 Χ 56)g = 168g  (4 Χ 2)g= 8g

The amount of H₂ obtained from 168g red hot Fe is 8g

∴ Amount of H₂ obtained from 21g red hot Fe is \(\frac{8 \times 21}{168} g\) = 1g

Again, the volume of 2g H₂ at STP is 22.4L

∴ Volume of 1g H₂ at STP is \(\frac{22.4}{2} \mathrm{~L}\) = 11.2L

Question 36. Calculate the amount of (1) magnesium chloride and (2) hydrogen (in grams) formed by the reaction between 36g magnesium and an excess amount of dil. HCI.
Answer: Balanced Chemical equation: \(\mathrm{Mg}+2 \mathrm{HCl} \rightarrow \mathrm{MgCl}_2+\mathrm{H}_2\)

24g   (24 +2 Χ 35.5)g =95g   2g

1. 24g of Mg forms 95g of MgCl₂

∴ 36g of Mg forms \(\frac{95 \times 36}{24} \mathrm{~g}\) = 142.5g of MgCl₂

2. Again, 24g of MgCl₂ Mg forms 2g of

∴ 142.5 g of and 3g of H₂ will be obtained by the reaction between 36g and an excess amount of dil HCL.

Question 37. What amount of KNO₃ should be heated to obtain oxygen exactly in the same amount that is obtained by heating 49g KClO₃?

Answer: Balanced chemical equation for heating KClO₃:

2\(\mathrm{KClO}_3 \rightarrow 2 \mathrm{KCl}+3 \mathrm{O}_2\)

2 Χ 39 + 35.5 + 16 Χ 3)g = 245g  3 Χ 32g + 96g

According to the reaction, 245g KClO₃ produced 96g of O_2.   

∴ 49g KClO₃ produces \(\frac{96 \times 49}{245} \mathrm{~g}\) = 19.2 g of O_2.

Balanced chemical equation for heating: 

⇒ \(\begin{aligned}
& 2 \mathrm{KNO}_3 \rightarrow 2 \mathrm{KNO}_2+\mathrm{O}_2 \\
& 2 \times(39+14+16 \times 3) \mathrm{g} \quad 32 \mathrm{~g} \\
& =202 \mathrm{~g} \\
&
\end{aligned}\)

According to the reaction, 32g O_2. is obtained by heating 202g of KClO_3.

∴ 19.2 g O₂ is obtained by heating \(\frac{202 \times 19.2}{32} \mathrm{~g}\) = 121.2g of KClO_3.

∴ 121.2g of KNO₃ should be heated to obtain exactly in the same amount that is obtained by heating 49g of KClO_3.

Question 38. What is the percentage of phosphorus in calcium phosphate?
Answer:

Molecular mass of Ca₃(PO₄)₂ = 40 Χ 3 +2 (31 +16 Χ 4) = 310

∴ Amount of phosphours in 100g Ca₃(PO₄)₂ = \(\frac{62 \times 100}{310}\) = 20g

∴ percentage of phosphours in Ca₃(PO₄)₂ is 20%

Question 39. Calculate the amount of magnesium oxide formed by the complete combustion of 24g magnesium.

Answer: Balanced chemical equation:

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

2 Χ 24g = 48g  2 Χ (24+16) g = 80g

According to the equation, amount of MgO obtained from 48g Mg is 80g

∴ = \(\left(\frac{80 \times 24}{48}\right) \mathrm{g}\) = 40g

∴ 40g of MgO will be obtained by the complete combustion of 24g Mg.

Question 40. SO₂ is produced by the combustion of iron pyrites for the industrial preparation of sulphuric acid. The chemical equation is as follows:^4FeS₂ ^{+ 11O}₂^→2Fe₂^O₃ ^{+ 8SO}₂ ^. Calculate the amount of FeS₂ needed to produce 512g SO₂.

Answer: Balanced chemical equation:

4\(\mathrm{FeS} 3+11 \mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3+8 \mathrm{O}_2\)

4 Χ (56 + 32 Χ 2)= 480g      8x(32+16×2) = 512g

To produce 512g SO₂ , 480g of FeS₂ is required.

Question 41. 112g of a metal oxide and a gaseous compound are formed when 200g of a metal carbonate is heated. Vapour density of the gaseous compound is 22.Calculate the number of moles of the gaseous compound formed in the reaction.
Answer:

Mass of the gaseous compound formed in the reaction = (200- 112)g = 88g

Given, vapour density of the gaseous compound = 22

∴ Molecular mass of the gaseous compound = 2 x 22 = 44

∴ Number of moles of the gaseous compound = 88/44 = 2

∴ 2 mol of the gaseous compound will be formed.

Wbbse Class 10 Physical Science Solutions

Question 42. Commercial zinc contains 35% impurity. Calculate the amount of hydrogen gas formed when 50g of such zinc reacts with excess amount of dii H₂SO₄.

Answer: Percentage of impurity = 35%

∴ Percentage of pure zinc = (100- 35)% = 65%

∴ Amount of pure zinc in 50g of impure sample = 65/100 Χ 50g = 32.5g

Now, balanced chemical equation:

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

Now, 65g of pure zinc forms 2g of H₂

∴ 32.5g of pure zinc forms \(\frac{2 \times 32.5}{65} \mathrm{~g}\) = 1 g of H_2.

∴ 1 g hydrogen gas will be formed

Question 43. The oxygen formed by heating a mixture of 75g potassium chlorate and manganese dioxide reacts completely with 3g hydrogen to form water. Calculate the amount of potassium chlorate and manganese dioxide.

Answer: Balanced chemical equation for the formation of water from hydrogen and oxygen is

⇒ \(\begin{gathered}
2 \mathrm{H}_2+\mathrm{O}_2 \rightarrow \mathrm{H}_2 \mathrm{O} \\
(2 \times 2) \mathrm{g}=4 \mathrm{~g}(16 \times 2) \mathrm{g}=32 \mathrm{~g}
\end{gathered}\)

According to the equation,  4g of hydrogen reacts with 32g of O₂

∴ 3g of hydrogen reacts with \(\frac{32 \times 3}{4} \mathrm{~g}\) = 24 g of O_2.

Now, balanced chemical equation for preparation of oxygen from potassium chlorate is

2\(\mathrm{KClO}_3+\left[\mathrm{MnO}_2\right] \rightarrow 2 \mathrm{KCl}+3 \mathrm{O}_2+\left[\mathrm{MnO}_2\right]\)

2(39+35.5+16×3) = 245g   3 Χ 32 = 96g

According to the equation, to produce 96g oxygen, KClO₃ needed

∴ To produce 24g oxygen, KClO_3. needed \(\frac{245 \times 24}{96} \mathrm{~g}\) = 61.25g

∴ Amount of KClO₃ in the mixture = 61.25 g

∴ Amount of MnO₂ in the mixture = (75 – 61.25)g = 13.75g

Question 44. SO₂  is oxidised to form SO₃ by O₂ in presence of catalyst. What amount of SO₂  (in g) will be needed to produce 40g SO₃

Answer: Balanced chemical equation,

⇒ \(\begin{array}{ccc}
2 \mathrm{SO}_2+\mathrm{O}_2 & \rightarrow & 2 \mathrm{SO}_3 \\
2 \times(32+2 \times 16) \mathrm{g} & & 2 \times(32+3 \times 16) \mathrm{g} \\
=128 \mathrm{~g} & & =160 \mathrm{~g}
\end{array}\)

According to the equation, to produce 160g SO₃, SO₂ needed = 128g

∴ To produce 40g SO₃, SO₂ needed = \(\frac{128 \times 40}{160} \mathrm{~g}\) = 32g

∴ 32g of SO₂ will be needed.

Question 45. To convert 25.64g of a solid element A into a gaseous oxide, 11.2L of oxygen (at STP) is needed. Volume of the gaseous oxide (at STP) is 2-1.2L What is the vapour density of the gaseous oxide?
Answer: At STP, the mass of 22.4L oxygen = 32g

∴ Mass of 11.2L oxygen = \(\frac{32 \times 11.2}{22.4} \mathrm{~g}\) = 16g

∴ Mass of the gaseous oxide = (25.64 + 16)g = 41.64g

According to the question, 21.2L of the gaseous oxide weight 41.64g

∴ 22.4L of the gaseous oxide weights \(\frac{41.64 \times 22.4}{21.2} \mathrm{~g}\) = 44g

∴ Molecular mass of the gaseous oxide = 44

∴ Vapour density of the gaseous oxide = 44/2 =22

Wbbse Class 10 Physical Science Solutions

Question 46. What amount of HgO should be heated to produce 0.5 mol oxygen gas?
Answer: Balanced chemical equation of the reaction:

2\(\mathrm{HgO} \rightarrow 2 \mathrm{Hg}+\mathrm{O}_2\)

2 Χ (200 +16)g = 432g

Now according to the equation, 1 mol O₂ is obtained by heating = 432 Χ 0.5g = 216g of HgO

∴ 0.5 mole O₂ is obtained by heating =432 Χ 0.5g = 216g of HgO

∴ 216g HgO should be heated to produce 0.5 mol oxygen gas.

Question 47. What amount of ammonia will be formed by the reaction of 5.35g NH₄CI with an excess amount of lime?
Answer: Balanced equation of the reaction:

2\(\mathrm{NH}_4 \mathrm{Cl}+\mathrm{CaO} \rightarrow \mathrm{CaCl}_2+2 \mathrm{NH}_3+\mathrm{H}_2 \mathrm{O}\)

2(14+1×4+35.5)g V = 107g   2(14+3×1)g = 34g

According to the reaction, 107g NH₄CI produced 34g NH₃   

∴ 5.35g NH₄CI produces \(\frac{34 \times 5.35}{107} \mathrm{~g}\) = 1.7g NH₃ 

∴ 1.7 g of NH₃ will be produced.

Question 48. What amount of Al will be needed (in grams) to produce 558g of Fe by the reduction of Fe₂O₃ by Al at high temperatures? Calculate the number of moles of Fe₂O₃ needed in the reaction.
Answer: Balanced chemical equational of the reaction:

⇒ \(\begin{aligned}
\mathrm{Fe}_2 \mathrm{O}_3+ & 2 \mathrm{Al} \rightarrow 2 \mathrm{Fe}+\mathrm{Al}_2 \mathrm{O}_3 \\
(2 \times 27) \mathrm{g} & (2 \times 55.8) \mathrm{g} \\
=54 \mathrm{~g} & =111.6 \mathrm{~g}
\end{aligned}\)

According to the equation of the reaction,  to produce 111.6g Fe, Al needed =54g

∴ To produce 558g Fe, Al needed = \(\frac{54 \times 558}{111.6} g\) = 270g

∴ 270g of Al will be required.

Again, according to the equation of the reaction, 54g of Al reacts with 1 mol Fe₂O₃

∴ 270g of AI will react with \(\frac{1 \times 270}{54}\) mol Fe₂O₃= 5 mol Fe₂O₃

∴ 5 mol Fe₂O₃ will be required for the reaction.

Question 49. How many grams of H₂SO₄ should react with zinc to form 44.8 litre of H₂ gas at NTP?

Answer: Balanced equation of the reaction—\(\mathrm{Zn}+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{ZnSO}_4+\mathrm{H}_2\)

2 Χ 1+ 32+ 16 Χ 4 = 98g    22.4L

To form 22.4L H₂, H₂SO₄ needed =98g

∴ To form 44.8L H₂, H₂SO₄ needed = \(\frac{98 \times 44.8}{22.4} \mathrm{~g}\) = 196g

Question 50. Upon heating with calcium hydroxide, 32.1g ammonium chloride produces 10.2g NH₃, 33.3g CaCl₂ and 10.8g H₂O. What is the amount of calcium hydroxide (in gram) that participates in the reaction? Calculate the number of moles of NH₃ and the volume of NH₃ at STP, produced in the reaction.

Answer: According to the law of conservation of mass, mass of NH₄CI + mass of Ca(OH)₂ = mass of NH₃ + mass of CaCl₂ + mass of H₂0

∴ 32.1 g + mass of Ca(OH)₂ = 10.2g + 33.3g + 10.8g

∴ Mass of Ca(OH)₂ = (10.2 + 33.3 + 10.8)g- 32.lg = 22.2g

∴ 22.2g of Ca(OH)₂ participates in the reaction. = 14 +1×3 = 17

Wbbse Class 10 Physical Science Solutions

Question 51. lg limestone, in reaction with an excess amount of HCI, produces 0.4g CO₂ gas. Calculate the percentage of purity in the lime stone sample.

Answer: A balanced chemical equation of the reaction is \(\mathrm{CaCO}_3+2 \mathrm{HCl} \rightarrow \mathrm{CaCl}_2+\mathrm{CO}_2+\mathrm{H}_2 \mathrm{O}\)

(40+12+3Χ 16)g=100g  (12+2Χ16)g = 44g

According to the equation, 44g of CO₂ is obtained from 100g of CaCO₃

∴ 0.4g of CO₂ is obtained from \(\frac{100 \times 0.4}{44} \mathrm{~g}\) = 0.909g of CaCO₃

∴ 1 g of CaCO₃ sample contains 0.909g pure CaCO₃

∴ 100g of CaCO₃ sample contains = (0.909 x 100)g = 90.9 g of pure CaCO₃

Question 52. How many grams of CuO can be reduced completely by hydrogen gas obtained from the reaction of 14g of Fe and dii HCI?

Answer: Balanced equation of the chemical reaction is

⇒ \(\underset{56 \mathrm{~g}}{\mathrm{Fe}}+2 \mathrm{HCl} \rightarrow \mathrm{FeCl}_2+\mathrm{H}_2\)

According to the reaction, 56g Fe produces 2g if H₂ in reaction with dil Hcl

∴ 14g Fe produces \(\frac{2 \times 14}{56} g\) = 0.5g H₂ in reaction with dil HCl

Now, the balanced chemical equation of the reduction of CuO is

⇒ \(\underset{(63.5+16) \mathrm{g}=79.5 \mathrm{~g}}{\mathrm{CuO}}+\underset{2 \mathrm{~g}}{\mathrm{H}_2} \rightarrow \mathrm{Cu}+\mathrm{H}_2 \mathrm{O}\)

According to the reaction, 2g H₂ can reduce = \(\frac{79.5 \times 0.5}{2} \mathrm{~g}\) = 19.875g of CuO

∴ H₂ gas obtained from the reaction of 14g Fe and dil. HCL can completely reuce 19.875g of CuO.

Question 53. What amount (in gram) of cupric oxide will be reduced by 3.4g NH₃? What will be the volume of the gas produced in this case at NTP?
Answer: Balanced chemical equation of the reaction is—

3\(\mathrm{CuO}+2 \mathrm{NH}_3 \rightarrow 3 \mathrm{Cu}+\mathrm{N}_2+3 \mathrm{H}_2 \mathrm{O}\)

3x(63.5+16)g = 238.5g  2x(14+3)g= 238.5g = 34g  22.4L

∴ 34g NH₃can reduce 238.5g CuO

∴ 3.4g NH3 can reduce \(\frac{238.5 \times 3.4}{24} \mathrm{~g}\) = 23.85 g CuO

Again N₂ gas is formed during the reaction, according to the reaction

34g NH₃ produces 22.4L N₂ at NTP

∴ 3.4g NH₃ produces \(\frac{22.4 \times 3.4}{34} L\) = 2.24 L N₂ at NTP.

Question 54. What amount of Fe₂O₃ will be needed to produce 22.4l Of oxygen gas according to the following reaction? \(\mathrm{Fe}_2 \mathrm{O}_3 \rightarrow \mathrm{Fe}+\mathrm{O}_2\)

Answer: Balanced chemical equation is

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

2 Χ (2 Χ 56 + 3 Χ 16)g = 320g   3 Χ 22.4L(at STP)

According to the reaction,  3 Χ 22.4L O₂ will be produced by 320g Fe₂O₃

∴ 22.4L O₂ will be produced by \(\frac{320 \times 22.4}{3 \times 22.4} \mathrm{~g}\)

= 106.67g Fe₂O₃

∴ 106.67g Fe₂O₃ of will be needed.

Chapter 3 Chemical Calculations Miscellaneous Type Questions Match The Column:

Question 1.

Column A	Column B
1. Mass-energy equivalence	1. Vapour density
2. Calculation of relative quantities of reactants and products	2. Lavoisier
3. Conservation of mass	3. E=mc2
4. The quantity whose value is half the molecular mass of a gas	4. Stoichiometry

 

Answer:

1. Mass-energy equivalence: 3. E=mc²
2. Calculation of relative quantities of reactants and products: 4. Stoichiometry
3. Conservation of mass: 2. Lavoisier
4. The quantity whose value is half the molecular mass of a gas: 1. Vapour density

Question 2.

Column A	Column B
1. Amount of CO2 produced on heating 1 mol CaCO3	1. 32g
2. 0.01 mol C2H4	2. 71g
3. 22.4L Cl2(g) at STP	3. 0.28g
4. Mass of 6.022 X 1023 O2 molecules	4. 44 g

 

Answer:

1. Amount of CO₂ produced on heating 1 mol CaCO₃: 4. 44 g
2. 0.01 mol C₂H₄: 3. 0.28g
3. 22.4L Cl₂(g) at STP: 2. 71g
4. Mass of 6.022 X 10²³ O₂ molecules: 1. 32g

WBBSE Solutions for Class 10 Physical Science and Environment

Chapter 1 Environmental Concern Topic A Atmosphere Synopsis:

Atmosphere: The gaseous layer surrounding the surface of the earth that extends up to about 1600 km, is known as the atmosphere.

Different layers of the atmosphere: On the basis of temperature and pressure, the atmosphere is divided into five different layers namely troposphere, stratosphere, mesosphere, thermosphere, and exosphere.

Convection current: On heating, fluids (liquids and gases) undergo expansion, and hence their density decreases. So, the heated liquid or gas becomes lighter and moves upward. On the other hand, the cool heavy part of the upper region comes down. This results in the formation of circular current which is known as convection current.

Read and Learn Also WBBSE Solutions for Class 10 Physical Science and Environment

Wind and storm: Wind blows from a region of comparatively higher pressure to a region of comparatively lower pressure. Earth surface, under certain circumstances, gets heated up. Hence the air adjacent to this earth surface also gets heated up and moves upward.

This creates a low pressure zone in that, region. Cooler air from surroundings rushes towards that region. Greater the difference in air pressure, faster the movement of air from the higher to lower pressure region. Thus wind is created. When wind blows above a specified speed, it is termed as a storm.

Chapter 1 Environmental Concern Topic A Atmosphere Short And Long Answer Type Questions:

Question 1. What is the atmosphere? What are the different regions of the atmosphere?
Answer:

Atmosphere:

The gaseous layer surrounding the surface of the earth which extends up to about 1600 km, is known as the atmosphere. Under the influence of gravity, this surrounding remains attached to the surface of the earth. On the basis of height and temperature, the atmosphere is divided into five different regions.

Layers	Height from the earth’s surface	Temperature Range (°C)
Tropo- sphere	0- 12 km	+15 to -60 (temperature increases with the. increase in height)
Strato- sphere	12 – 45 km	-60 to 0 (temperature increases with the increase of height)

 

Layers	Height from the earth’s surface	Temperature Range (°c)
Meso- sphere	45-85 km	0 to -100 (temperature increases with the increase in height)
Thermo- sphere	85 – 500 km	-100 to +1200 (temperature increases with the increase in height)
Exo- sphere	500 -1000 km	> 1200

 

Wb Class 9 Physical Science Question and Answers

Question 2. What are the ‘homosphere’ and ‘heterosphere’ of the atmosphere?
Answer:

The ‘homosphere’ and ‘heterosphere’ of the atmosphere:

The lower part of the atmosphere (extends up to an approximate height of 85 km above the earth surface) is homogeneous in nature. The gaseous components ( N_2, O_2, Ar, CO_2, water vapor etc.) are evenly distributed in this, region. This part of the atmosphere is called the homosphere.

The remaining part of the atmosphere above the homosphere where the components are unevenly distributed is called the heterosphere. This part starts just above the homosphere region and extends between 85km above the earth’s surface up to a distance of 1000 km.

Question 3. Define troposphere. Why is it called turbulent sphere’?
Answer:

Troposphere:

The lowest layer of the atmosphere is called the troposphere. It extends from the sea level up to a height of about 12 km. This layer contains dust particles, water vapour, clouds etc. Different natural phenomena like storms, rains, lightning, thunderstorms etc., occur in this region. So, it is also known as the turbulent sphere.

Question  4. What is the stratosphere? Why is it dynamically stable?
Answer:

Stratosphere:

The atmospheric layer that extends up to a height of 45 km above the troposphere is known as stratosphere.

There is very little air in this region and dust particles, water vapor etc. are absent in this layer. Hence, this layer is free of several associated turbulence like clouds, rains, lightning, thunderstorms etc. Thus, it is dynamically stable.

Environment Solutions

Question 5. What is the mesosphere? Name the top layer of the mesosphere.
Answer:

Mesosphere:

The atmospheric layer present just above the stratosphere and extends up to a height of 85 km below the thermosphere is known as the mesosphere. The top layer of the mesosphere is known as mesopause. Temperature remains fixed (-92°C) in this region.

Question 6. What is the thermosphere? Why is it named so?
Answer:

Thermosphere:

The atmospheric layer present just above the mesosphere and extends up to a height of 500 km is known as thermosphere.

The temperature of this layer tends to increase abruptly with increasing altitude. At a height of about 120 km the temperature is about 500°C, at a height of 200 km the temperature becomes almost 700°C and at 480 km the temperature is about 1232°C. Therefore, this layer is called thermosphere.

Wb Class 9 Physical Science Question and Answers

Question 7. What is ionosphere? Why Is it named so?
Answer:

Ionosphere:

The particular portion of thermosphere which contains several gaseous ions is known as ionosphere.

Cosmic radiation, X-rays, and gamma rays coming from the Sun ionizes the gases (mainly nitrogen and oxygen) of this region. As a result, this particular region contains large number of ions and free electrons. Due to the high concentration of ions in this region, this is called ionosphere.

Question 8. What is meant by ‘aurora’? What are ‘aurora Borealis and ‘aurora australis’?
Answer:

‘Aurora’:

Cosmic radiations, X-rays, gamma rays coming from the sun ionizes the gases present in ionosphere (a specific portion of thermosphere) and produce large number of ions as well as free electrons.

These charged particles undergo interactions with the magnetic field of the earth. As a result of this interactions, bright lights appear in the sky of polar regions. These polar lights or natural light displays in the earth’s sky is commonly known as aurora.

The aurora formed in the sky of north pole is known as aurora borealis and the aurora formed in the sky of south pole is called aurora australis.

Question 9. Give a brief description of the exosphere layer of atmosphere.
Answer:

The exosphere layer of atmosphere:

The atmospheric layer situated at a height of more than 500 km with respect to the surface of earth is called exosphere. It extends up to an altitude of about 1000 km. The temperature in this region is more than 1200°C. This layer contains hydrogen and. helium gases. Artificial satellites and space stations are located at this layer.

Wb Class 9 Physical Science Question and Answers

Question 10. How does atmospheric pressure change with the increase in altitude?
Answer:

As we move up from the sea level, density of air decreases, and hence, atmospheric pressure also decreases.

Generally, with every 110 m rise in height, air pressure decreases by 1cm. However, the rate of change of pressure is not uniform at the upper layer of the atmosphere.

Question 11. Discuss how temperature change in different layers of the atmosphere as the distance from the earth’s surface increases.
Answer:
1. The troposphere gets colder with the increase in altitude. With each km rise in altitude, the temperature falls by 6.5°C. The top layer of troposphere has a temperature of about -56°C.

2. With increasing altitude, the temperature of stratosphere gradually increases and temperature becomes 0°C at the top.

3. Mesosphere is the coldest region of atmosphere. Temperature decreases as we move upward along this layer.

4. In the thermosphere, temperature gradually increases with increasing height. This layer absorbs cosmic radiation coming from the Sun and is heated up. ConseQuestionuently, its temperature reaches up to 1200°C. The temperature of exosphere is higher than 1200°C.

Wb Class 9 Physical Science Question and Answers

Question 12. Why does temperature increase with increasing height at the stratosphere?
Answer:

The oxygen molecules present in stratosphere is dissociated into atomic oxygen by absorbing the UV rays of the sun. These oxygen atoms then combine with molecular oxygen to form ozone (O₃) molecules.

The last step is highly exothermic in nature and produces a large amount of heat. As a result, the temperature of this layer increases gradually with increase in height and reaches at a temperature around 0°C at the top of the layer.

Question 13. What is convection? What is convection current?
Answer:

Convection: The process by which the heated molecules of a liquid or a gas move from a region of higher temperature to a region of lower temperature thus carrying heat energy, is called convection.

Convection current: When a liquid or gas is heated, its density decreases due to expansion The heated liquid or gas thus becomes lighter and moves upward. Comparatively colder and heavier liQuestionuid or gas from the surroundings then rushes towards that region to fill up the vacuum. Such circulation of gaseous or liQuestionuid layer results in the formation of a convection current.

Question 14. How is convection current formed in air?
Answer: If the earth’s surface becomes hot the air adjacent to the earth’s surface also gets heated up. Hence it becomes lighter and rises up. This creates a low-pressure zone in that region. Cold air from surroundings then rushes towards that region to fill up the vacuum. This creates a convection current.

Question 15. Give two natural phenomena where a convection current is observed. When do they have maximum intensity?
Answer:

Convection current is observed in land breeze and sea breeze. The inland breeze blows during the night and intensifies around the dawn. On the other hand, sea breeze blows at the morning and intensifies around dusk.

Question 16. Describe the formation of land breeze.
Answer:

The formation of land breeze:

At night, the landmass in the coastal area loses heat at a faster rate than the sea water. Hence, the water remains comparatively warmer at night. ConseQuestionuently, warm air above the sea becomes lighter and rises up while cold and dense air from the landmass starts moving towards the sea. This is known as land breeze.

Question 17. Describe the formation of the sea breeze.
Answer:

The formation of the sea breeze:

The specific heat of water is high, due to this the land in the coastal areas get heated up in a Questionuicker manner compared to the seawater. As a result, the air above the landmass becomes lighter and rises up, thereby creating a low-pressure zone. The cold, denser air from the sea then flies into this region resulting in the sea breeze.

Question 18. How is a storm formed?
Answer:

Formation Of Storm

Wind blows from a higher-pressure region to a lower pressure region. The earth surface under certain circumstances get heated up so Questionuickly that the air adjacent to the earth’s surface becomes warmer and lighter. Hence it goes upward creating a vacuum. As a result a low pressure zone is created in that region.

The air from the surroundings then rushes towards that region. Greater the difference in air pressure, the faster the air blows from higher to the lower-pressure region. When wind blows above a specified speed, it is termed as a storm.

Chapter 1 Environmental Topic A Atmosphere Very Short Answer Type Questions Choose The Correct Answer:

Question 1. The layer of the atmosphere having the highest density is
1. Troposphere
2. Stratosphere
3. Mesosphere
4. Thermosphere

Answer: 1. Troposphere

Question 2. The layer of atmosphere in which natural phenomena such as storms, rains, and lightning occur is
1. Thermosphere
2. Troposphere
3. Stratosphere
4. Mesosphere

Answer: 2. Troposphere

Question 3. In the atmosphere, the percentage (by mass) occupied by the troposphere is
1. 25
2. 50
3. 75
4. 60

Answer: 3. 75

Question 4. The layer which is characterized by its calmness is
1. Troposphere
2. Stratosphere
3. Mesosphere
4. Thermosphere

Answer: 2. Stratosphere

Question 5. The temperature at the topmost level of the stratosphere is
1. 50°C
2. -50°C
3. 0°C
4. -30°C

Answer: 3. 0°C

Question 6. The atmospheric layer that absorbs ultraviolet rays is
1. Troposphere
2. Ozonosphere
3. Thermosphere
4. Mesosphere

Answer: 2. Ozonosphere

Wbbse Class 10 Physical Science And Environment

Question 7. The hottest layer of the atmosphere is
1. Exosphere
2. Ozonosphere
3. Ionosphere
4. Thermosphere

Answer: 1. Exosphere

Question 8. In which atmospheric region does aurora form?
1. Ozonosphere
2. Ionosphere
3. Tropopause
4. Mesosphere

Answer: 2. Ionosphere

Question 9. The atmospheric layers where temperature increases with increasing altitude are
1. Troposphere and thermosphere
2. Mesosphere and stratosphere
3. Thermosphere and mesosphere
4. Stratosphere and thermosphere

Answer: 4. Stratosphere and thermosphere

Question 10. The atmospheric layer in which the radio waves get reflected is
1. Stratopause
2. Ozonosphere
3. Ionosphere
4. Mesosphere

Answer: 3. Ionosphere

Question 11. For each kilometer rise in altitude from the earth’s surface, the troposphere gets colder by
1. 5.6°C
2. 6.5°C
3. 3.5°C
4. 4.6°C

Answer: 2. 6.5°C

Question 12. Artificial satellites and space stations are located at the
1. Eexosphere
2. Thermosphere
3. Mesosphere
4. Stratosphere

Answer: 1. Eexosphere

Wbbse Class 10 Physical Science And Environment

Question 13. Which layer of the atmosphere controls temperature and water cycle?
1. Stratosphere
2. Thermosphere
3. Troposphere
4. Mesosphere

Answer: 3. Troposphere

Question 14. The intensity of land breeze increases
1. Around dawn
2. Around noon
3. Around dusk
4. At night

Answer: 1. Around dawn

Question 15. The layer of the atmosphere that can be called as ‘Natural Solar Screen’ is
1. Troposphere
2. Ozonosphere
3. Thermosphere
4. Exosphere

Answer: 2. Ozonosphere

Wbbse Class 10 Physical Science And Environment

Question 16. Several ions are present in
1. Troposphere, stratosphere
2. Troposphere, mesosphere
3. Mesosphere, thermosphere
4. Stratosphere, exosphere

Answer: 3. Mesosphere, thermosphere

Question 17. The average height of the troposphere from earth’s crust
1. 18 km
2. 15 km
3. 12 km
4. 20 km

Answer: 3. 12 km

Question 18. In which layer of the atmosphere the amount of water vapor is the highest
1. Troposphere
2. Stratosphere
3. Thermosphere
4. Mesosphere

Answer: 1. Troposphere

Question 19. X-ray is absorbed in the layer called
1. Troposphere
2. Mesosphere
3. Ionosphere
4. Exosphere

Answer: 3. Ionosphere

Question 20. Gases that are present in the exosphere are
1. He, H₂
2. N_2, O₂
3. N₂, H2
4. H₂, O₂

Answer: 1. He, H₂

Question 21. The coldest region of the atmosphere is
1. Stratosphere
2. Exosphere
3. Mesosphere
4. Thermosphere

Answer: 3. Mesosphere

Chapter 1 Environmental Concern Topic A Atmosphere Answer In Brief

Question 1. State the range of troposphere in the polar region and eQuestionuatorial region.
Answer: In the polar region, troposphere extends from sea level up to a height of 8-9 km and in the eQuestionuatorial region it extends up to 16-18 km.

Question 2. What is tropopause?
Answer:

Tropopause

The narrow region, at a height of 12km from the earth’s surface, where troposphere meets stratosphere is known as the tropopause. In this region, the temperature remains unaltered with the increase in height.

Question 3. What is the air pressure at the lowest part of the troposphere?
Answer: The air pressure at the lowest part of the troposphere is eQuestionual to the pressure of 76 cm Hg column.

Question 4. What is stratopause?
Answer:

Stratopause

The thin layer between the stratosphere and mesosphere where. the change in temperature becomes almost negligible is known as the stratopause.

Question 5. Which part of the atmosphere is the coldest?
Answer: Mesosphere.

Question 6. Which gases are predominantly found in thermo sphere?
Answer: Nitrogen and oxygen.

Wbbse Class 10 Physical Science And Environment

Question 7. The temperature gradually decreases at a definite rate as we move upward in the atmosphere from the earth’s surface. What is this phenomenon called?
Answer: Normal lapse rate.

Question 8. What is the temperature at the upper end of the troposphere?
Answer: At the upper end of the troposphere, the temperature ranges between -56°C to -60°C.

Question 9. Give examples of local winds.
Answer: Land breeze and sea breeze.

Question 10. Name the lowest layer of atmosphere.
Answer: Troposphere

Question 11. In which layer of the atmosphere do storms and rain occur?
Answer: Troposphere.

Question 12. The living world of the earth’s crust remains in direct contact with which layer of the atmosphere?
Answer: Troposphere.

Question 13. In which layer of the atmosphere does the strongly moving ‘Jet stream’ flow?
Answer: The ‘Jet Stream’ flows in the tropopause region.

Question 14. Through which layer of the atmosphere can a jet plane fly?
Answer: Jet planes can fly through stratosphere.

Question 15. How does the temperature change with an increase in height of the stratosphere?
Answer: The temperature increases with increase in height in the stratosphere.

Question 16. Name some gaseous components present in stratosphere.
Answer: Nitrogen, oxygen, and ozone.

Question 17. What is the temperature of the mesopause part of the mesosphere?
Answer: -92°C.

Question 18. Aurora Borealis is found in which region of the atmosphere?
Answer: Ionosphere region of thermosphere.

Wbbse Class 10 Physical Science And Environment

Question 19. Mention some components of the mesosphere and thermosphere.
Answer:
Some components of the mesosphere: N_2, O_2,
Some components of thermosphere: of, O ^⊕ NO ^⊕

Question 20. Which layer is used for communication with the help of radio waves?
Answer: Ionosphere of thermosphere .

Question 21 Mention the maximum temperature of thermosphere.
Answer: Almost 1200°C.

Question 22. Which layer of the atmosphere is suitable for the installation of artificial satellites?
Answer: Exosphere.

Chapter 1 Environmental Concern Topic A Atmosphere Fill In The Blanks:

Question 1. The _________ is the most dense layer of the atmosphere.
Answer: Troposphere

Question 2. The layer separating troposphere stratosphere is called ____________.
Answer: Tropopause

Question 3. When a meteorite enters the earth’s atmosphere, it burns due to it’s collision with the air of __________ layer.
Answer: The Mesosphere

Wbbse Class 10 Physical Science And Environment

Question 4. The lower part of the ___________ ionosphere.
Answer: Thermosphere

Question 5. Due to the presence of excess amount of _________ gas in the stratosphere, the temperature of this region rises with altitude.
Answer: Ozone

Question 6. As the distance from the earth surface increases, the __________ of air decreases.
Answer: Density

Question 7. _______ breeze helps in sailing ships offshore.
Answer: Land

Question 8. The term troposphere was first used by the scientist 
Answer: Teisserenc de Bort.

Question 9. The word troposphere means __________ region.
Answer: Turbulent

Question 10. The temperature of the troposphere ___________ with increases in height.
Answer: Turbulent

Question 11. The air layer above the troposphere is called ____________
Answer: Stratosphere

Question 12. The temperature of stratosphere _______ with increase in height.
Answer: Increases

Question 13. The layer of the atmosphere in which Aurora can be seen is _______
Answer: Thermosphere

Question 14. The effect of charged molecules in the ionosphere of atmosphere results in the formation of _________
Answer: Aurora

Question 15. The water cycle of the earth is controlled by ________
Answer: Troposphere

Chapter 1 Environmental Concern Topic A Atmosphere State Whether True Or False:

Question 1. In troposphere, with increase in height, the pressure decreases.
Answer: True

Question 2. Stratosphere is the most dense layer of the atmosphere.
Answer: False

Question 3. Convection currents are formed in solid, liquid, and gaseous states of matter.
Answer: False

Question 4. Heterosphere starts from a height of 85 km. relative to the surface.
Answer: True

Wbbse Class 10 Physical Science And Environment

Question 5. Air pressure decreases with increase in height in troposphere.
Answer: True

Question 6. Water vapour, clouds, etc. do not exist in troposphere.
Answer: False

Question 7. Due to the presence of dust particles in troposphere, the sky looks blue.
Answer: True

Question 8. Aurora can be seen in ozonosphere.
Answer: False

Question 9. The maximum temperature of thermosphere is around 1200°C.
Answer: True

Question 10 Atmospheric pressure in Darjeeling is greater than that in Kolkata.
Answer: False

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming Synopsis

Ozonosphere: The harmful UV rays of the Sun are absorbed by the ozone layer or ozonosphere present in the stratosphere. In this way, the ozone layer protects living organisms from their harmful effects.

Ozone hole: Increased use of chemicals and technology for the development of civilization has adverse effects on the ozone layer. Due to the constant decomposition of ozone molecules, ozone holes are found to be formed at different parts of the ozone layer.

Compounds responsible for ozone layer depletion: Chlorofluorocarbons (CFCs) and oxides of nitrogen (mainly NO) are mainly responsible for the depletion of the ozone layer.

Greenhouse effect: Gases like CO₂, CH₄, N₂O, CFC, and water vapor traps the retransmitted infrared radiations of higher wavelengths. This helps to keep the earth’s surface warm and thus a favorable environment is created for the survival of living beings. This is called the greenhouse effect.

Global warming: Gradual increase in the average temperature of the earth’s atmosphere . due to several natural phenomena as well as man-made activities is called global warming. Greenhouse effect plays an important role to create global warming.

Causes of global warming: Unlimited use of fossil fuels, emission of harmful gases by industries, and deforestation enhance the amount of several greenhouse gases like CO₂, CH₄, CFC, etc. in the atmosphere. The increased amount of such greenhouse gases is one of the main reasons of global warming.

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming Short And Long Answer Type Questions

Question 1. What is ozonosphere? Why is this layer so essential for us?
Answer:

Ozonosphere:

Stratosphere contains an ozone gas layer which is extended from 16 km to 30 km with respect to the earth’s surface and the concentration of ozone gas is maximum in this layer. This particular layer is called ozonosphere.

The harmful UV rays of the Sun are absorbed by ozone layer. The ozone layer in the stratosphere prevents the UV rays from reaching the earth’s surface and thus, protects living organisms from its harmful effects. That is why ozonosphere is so essential.

Question 2. What is ozone hole?
Answer:

Ozone hole:

Excessive use of different chemicals and technological devices for the development of human civilization leads to a gradual depletion of the ozone layer. The rate of decomposition of ozone is much higher than the rate of its formation of it. As a result, hole is formed in certain regions of the ozone layer which is known as the ozone hole.

Question 3. Name some man-made chemical substances which are responsible for ozone layer depletion.
Answer: Chlorofluorocarbons (CFCs) and some oxides of nitrogen such as nitric oxide (NO) and nitrogen dioxide NO₂ are mainly responsible for depletion of the ozone layer.

Question 4. What is the unit used to measure concentration of ozonosphere? Define the unit.
Answer: The concentration of ozone layer is measured in Dobson unit. At 0°C temperature and 760 mm of Hg pressure, the density of 0.01 mm thick ozone gas layer is called 1 Dobson Unit (1 DU).

Question 5. Describe the significance of ozone layer on the environment.
Answer:

The significance of ozone layer on the environment:

Ozone layer in stratosphere acts as a protective shield for the earth. If the ozone layer was not present in the stratosphere, the ultraviolet rays would reach the earth surface.

As a result- the overall temperature of the earth surface would increase. The increased temperature would lead to extinction of plants and animals on the earth. direct exposure to ultraviolet rays may cause a number of diseases including skin cancer, premature cataract etc.

Wbbse Class 10 Physical Science And Environment

Question 6. How is ozone layer formed?
Answer: UV-C (wavelength: 100-280 nm) along with some UV-B. (wavelength: 280-315 nm) decomposes oxygen molecules present in the stratosphere into atomic oxygen. These oxygen atoms then combine with molecular oxygen to form ozone (O3) molecules. The last step is highly exothermic in nature, i.e., large amount of heat is produced in this step.

\(\mathrm{O}_2 \stackrel{(\text { UV-C or UV-B) }}{\longrightarrow} \mathrm{O}+\mathrm{O}\) \(\mathrm{O}_2+\mathrm{O} \longrightarrow \mathrm{O}_3\)

Question 7. Discuss the natural causes that decompose the ozone present in ozone- sphere. How is the total amount of ozone gas maintained in the ozone layer?
Answer: Dissociation of ozone molecules in the stratosphere occurs simultaneously with its formation. Ultraviolet rays of larger wavelength (UV-A; wavelength: 315-400 nm) decompose ozone molecules into molecular oxygen and atomic oxygen.

\(\mathrm{O}_3 \stackrel{\mathrm{UV}-\mathrm{A}}{\longrightarrow} \mathrm{O}_2+\mathrm{O}\)

In stratosphere, formation of ozone molecules and decomposition of ozone molecules occur simultaneously in a cyclic manner and a dynamic equilibrium is maintained between these two processes. As a result, the amount of ozone gas in the stratosphere remains constant.

Question 8. Ozone gas present in stratosphere is important but presence of ozone gas in to- posphere is harmful-justify the statement.
Answer: The harmful UV rays of the Sun are absorbed by ozone layer present in the stratosphere. Hence, this stratospheric ozone is protecting the earth from UV rays and their harmful effects as a natural umbrella.

On the other hand, ozone gas present in troposphere acts as a greenhouse gas (it contributes about 7-8% to the greenhouse effect). The greenhouse effect plays an important role in the gradual increase of the average temperature of the earth.

Question 9. How is the dynamic equilibrium maintained in the stratosphere during the formation and depletion of the ozone layer?
Answer:
Formation of ozone layer: Oxygen molecules. present in the stratosphere absorbs UV rays and are dissociated to form oxygen atoms.

\(\mathrm{O}_2+\text { UV-rays } \rightarrow \mathrm{O}+\mathrm{O}\)

This atomic oxygen combines with molecular oxygen to form ozone molecules-

Wbbse Class 10 Physical Science And Environment

\(\mathrm{O}_2+\mathrm{O} \rightarrow \mathrm{O}_3\)

Depletion of ozone layer: Ozone decomposes to 0.12 Discuss the harmful effects of ozone oxygen molecules by absorbing UV rays.

\(\mathrm{O}_3+\text { UV-ray } \rightarrow \mathrm{O}_2+\mathrm{O} ; \mathrm{O}_3+\mathrm{O} \rightarrow 2 \mathrm{O}_2\)

Formation of ozone molecules and decomposition of ozone molecules-these two opposite processes continue in cycles at the ozone layer and an equilibrium is established. It is for this equilibrium that the amount of ozone gas in the stratosphere remains constant.

Question 10. Discuss the role of chlorofluorocarbons (freons) in the depletion of ozone layer.
Answer:

The role of chlorofluorocarbons (freons) in the depletion of ozone layer:

Chlorofluorocarbons (CFCs) act as catalysts in the depletion of ozone layer in the stratosphere. In presence of ultraviolet rays, CFCs dissociate to form active chlorine atoms (CI).

Active chlorine atom reacts with ozone molecule to form oxygen and chlorine monoxide. Chlorine monoxide further reacts with ozone molecules to form oxygen and active chlorine. This process continues in a cyclic manner. As a result, the density of ozone layer decreases.

\(\mathrm{CFCl}_3 \stackrel{\mathrm{UV} \text {-ray }}{\longrightarrow} \dot{\mathrm{C}} \mathrm{FCl}_2+\dot{\mathrm{C}}\) \(\dot{\mathrm{Cl}}+\mathrm{O}_3 \longrightarrow \dot{\mathrm{C}} \mathrm{OO}+\mathrm{O}_2\) \(\dot{\mathrm{ClO}}+\mathrm{O}_3 \longrightarrow 2 \mathrm{O}_2+\dot{\mathrm{C}}\)

Question 11. How are supersonic jet planes responsible for depletion of ozone layer?
Answer: Supersonic jet planes flying through the stratosphere emits large amount of nitric oxide. Nitric oxide (NO) reacts with ozone to form nitrogen dioxide (NO₂) and oxygen.

On the other hand in stratosphere, ozone molecules are decomposed by the UV rays to produce atomic oxygen, which reacts with NO₂ to produce back NO. Thus, the amount of NO never decreases and causes gradual destruction of ozone molecules. As a result depletion of ozone layer takes place.

\(\mathrm{NO}+\mathrm{O}_3 \longrightarrow \mathrm{NO}_2+\mathrm{O}_2\) \(\mathrm{O}_3 \stackrel{\mathrm{UV}}{\longrightarrow} \mathrm{O}_2+\mathrm{O}\) \(\mathrm{NO}_2+\mathrm{O} \longrightarrow \mathrm{NO}+\mathrm{O}_2\)

Question 12. Discuss the harmful effects of ozone layer depletion.
Answer:

The harmful effects of ozone layer depletion:

1. Ozone layer in the atmosphere absorbs harmful UV rays which would otherwise, reach the earth’s surface. This would cause an increase in the earth’s temperature resulting in melting of snow-caps in polar regions. Thus the water level of the seas and oceans would increase and flood the coastal regions.

Exposure to UV rays may cause skin cancer and premature cataract in the eye, and natural immunity also decreases. These rays are
responsible for formation of photochemical Ozone layer in the atmosphere absorbs. harmful UV rays which would otherwise, reach the earth’s surface.

This would cause an increase in the earth’s temperature resulting in the melting of snow-caps in polar regions. Thus the water level of the seas and oceans would increase and flood the coastal regions.

2 Exposure to UV rays may cause skin cancer and premature cataract in the eye, and natural immunity also decreases. These rays are responsible for the formation of photochemical smog.

3 UV rays inhibit the process of photosynthesis in plants, hence production of crop gets reduced.

Question 13. What are greenhouse gases?
Answer:

Greenhouse gases:

Some gases present in the atmosphere absorb radiations of comparatively greater wavelengths (re-radiated from the earth surface) and reflects back the remaining radiations, thus they keep the earth’s atmosphere warm. These gases are known as greenhouse gases.

Some examples of greenhouse gases are carbon dioxide (CO₂), methane ( CH₄), chlorofluorocarbon (CFC), ozone (O₃), nitrous oxide (N₂O), water vapor etc.

Wbbse Class 10 Physical Science And Environment

Question 14. What is greenhouse effect?
Answer:

Greenhouse effect:

Infrared (IR) light coming from the Sun accounts for the heating of earth. As it has a smaller wavelength the gases present in the atmosphere are unable to absorb this incident radiation. The IR light re-radiated from the earth surface is of longer wavelengths.

Greenhouse gases like CO₂, CH₄, CFC, N₂O, O₃ etc. absorb these re-radiated wavelengths and thus prevent the heat of the Sun to escape from the earth’s atmosphere. As a result the earth’s atmosphere stays warm and thus a favourable environment is created for the survival of living beings. This phenomenon is known as greenhouse effect.

Question 15. Discuss the importance of greenhouse effect.
Answer:

The importance of greenhouse effect:

Greenhouse gases prevent the reflected radiations to escape from the earth. As a result, the earth’s atmosphere stays warm and thus a suitable environment is created for the survival of living beigns.

If there were no greenhouse gases in the atmosphere, there would have been no greenhouse effect and the earth’s average temperature would fall to -30°C.

Question 16. How do greenhouse gases increase the temperature of the atmosphere?
Answers: IR rays (of shorter wavelengths) of the Sun can easily penetrate the lining of the greenhouse gases (CO₂, CH₄,O₃, N₂O, CFC etc.) present in the atmosphere. The earth’s surface absorbs some part of this radiation and becomes heated. Remaining part of this radiation is retransmitted.

IR rays radiated from the heated earth’s surface is of greater wavelengths and got absorbed by the screen of greenhouse gases. Thus the layer of greenhouse gases maintains the warmth of the atmosphere within a certain range by trapping the retransmitted IR rays. But the amount of greenhouse gases in the atmosphere is gradually increasing due to human activities as well as by natural causes.

As a result greater amount of surface radiated IR rays (of longer wavelengths) getting trapped in the atmosphere which in turn increases the mean [temperature of the earth. Thus greenhouse effect influences rather accelerates the phenomenon of global warming.

Question 17. Name the sources of CO₂ in atmosphere. Discuss its role in creating greenhouse effect.
Answers:
Sources of CO₂ in the atmosphere:
1. Large quantities of CO₂ is added to the atmosphere due to uncontrolled use of fossil fuels in automobiles and factories.
2. In different industries, mainly in case of cement production, a large amount of CO₂ is produced.
3. Cutting down of trees in indiscriminate manner results in the rise of the amount of CO₂ in atmosphere.

Role of CO₂ in creating greenhouse effect:
Amount of CO₂ in atmosphere is maximum, so it has the highest contribution in creating greenhouse effect. Its contribution in greenhouse effect is more than 50%.

Wbbse Class 10 Physical Science Solutions

Question 18. Name the sources of methane gas. Discuss the role of methane gas in greenhouse effect.
Answer:
Sources of methane gas:
1. Methane gas is produced in wetlands due to anaerobic decomposition of dead plants and animals by the action of methanogenic bacteria (methanococcus, methanobacterium).
2. Methane gas is found in oil mines.
3. Decomposition of several organic wastes and the excreta of different animals release methane to the atmosphere

Role In greenhouse effect: The capability of a molecule of methane to trap the re-radiated heat energy is 25 times greater than that of a molecule of CO₂. But the amount of methane in the atmosphere is much lower than that of carbon dioxide. Hence the contribution of methane in greenhouse effect is lower than that of CO₂ which is nearly around 19-20%.

Question 19. What are the sources of chlorofluorocarbons? Discuss their role in causing greenhouse effect?
Answer:
Sources of chlorofluorocarbons: Chlorofluorocarbons are extensively used in refrigerator as refrigerant, fire extinguishers,
1. Aerosol sprays as propellant,
2. Foam and plastic productions etc. Chlorofluorocarbons from these sources are added to the atmosphere.

Role of CFCs in creating greenhouse effect: The capacity of a CFC molecule to trap the retransmitted heat energy is 15000-20000 times greater than that of a CO₂ molecule. These compounds are highly stable and remain in the atmosphere for a long period of time. Contribution of CFCs towards greenhouse effect is nearly around 16%.

Question 20. Discuss the consequences of greenhouse effect.
Answer:

Detrimental consequences of greenhouse effects are

1. Average temperature of earth’s surface and troposphere is constantly increasing which may result in melting of polar ice and snow caps of mountains. Water level of the seas and oceans would thus increase and flood the coastal regions. Saline water will make agricultural land infertile and crop production will significantly decrease.
2. Frequent drought will take place in the countries of northern hemisphere.
3. Frequency and intensity of super cyclones and tornadoes will increase.
4. Living organisms will gradually become extinct due to the intense heat. This will destroy the ecosystem.

Question 21. State some precautionary measures that should be taken to stop global warming.
Answer:
1. Use of fossil fuels has to be controlled.
2. Deforestation should be stopped immediately.
3. Use of non-conventional energy should chlorofluorocarbon carbons have to be banned.
4. Nitrogenous fertilizers should be used in a controlled manner. This will reduce the possibility of mixing nitrogenous oxides in air.
5. Awareness of common people regarding greenhouse effect and global warming will surely reduce the extent of global warming.

Wbbse Class 10 Physical Science Solutions

Question 22. Mention the name of an organic and an inorganic greenhouse gas and their percentage contributions towards global warming.
Answer: An organic greenhouse gas is methane (CH₄) and an inorganic greenhouse gas is carbon dioxide CO₂. The contributions of methane and carbon dioxide towards global warming are 19% and 50% respectively.

Question 23. Why are O₂ and N₂ not considered as greenhouse gases?
Answer: Gases that can absorb infrared radiations of longer wavelengths emitted from the earth’s surface are called greenhouse gases. O₂ and N₂ gases are unable to absorb these retransmitted IR radiations of longer wavelengths, so they do not cause the greenhouse effect. That is why they are not considered as greenhouse gases.

Question 24. What is global warming?
Answer:

Global warming:

Global warming is defined as the gradual increase in the average temperature of the earth’s atmosphere mainly due to the greenhouse effect caused by the uncontrolled increase of several greenhouse gases ( CH₄, CO₂, CFC, etc.) in the atmosphere.

Question 25. Why will the crop production be disrupted as a result of global warming? Or, Mention a detrimental effect of global warming.
Answer: Average temperature of the earth’s atmosphere is gradually increasing which may result in melting of polar ice and the snow caps of mountains. Water level of the seas and oceans would thus increase and flood the coastal regions. The saline water of seas will thus make the fertile agricultural lands unsuitable for cultivation. As a result, crop production will be severely disrupted.

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming Choose The Correct Answer:

Question 1. The coldest region of the atmosphere is
1. Tropopause
2. Mesopause
3. Stratopause
4. Thermopause

Answer: 3. Stratopause

Question 2. Heat reaches the earth surface from the Sun by
1. Conduction
2. Convection
3. Radiation
4. All of these

Answer: 3. Radiation

Question 3. The depletion of ozone layer over the Antarctic region is maximum in the month of
1. March-August
2. September-November
3. January-March
4. April-June

Answer: 3. January-March

Question 4. Oxides of which element is responsible for the decomposition of ozone molecules?
1. Carbon
2. Nitrogen
3. Hydrogen
4. Sulfur

Answer: 2. Nitrogen

Question 5. CFC dissociates in the presence of ultraviolet rays to produce
1. Active carbon atoms
2. Active chlorine atoms
3. Active fluorine atoms
4. Active hydrogen atoms

Answer: 2. Active chlorine atoms

Wbbse Class 10 Physical Science Solutions

Question 6. Which oxide of nitrogen emitted from the supersonic jet planes causes depletion of ozone layer?
1. NO₂
2. N₂O
3. NO
4. N₂O₄

Answer: 3. NO

Question 7. The concentration of ozone in ozonosphere Stain is around
1. 30 ppm
2. 10 ppm
3. 40 ppm
4. 20 ppm

Answer: 2. 10 ppm

Question 8. Which of the following is a greenhouse gas?
1. O₂
2. N₂
3. O₃
4. H₂

Answer: 3. O₃

Question 9. Which of the following is not a greenhouse Jotolygas?
1. CO₂
2. NO₂
3. CH₄
4. N₂O

Answer: 2. NO₂

Question 10. The contribution of the CO₂ Greenhouse effect is almost
1. 30%
2. 20%
3. 50%
4. 60%

Answer: 3. 50%

Wbbse Class 10 Physical Science Solutions

Question 11. Naturally, ozone gas is produced due to the reaction between
1. CFC and O₂
2. UV rays and O₂
3. IR rays with O2ollatay
4. O₂ and water vapor

Answer: 2. UV rays and O₂

Question 12. The correct order of contribution of the given gases towards greenhouse effect is
1. CO₂>CH₄>CFC>O₃
2. CFC>O₃>CH₄>CO₂
3. O₃>CH₄> CFO>CO₂
4. CH₄>CO₂>O₃>CFC

Answer: 1. CO₂>CH₄>CFC>O₃

Question 13. Which would inhibit photosynthesis if it reaches the earth’s surface?
1. Infrared radiation
2. Ultraviolet radiation
3. Radio radiation
4. None of these

Answer: 2. Ultraviolet radiation

Question 14. Source of CFC is
1. A refrigerator
2. Vehicles
3. Agricultural fields
4. Water-bodies

Answer: 1. A refrigerator

Question 15. Which of the following greenhouse gas has the maximum contribution in increasing the temperature of the earth?
1. N₂O
2. CO₂
3. CH₄
4. H₂O vapour

Answer: 3. CH₄

Wbbse Class 10 Physical Science Solutions

Question 16. Dissociation of ozone molecules is done by
1. CI radical
2. OH radical
3. I radical
4. CH₃ radical

Answer: 1. CI radical

Question 17. Which of the following is required to form ozone layer?
1. Ultraviolet ray
2. Visible light
3. Microwave
4. Radiated heatwave

Answer: 1. Ultraviolet ray

Question 18. The ozone layer of the atmosphere absorbs y-Raya-ray UV-ray B-ray 19 The main greenhouse gas is
1. γ-ray
2. α-ray
3. UV-ray
4. β-ray

Answer: 3. UV-ray

Question 19. The main greenhouse gas is
1. CH₄
2. CO₂
3. CFC
4. NO₂

Answer: 2. CO₂

Question 20. Which one of the following plays the key role in depletion of ozone layer?
1. NO₂
2. O₂
3. NO
4. CFC

Answer: 4. CFC

Question 21. Mars gas formed in water-bodies is
1. CO₂
2. CO
3. CH₄
4. C₂H₂

Answer: 3. UV-ray

Wbbse Class 10 Physical Science Solutions

Question 22. Which of the following is responsible for greenhouse effect?
1. γ-ray
2. UV-ray
3. IR-ray
4. X-ray

Answer: 3. IR-ray

Question 23. The full name of CFC is
1. Chlorofluorocarbon
2. Cold fluorinated carbon
3. Carbonfluro chloride
4. Chlorofluoro carbonate

Answer: 1. Chlorofluorocarbon

Question 24. Which of the following does not deplete the ozone layer?
1. NO
2. N₂O
3. CO₂
4. DCFC

Answer: 3. CO₂

Question 25. The ozone layer is formed in the atmosphere due to
1. Chemical reactions
2. Photochemical reactions
3. Nuclear reactions
4. Electrochemical reactions

Answer: 2. Photochemical reactions

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming Answer In Brief

Question 1. Give an example of a gas which is responsible for the gradual increase of atmospheric temperature.
Answer: Carbon dioxide.

Question 2. What percentage of ozone gas found in the atmosphere is present in ozonosphere?
Answer: Ozonosphere contains almost 90% of total ozone gas present in the atmosphere.

Question 3. What happens when ozone molecules absorb ultraviolet rays?
Answer: They decompose to form oxygen molecules.

Question 4. How are UV rays harmful to the eyes?
Answer: Ultraviolet (UV) rays damage the retina of the eyes and may cause premature cataracts.

Question 5. Which instrument is used to measure the density of the ozone gas in atmosphere?
Answer: Dobson spectrometer.

Question 6. How many ozone molecules can be destroyed by an active chlorine atom?
Answer: Over 100,000 ozone molecules.

Question 7. What would be the average temperature of the earth if the greenhouse gases were not present in atmosphere?
Answer: In absence of the greenhouse gases, the average temperature of the earth would have been -30°C.

Question 8. What is the average temperature of the atmosphere due to the presence of greenhouse gases?
Answer: Due to the presence of greenhouse gases, the average temperature of the atmosphere is maintained around 15°C.

Question 9. Which diseases may flourish due to global warming?
Answer: Diseases like dengue, malaria, encephalitis may flourish due to global warming.

Question 10. What do you mean by global warming?
Answer:

Global warming

The phenomenon of gradual increase of earth’s temperature due to greenhouse effect is known as global warming.

Wbbse Class 10 Physical Science Solutions

Question 11. What is the ‘natural sunscreen’ or ‘natural solar screen’ or the ‘umbrella of the earth’?
Answer: The ozonosphere of the stratosphere.

Question 12. Incidence of which ray from the Sun is prevented by the ozone layer?
Answer: UV-ray.

Question 13. Which layer absorbs the Sun’s ultraviolet rays?
Answer: Ozone layer present in the stratosphere.

Question 14. Write down the full form of ODS.
Answer: Ozone Depleting Substances.

Question 15. What is ‘Chapman Cycle’?
Answer:

‘Chapman Cycle’

The cycle of creation and destruction of ozone in the stratosphere is called the Chapman Cycle.

Question 16. What is the full form of ODP?
Answer:

Full form of ODP

Ozone Depletion Potential.

Question 17. Which atom of CFC is responsible for the destruction of the ozone layer?
Answer: The cl atom of CFC is responsible for this.

Question 18. Which gas breaks down the ozone layer during lightning?
Answer: Nitric oxide (NO) gas.

Question 19. What is the role of NO in the depletion of ozone layer?
Answer: NO reacts with O₃ to form NO₂ and O₂ and thus decomposes O_3. As a consequence depletion of ozone layer occurs.
NO + O₃→ NO₂+O₂+hγ

Question 20. In which year was the Montreal Protocol signed?
Answer: Montreal Protocol was signed in 1987.

Question 21. Write the name of an organic greenhouse gas.
Answer: Methane (CH₄).

Question 22. Which is the main greenhouse gas?
Answer: Carbon dioxide (CO₂).

Question 23. Write the full form of GWP.
Answer: Global Warming Potential.

Question 24. Mention a detrimental effect of global warming.
Answer: As global warming increases, polar ice caps will melt and as a result, coastal areas will be flooded.

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming Fill In The Blanks

Question 1. Ozonosphere is present in the  _______.
Answer: Stratosphere

Question 2. ________ was the first to observe the formation of holes in the ozone layer.
Answer: Joe Farman

Question 3. The full form of CFC is a ___________
Answer: Chloroflurocarbon

Question 4. Ozone is a gas with pale blue, _________ Color and a smell.
Answer: Pungent

Question 5. Ozone gas is an allotrope of _______ gas.
Answer: Oxygen

Question 6. The density of ozone gas in atmosphere is measured in the_______ unit.
Answer: Dobson

Question 7. Nitrous oxide and water vapor are _______ gases.
Answer: Greenhouse

Question 8. _______ layer of the atmosphere is also known as the ‘earth’s protective umbrella’.
Answer: Ozone

Question 9. The capacity of a molecule of methane to retain the absorbed heat energy is ______ times more than that of a CO₂ molecule.
Answer: 25

Question 10. The wavelength of infrared radiation______ with decreasing energy.
Answer: Increases

Question 11. The rate of increase of CFC gas in the atmosphere per year is ______
Answer: 5%

Question 12. The gas emitted by supersonic aircraft breaks down the ozone layer is _______
Answer: Nitric Oxide

Question 13. The destruction of the ozone layer can cause diseases like ________
Answer: Skin cancer

Question 14. An inorganic greenhouse gas is _________
Answer: Carbon dioxide

Question 15. The _____________ present in the body spray is a greenhouse gas.
Answer: Chloroflurocarbon

Question 16. The greenhouse effect is causing ____________
Answer: Global Warming

Chapter 1 Environmental Concern Topic B Ozone Layer, Greenhouse Effect and Global Warming State Whether True Or False:

Question 1. Nitrogen is a greenhouse gas.
Answer: True

Question 2. Supersonic jets release huge amount of nitric oxide.
Answer: False

Question 3. Scientist Joe Farman first noticed the formation of the ozone hole over Antarctica.
Answer: False

Question 4. Under the influence of UV rays, CFC splits and produces methane gas.
Answer: True

Question 5. The Montreal Protocol was signed in the year 1987.
Answer: False

Question 6. Damage to the ozone layer will increase the chance of cataracts.
Answer: False

Question 7. It would have been easier for the fauna to survive if there were no greenhouse gases in the atmosphere.
Answer: True

Question 8. The two gases responsible for the greenhouse effect are nitrogen and oxygen.
Answer: True

Question 9. Among all greenhouse gases the contribution of carbon dioxide is maximum in creating greenhouse effect.
Answer: False

Question 10. If there were no greenhouse gases, the temperature of the earth would be 50°C.
Answer: True

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development Synopsis

Calorific value: Heat energy produced due to the complete combustion of 1 kg of a fuel is known as its calorific value. Unit: kJ • kg-1 (in SI).

The necessity of conservation of fossil fuels:
The incessant use of fossil fuels (coal, petroleum, natural gas) is gradually depleting the reserve of these fuels. To conserve these fuels, we must encourage the use of alternative fuels which are basically non- conventional renewable energy sources.

Alternative fuels: Alternative fuels are those which are being used to restrict the use of fossil fuels in order to conserve their sources as these are gradually running out. Examples: Solar energy, wind energy, tidal energy etc.

Sustainable development: Controlled and systematic use of natural resources to meet the need of present generation as well as sustaining them for future generation is called sustainable development. First step of sustainable development is to encourage the use of alternative sources of energy that are renewable and eco-friendly.

Purposes of sustainable development: Social welfare, Economic development and Ecological development are the major three purposes of sustainable development.

Conventional/non-renewable sources of energy: These sources are gradually exhausting due to their prolonged uses and these sources cannot be renewed.
Examples: Coal, petroleum, natural gases etc.

Non-conventional/renewable sources of energy: These sources have presently been used as alternative energy sources and these are inexhaustible and renewable.

Examples: Solar energy, wind energy, biomass energy, tidal energy etc.

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development Short And Long Answer Type Questions:

Question 1. What is sustainable development?
Answer:

Sustainable development:

The controlled and systematic use of natural resources to meet the need of present generation as well as sustaining resources for the future generation is called sustainable development.

Question 2. Discuss the purposes of development’.
Answer:
Purposes of sustainable development are Social welfare:
Major goal of sustainable development is to establish social equality and eradicate the differences that arise from social misconceptions, superstitions, backward traditions etc.

Economic development: No society can prosper without economic development. It is required in the field of education, health, agriculture and industries. Hence, one of the major goals of sustainable development is to achieve economic development.

Ecological development: Ecology of earth depends on the constant interaction between living and non-living things. Sustainable development also aims at maintaining proper balance in nature and building a stronger relationship between living & non-living world. of ‘sustainable

Question 3. Why is sustainable development considered to be of utmost importance for the present as well as a future generations?
Answer: The aim of sustainable development is to use natural resources without affecting the environment. This will enable both present and future generations to preserve natural resources and utilize them for a long time.

It gives importance not only to economic development but also to cultural, social & geopolitical development have been able to control the decline of environmental conditions as it encourages to decrease the use of conventional sources of energy and increases the use of alternative resources.

Question 4. What are fossil fuels?
Answer:

Fossil fuels:

Fossil fuels are natural fuels formed due to the action of geothermal heat, pressure, and bacterial degradation of the dead plants and animals buried beneath the earth’s crust for many years. Coal, petroleum, natural gas are examples of fossit fuels. These are exhaustible and non¬ renewable sources of energy.

Question 5. What do you mean by conventional or non-renewable sources of energy? Give examples.
Answer:

The sources of energy which- are gradually exhausting due to their prolonged uses and which cannot be renewed are called conventional or non-renewable sources of energy.

These have been used as the major energy source till date and hence their reserves are gradually running out. Coal, petroleum, natural gases are examples of conventional or non-renewable sources of energy.

Question 6. What are non-conventional? orrenewable sources of energy? Give examples.
Answer:

The sources of energy that have presently been used as alternative energy sources to minimize the uses of fossil fuels and which are in exhaustible and renewable are known as non-conventional or renewable sources of energy. Solar energy, wind energy, biomass energy, tidal energy, etc. are examples of non-conventional or renewable sources of energy.

Question 7. Write three differences between conventional and non-conventional sources of energy.
Answer:

The major differences are listed below:

Conventional sources of energy	Non-conventional sources of energy
1. These sources of energy are extensively used and their reserves are gradually exhausting. Examples—are coal,    petroleum, natural gas, etc.	1. These sources of energy are not being used extensively. These sources are exhaustible and they are mostly reusable in nature. Example-wind energy, are solar energy, tidal energy, etc.
2. These sources of energy cause pollution.	2. These sources are non-polluting in nature.
3. The sources of energy are not evenly distributed throughout the world.	3. Non-conventional energy sources are almost evenly distributed.

Question 8. Why should fossil fuels?
Answer:

The natural reserves of fossil fuels are gradually running out. The sources of such fuels are abruptly exhausting due to their indiscriminate usage. As a result, the world may face a severe energy crisis and development of human civilization may cease in near future.

These sources of energy are non-renewable. Combustion of fossil fuels produces several greenhouse gases. Hence controlled use of these fuels will minimize the chances of global warming.

Question 9. How is electricity generated from wind energy?
Answer:

Practically, the kinetic energy of the wind is harnessed to obtain wind energy. Using windmills, wind energy can be converted into electrical energy.

The blades of the windmills are connected to the turbines of an electrical generator. When the blades rotate by the action of wind, the turbines also rotate and electricity is generated.

Question 10. How is electricity generated using tidal energy?
Answer:

During high tides, the surging water is stored in a dam and during the low tides this stored water is released, which comes down with a high velocity.

Turbines of electrical generators are placed in the path of moving water and thereby the kinetic energy of water can be converted into electrical energy.

Question 11. How is electricity generated using geothermal energy?
Answer:

During volcanic eruptions, most of the magma coming out from the earth, gets stuck in the rocks near the earth surface at a depth of 5 to 10 km. These rocks, in contact with the magma, gets heated up.

When water passes through these hot rocks, it vaporizes to form steam. The steam is piped out directly from the underground wells to the power plants. The kinetic energy of the steam rotates the turbines and generates electricity.

Question 12. What are methanogenic bacteria? Give examples.
Answer:

Methanogenic bacteria

Methanogenic bacteria can be defined as microorganisms that produce methane gas by the anaerobic decomposition of biomass. Some examples of methanogenic bacteria are Methanococcus, Methanobacterium etc.

Question 13. Discuss the role of methanogenic bacteria in increasing the amount of methane in the atmosphere.
Answer:

Methanogenic bacteria decompose dead plants in wetlands, rain forests etc. to produce methane gas. Thus, the amount of methane in atmosphere increases.

Question 14. What is biomass energy? What are the disadvantages of using biomass energy?
Answer:

Biomass energy:

Biomass refers to dead plants and plant remains, domestic wastes, agricultural wastes, animal excreta, animal carcasses etc. Energy stored in biomass is known as biomass energy.

Biomass energy is considered to be one of the major alternative sources of energy to replace fossil fuels. However, it has some harmful effects on environment as the combustion of biomass causes air pollution.

Question 15. What is biogas? What are its constituents?
Answer:

Biogas:

Biomass refers to dead plants and plant remains, domestic wastes, agricultural wastes, animal excreta, animal carcasses, etc. The biomass is enclosed in a large chamber and decomposed by the action of methanogenic bacteria to produce a combustible gas, known as biogas.

The major components of biogas is methane (CH₄). Apart from methane, carbon dioxide CO_2, and traces of H₂, N₂, CO, O_2, H₂S, and water vapor are also present in it.

Question 16. Write some uses of biogas.
Answer:

Some uses of biogas:

Biogas can be used as a fuel to light lamps, to heat water, and for cooking purposes. In many countries, it is used to dry bricks, ceramic tiles, tobacco etc. Electricity generated from biogas can be used to run small pumps and small-scale industries.

Question 17. What is biofuel? Give examples.
Answer:

Biofuel:

The fuel produced from biomass is known as biofuel. Bioethanol produced by fermentation of corn or sugar cane is an example of biofuel. It is mixed with petrol and used as automobile fuel.

Question 18. What is methane hydrate?
Answer:

Methane hydrate:

Methane hydrate is a crystalline solid substance (clathrate compound) in which a large amount of methane molecules are trapped inside the cage-like lattice of ice. It is represented by the formula 4CH₄ 23H₂O. It exists only under suitable temperatures and pressure.

Question 19. How is methane present in coal mines?
Answer: Due to bacterial action or geothermal heating, methane gas is produced in coal seams. The coal seams are saturated by ground water and the produced water pressure causes methane gas to get absored in the coal seam.

Question 20. Why is methane gas extracted before the mining of coal?
Answer: Before the mining of coal, methane gas should be extracted as much as possible from the coal bed. This is because- it will reduce the probability of catching fire and chances of accidents can be averted, the extracted methane gas can be used as a potential fuel.

Question 21. Why is methane hydrate considered as an important source of energy?
Answer: Recent discovery of methane hydrate explores a new dimension of energy resources. Scientists have estimated that methane hydrate contains more carbon than all the fossil fuel available on the earth, combined together.

1 Litre of methane hydrate is found to contain 170 Litre of methane gas in STP. When the world is facing an energy crisis due to exhaustion of fossil fuels, methane hydrate provides a complete new direction towards energy resources.

Question 22. State the concerns of using methane?
Answer:

Concerns of using methane

Methane hydrate is not as stable as other minerals. Due to increase in temperature or decrease in pressure, it may be transformed into water and methane gas may be released from the mineral. This may cause landslide beneath the earth surface. Methane is a major greenhouse gas. Heating methane hydrate may release methane gas to the atmosphere which may lead to global warming.

Question 23. Discuss some advantages of using solar energy.
Answer:

Some advantages of using solar energy:

Solar energy is inexhaustible and renewable. Scientists consider it as the most important source of alternative energy. Unlike fossil fuels, solar energy does not cause pollution. source of energy is very effective in less populated areas and in those regions where the .conventional sources of energy are not widely available. The production of solar energy requires minimum natural resources.

Question 24. Mention some disadvantages of using solar energy.
Answer:

Some disadvantages of using solar energy:

Production cost of solar energy is very high as it requires complex technological support. So, under-developed countries find this energy unsuitable for use.

The initial cost of installing solar panels is very high. This is one of the major disadvantages of solar energy. Areas which do not get sufficient sunlight, are not suitable for utilising solar energy. Solar energy centers produce solar energy in limited quantities which can satisfy only the local need.

Question 25. Mention some advantages of wind energy.
Answer:

Some advantages of wind energy:

Wind energy is renewable and is source is inexhaustible. Even it’s incessant use will not create energy crisis. Even though initial cost of installation is somewhat high, regular cost is very create energy crisis. Recurring cost is very low. This energy source does not cause pollution.

Question 26. Mention some advantages of tidal energy.
Answer:

Some advantages of tidal energy:

Tidal energy is renewable and is source is inexhaustible. Even it’s incessant use will not create energy crisis. Recurring cost is very slow. This energy sorce does not cause pollution.

Question 27. Discuss some advantages of geothermal energy.
Answer:

Some advantages of geothermal energy:

1. This source of energy is inexhaustible.
2. Being inexhaustible, it can be supplied continuously throughout the day.
3. Direct use of this energy inhibits the release of greenhouse gases into the atmosphere, thereby preventing pollution.

Question 28. Mention some advantages of biomass energy.
Answer:

Some advantages of biomass energy:

Biomass energy is a renewable source of energy. The residue left after production of biogas from biomass can be used as fertilizers.

Biogas can be used as fuels and for the generation of electricity. Pollution caused by the waste material during biomass production can be controlled by using them for production of biogas.

Question 29. How is electricity produced from domestic wastes?
Answer: Initially, substances that can be recycled are separated from the domestic wastes. Remaining wastes are then placed into a furnace through a hopper. These are then burnt in the combustion chamber. The heat so produced vapourises the water taken in a tank above the furnace.
The vapour is used to rotate the turbine which in turn generates electricity.

Question 30. Mention some uses of coalbed methane (CBM).
Answer:

Some uses of coalbed methane (CBM)

Coalbed methane (CBM) can be used as a fuel for domestic as well as industrial purposes. It can be used as a substitute of natural gas in power stations. It can be liquefied by applying pressure. The liquefied gas can be used in automobiles.

Question 31. Classify biofuel. Give example of each Characteristics of each type?
Answer:

Biofuel:

The fuel produced from biomass is called biofuel. It is classified into three groups- solid biofuel, liquid biofuel and gaseous biofuel. Solid biofuel: Wood, straw, domestic wastes etc., are examples of solid biofuels.

Liquid biofuel: Bioethanol (produced by fermentation of corn or sugarcane), biodiesel (produced by transesterification of oil or fat) are examples of liquid biofuels.

Gaseous biofuel: Biogas or gobar gas is an example of gaseous biofuel.

Question 32. How is methane gas extracted from coal mines?
Answer: A vertical hole is drilled through the ground into coal seam. Coal seam remains saturated in When the water is pumped out, the pressure over the coal seam decreases and methane gas is released from the coal seam. The gas then comes out through the hole and its collected and supplied through pipelines.

Question 33. What is ignition point of a fuel? What are the characteristics of a good fuel?
Answer:

The ignition point of a fuel can be defined as the lowest temperature at which the fuel catches fire and undergoes combustion.

Characterisitics of a good fuel:

1. It should possess a high calorific value.
2. It should have a proper ignition temperature. The ignition temperature of the fuel should neither be too low or too high.
3. It should not produce poisonous products during combustion. In other words, it should not cause pollution during the time of combustion.
4. It should have a moderate rate of combustion. Combustion should easily be controlled i.e.
5. It should be easier to start or stop the combustion of the fuel when required. Combustion of the fuel should not produce. much ash.
6. It should be cheap and well available. The transportation and handling of the fuel must be easier.

Question 34. What do you mean by calorific value of fuel? Which one of coal and diesel has higher calorific value?
Answer:

Calorific value of fuel

The amount of heat produced due to the complete combustion of 1 kg of a fuel is known as its calorific value. Its value is represented in It is actually a metal box with a heat-insulating kl. kg^-1 unit in Sl.

The calorific value, of diesel (44800 kJ. kg^-1) is greater than that of coal (15000-34000 kJ.kg^-1).

Question 35. What are the uses of solar energy?
Answer:

Solar energy is a renewable, inexhaustible, and affordable form of energy.

It can be used in following ways-

1. Solar electricity can be generated through solar panels and is used to light streets, houses, health centres etc.
2. Solar water heaters are used to heat water by Solar water heater is used to heat water using using solar energy.
3. Solar space heating systems include radiant floors or a Force Hot Air (FHA) system to heat a home thus reduces the overall electricity cost.
4. Solar ventilation solution such as solar attic fans can reduce the burden of HVAC by helping to cool the home during summer.
5. Solar cookers can be used for cooking purposes.
6. Solar powered vehicles (e.g. buses, small cars etc.) are used in many cities.

Question 36. Mention the uses of solar cells.
Answer:

In a solar cell, solar energy is converted into electrical energy and the generated electricity can be stored in batteries or can be used directly.

1. Water can be pumped up with the help of this energy.
2. Motor cars can be driven by installing solar panels.
3. This energy can be used in artificial satellites, space ships or space stations.
4. Radio stations and TV relay stations can use this energy.
5. This energy can be used to supply powers to traffic signals, solar calculators etc.

Question 37. Briefly discuss about solar cooker.
Answer:

The properties of solar cooker are:

1. Solar cooker is a cheap, energy-saving, pollution-free tool used in cooking. the environment i.e. it is completely pollution free.
2. It is actually a mental box with a heat-insulting coating around it.
3. Black color is avery good heat absorber, so the inner surfaces are painted black.
4. With the help of a reflector, the sunlight is concentrated and applied to the box to create high temperature.
5. The surface of the cooker is covered with a glass plate to increase the temperature inside

Question 38. Briefly discuss about solar water heater.
Answer:

The properties of water heater are:

1. Solar water heater is used to heat water using solar energy.
2. A black-painted metal box and a black- painted copper coil are placed inside it.
3. Sun rays are being reflected into the box with the help of a reflector to create a high temperature, which also heats the inner copper coil.
4. The water flowing through the hot copper coil is also heated in this way.

Question 39. Mention a few uses of wind energy.
Answer:

Few uses of wind energy are:

1. Wind power is used to generate electricity by winding the wind turbines.
2. With the help of windmill it is possible to grind different grains or spices, pump water etc.

Question 40. Mention the advantages and disadvantage of using geothermal energy.
Answer:

Advantages of using geothermal energy:

1. The supply of geothermal energy is inexhaustible, i.e., its stock will never run out despite being used repeatedly.
2. Using of geothermal energy causes no harm to the environment i.e. it is completely pollution free.

Disadvantages of using geothermal energy:

1. The construction of geothermal power generation plants is extremely expensive.
2. Due to the use of advanced technology, the cost per unit of this power generation is higher than that of conventional power plant.

Question 41. Write the differences between conventional and non-conventional sources of energy.
Answer:

Differences between conventional and non-conventional sources of energy:

Conventional sources  of energy	Non-conventional Sources of energy
1. Conventional sources of energy (e.g. coal, petroleum, and natural gas) are nonrenewable sources of energy and that is why there is a probability of that such sources might be exhausted by excessive use.	1. Non-conventional sources of energy (e.g. solar energy, wind energy) are renewable sources of energy and thus they can not be exhausted.
Z- These energy sources have been used for a long time. For. example, the use of coal and petrol has been around for a . long time.	2. The use of these energy sources (e.g., solar energy, wind energy, tidal energy) is not very common. Their use has recently become popular.
3. Use of these energy sources causes environmental pollution.	3. Use of these energy sources does not cause environmental pollution at all.
4. The use of these energy sources is not conducive to sustainable development.	4. The use of these energy sources is conducive to sustainable development.
5. A variety of byproducts can be obtained during the extraction of these energies. For example, naphthalene, coal tar, etc. are obtained during the extraction of coal.	5. No by-products can be obtained during the extraction of these energies.

Question 42. In which type of natural environment is methane hydrate found?
Answer: The topography and environment of some parts of the world are favorable to the formation and sustainability of methane hydrate. Those are:

1. Sedimentary rock layers below the frozen soil layer (Permafrost) of the polar region.
2. The sedimentary layer is present at the bottom of large lakes and the seas in the subcontinental region.
3. The bottom region of frozen ice in the polar region.

Question 43. Why is methane hydrate termed as ‘fiery ice’?
Answer:

Methane hydrate is a solid, crystalline substance having the composition of 4CH₄. 23H₂O. A large amount of methane is trapped in these ice-like crystals. Methane is a combustible substance, hence methane hydrate starts to burn on contact with fire. Only water is left after combustion. That is why methane hydrate is termed as ‘fiery ice’.

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development Choose The Correct Answer:

Question 1. Which of the following is a fossil fuel?
1. Coal
2. Petroleum
3. Natural gas
4. All of these

Answer: 3. Natural gas

Question 2. Among petrol, diesel, kerosene, and LPG, which one has the highest calorific value?
1. Peterol
2. Kerosene
3. Diesel
4. LPG

Answer: 4. LPG

Question 3. Which of the following represents the correct order of calorific values of the corresponding fuels?
1. Coal < diesel < hydrogen < LPG
2. Coal < hydrogen < diesel < LPG
3. Coal < diesel < LPG < hydrogen
4. Coal < LPG < diesel < hydrogen

Answer: 3. Coal < diesel < LPG < hydrogen

Question 4. Which of the following is a fossil fuel?
1. Coal
2. Petroleum
3. Natura! gas
4. All of these

Answer: 4. All of these

Question 5. Fossil fuels need to be conserved because
1. These are very expensive
2. These are not easily available
3. These are non-renewable
4. None of these

Answer: 3. These are non-renewable

Question 6. Which of the following elements is the major constituent of solar cells?
1. Li
2. Na
3. Si
4. Cu

Answer: 3. Si

Question 7. The major source of energy which meets our daily demands is
1. Solar energy
2. Wind energy
3. Tidal energy
4. Fossil fuels

Answer: 4. Fossil fuels

Question 8. Solar energy gets converted into electrical energy in a/an
1. Photo voltaic cell
2. Solar cooker
3. Pressure cooker
4. Electric motor

Answer: 1. Photo voltaic cell

Question 9. The least polluting fossil fuel is
1. Diesel
2. Coal
3. Kerosene
4. Natural gas

Answer: 4. Natural gas

Question 10. Source of non-renewable energy is
1. Wind energy
2. Petroleum
3. Tidal energy
4. Solar energy

Answer: 2. Petroleum

Question 11. The energy stored in coal, petroleum, etc. is
1. Wind energy
2. Solar energy
3. Tidal energy
4. Renewable energy

Answer: 2. Solar energy

Question 12. Which gas found in coal mines is used as a fuel?
1. Ozone
2. Methane
3. Oxygen
4. Nitrogen

Answer: 2. Methane

Question 13. An example of methanogenic bacteria is
1. Methanococcus
2. Methanobacterium
3. Both 1 and 2
4. None of these

Answer: 3. Both 1 and 2

Question 14. Which of the following does work on the principle of the greenhouse effect?
1. Solar cooker
2. Radio station
3. Windmill
4. Biogas plant

Answer: 1. Solar cooker

Question 15. The largest windmill of India is situated in
1. Kanyakumari (approx. 380MW)
2. Surat (approx. 220MW)
3. Haridwar (approx. 460MW)
4. Sagardwip (approx. 410MW)

Answer: 1. Kanyakumari (approx. 380MW)

Question 16. The state with the largest installed wind power generation capacity in India is
1. Tamil Nadu
2. West Bengal
3. Kerala
4. Maharashtra

Answer: 1. Tamil Nadu

Question 17. Which of the following are eco-friendly fuels
(1)Coal,
(2)Kerosene,
(3)Natural gas,
(4)Biogas

1. (1) and (4)
2. (3) and (4)
3. (1) and (2)
4. (1) and (3)

Answer: 3. (1) and (2)

Question 18. Which type of coal has the highest calorific value?
1. Pit
2. Lignite
3. Anthracite
4. Bituminous

Answer: 3. Anthracite

Question 19. The main component of biofuel is
1. Carbon dioxide
2. Ethanol
3. Methanol
4. Ether

Answer: 2. Ethanol

Question 20. Solar cells are formed by
1. Conductors
2. Insulators
3. Semiconductors
4. Superconductors

Answer: 3. Semiconductors

Question 21. Source of the solar energy is
1. Nuclear fission
2. Coal
3. Nuclear fusion
4. Petroleum

Answer: 3. Nuclear fusion

Question 22. Formula of methane hydrate is
1. CH_3.H20
2. 4CH₃-23H₂O
3. CH₃OH
4. 4CH₄-23H₂O

Answer: 4. 4CH₄-23H₂O

Question 23. Fiery ice is
1. CO₂
2. Methane
3. Methane hydrate
4. Methanol

Answer: 3. Methane hydrate

Question 24. For which ray of sunlight does a solar cooker work?
1. Gamma ray
2. Infrared ray
3. Ultraviolet ray
4. Visible ray

Answer: 3. Methane hydrate

Question 25. What should be the minimum wind velocity for proper operation of a wind-power generator?
1. 15 km/h
2. 25 km/h
3. 35 km/h
4. 45 km/h

Answer: 1. 15 km/h

Question 26. The main component of biogas is
1. CO₂
2. CH₄
3. H₂
4. H₂S

Answer: 2. CH₄

Question 27. Which is not a source of biomass energy?
1. Agricultural waste
2. Cow dung
3. Wood
4. Nuclear fission

Answer: 4. Nuclear fission

Question 28. A source of non-conventional energy is
1. Petrol
2. Kerosene
3. Biogas
4. Coal

Answer: 3. Biogas

Question 29. In which of the following states does methane hydrate exists?
1. Solid
2. Liquid
3. Gas
4. Any physical state

Answer: 1. Solid

Question 30. The potential difference formed in solar cells
1. 0 V
2. 0.5 V
3. 220 V
4. 110 V

Answer: 2. 0.5 V

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development Answer In Brief

Question 1. What do you mean by alternative fuels?
Answer:

Alternative fuels

Alternative fuels are those which are being used to restrict the use of fossil fuels in order to conserve their sources as these are gradually running out.

Examples: solar energy, wind energy, etc.

Question 2. Mention two major sources of energy.
Answer:

Two major sources of energy

1. Conventional sources of energy and
2. Non-conventional sources of energy.

Question 3. Name a conventional and renewable source of energy.
Answer: Wind power.

Question 4. Write the name of a conventional and nonrenewable source of energy.
Answer: Coal (fossil fuel).

Question 5. Write the names of two non-conventional and renewable sources of energy.
Answer:
1. Solar energy,
2. Tidal energy.

Question 6. Write the names of two non-conventional and non-renewable sources of energy.
1. Geothermal energy,
2. Nuclear energy.

Question 7. Why is the inner surface of a solar heater painted black?
Answer: Black surface can absorb more heat from sunlight. So, the inner surface of a solar heater is painted black.

Question 8. Which fuel has the highest calorific value?
Answer: Hydrogen (150000 kJ/kg) has the highest calorific value.

Question 9. Why do the pipelines of natural gas often get blocked?
Answer: The pipelines of natural gas often get blocked due to the formation of solid methane hydrate.

Question 10. The gas trapped in coal mines is a potential source of energy. Name the gas. Why is it called firedamp? OR, What is firedamp?
Answer:
1. The gas is methane (CH₄).
2. Methane gas accumulated in coal mines forms an explosive mixture when mixed with air and often causes fire in coal mines. So this gas is named as firedamp.

Question 11. How is electricity produced in solar cells?
Answer: When photons present in sunlight are incident on the surface of the solar cell (made up of semiconductors, such as Si), the electrons get excited and ultimately becomes free. These free electrons generate electricity.

Question 12. What is geothermal energy?
Answer:

Geothermal energy

When groundwater passes over hotbeds of rocks, it gets vaporised. These vapours are utilized to move the plates of a turbine thereby producing electrical energy. This is known as geothermal energy.

Question 13. What is tidal energy?
Answer:

Tidal energy

During tides, the kinetic energy of water can be converted into electrical, energy using turbines. This energy is known as tidal power or tidal energy.

Question 14. Which is the major raw material used for the production of biogas?
Answer: Cow dung.

Question 15. What is the calorific value of fuel?
Answer:

Calorific value of fuel

The amount of heat produced due to the complete combustion of 1 kg of a fuel is known as its calorific value. Its value is represented in kJ • kg-1 unit in SI.

Question 16. How is biodiesel produced?
Answer: Biodiesel is produced due to the transesterification of vegetable oil and animal fats.

Question 17. Which seeds are used to produce biodiesel in India?
Answer: The seeds of the jatropha tree are used to produce biodiesel in India.

Question 18. What is ‘methane hydrate’?
Answer:

‘Methane hydrate’

Methane hydrate is a crystalline solid where methane molecules remain trapped inside a cage-like lattice of ice.

Question 19. Write down the unit of calorific value of a fuel.
Answer: The SI unit of calorific value of a fuel is J/kg and the CGS unit is cal/g.

Question 20. The calorific value of LPG is 50 kJ/g —Explain.
Answer: It means that 50 kJ of heat is obtained by the combustion of 1 g of LPG.

Question 21. Arrange in ascending order of calorific value—Kerosene, Wood, Coal, Hydrogen, LPG, Diesel.
Answer: Wood > Coal > Diesel > Kerosene > LPG > Hydrogen

Question 22. Write the name of fossil fuel.
Answer: Coal.

Question 23. Which gas is filled in the gas cylinder used for cooking?
Answer: Liquefied Petroleum Gas (LPG)

Question 24. What is the main component of LPG?
Answer: Butane.

Question 25. What is the main component of CNG?
Answer: Methane.

Question 26. Which of the following is a fossil fuel? Charcoal, Petrol, Ethanol.
Answer: Petrol

Question 27. Which gas is produced during the combustion of fossil fuels?
Answer: CO₂(Carbon dioxide).

Question 28. Write the name of a conventional renewable pollution-free energy source.
Answer: Wind power.

Question 29. Write the name of an environment-friendly non-conventional energy source.
Answer: Geothermal energy.

Question 30. Write the name of an alternative energy source.
Answer: Solar energy.

Question 31. Write the name of a device that uses solar cells.
Answer: Solar calculator.

Question 32. What is generated from the wind power of wind mill?
Answer: Electrical energy.

Question 33. Write the name of a place in West Bengal that has windmills.
Answer: Fraserganj.

Question 34. Name an energy source that does not come directly or indirectly from the sun.
Answer: Geothermal energy.

Question 35. What is available as residue in the controlled construction of wood?
Answer: Charcoal.

Question 36. Methanogenic Bacteria decompose substances to produce methane gas.
Answer: Biomass.

Question 37. Give an example of Methanogenic Bacteria.
Answer: Methanogens—Methanosarcina barkeri

Question 38. Mention the use of Biogas.
Answer: Biogas can be used as a fuel for cooking purposes.

Question 39. Write down the fuel form of CBM.
Answer: Coalbed methane.

Question 40. What is sweet gas?
Answer:

Sweet gas

Methane gas obtained from coal mines is called sweet gas.

Question 41. What is ‘fiery ice’?
Answer:

‘Fiery ice’

Methane Hydrate ((4CH₄• 23 H₂O ) is called ‘fiery ice’.

Question 42. How much volume of methane gas can be obtained from 1 Litre of methane hydrate at NTP?
Answer: Almost 170 Litre of methane gas can be obtained from 1 Litre of methane hydrate at NTP.

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development Fill In The Blanks:

Question 1. The calorific value of LPG is _________.
Answer:  55000 kJ/kg.

Question 2. The major component of biogas is ________
Answer: Methane

Question 3. Apart, from methane, another gas that is found in biogas is ________.
Answer: CO₂

Question 4. At STP, about __________ of methane is obtained from 1 L of methane hydrate.
Answer: 170L

Question 5. Methane hydrate is a solid _________ compound.
Answer: Clathrate

Question 6. _________ river basin in India has one of the largest reserves of methane hydrate.
Answer: Krishna-Godavari

Question 7. The idea of sustainable development is properly developed in the report of the __________ Commission.
Answer: Brundtland

Question 8. Biogas is produced due to the anaerobic decomposition of cow dung by __________ bacteria.
Answer: Methanogenic

Question 9. The calorific value of methane is _________ than petrol.
Answer: Less

Question 10. In photovoltaic cells, solar energy is directly converted into _________ energy.
Answer: Electrical

Question 11. In artificial satellites, the main source of energy is the energy produced by __________
Answer: Solar Cells

Question 12. The value of solar constant is approximately ________
Answer:  1.4kW/m²

Question 13. Biogas contains _______ of methane gas.
Answer: 50-80%

Question 14. Among wood and methane, the calorific value is more in case of __________
Answer: Methane

Question 15. The substance used to make solar cell is ________
Answer: Silicon

Question 16. Electrical energy is obtained from in solar cells ___________
Answer: Solar energy

Question 17. _________ is generated from wind power in wind mills.
Answer: Mechanical Energy

Question 18. _________ in West Bengal is a region that is rich in geothermal energy.
Answer:  Barkreshwar

Question 19. Vegetable wastes are decomposed by anaerobic bacteria in __________ of air in the biogas plants.
Answer: Absence

Question 20. To produce biogas___________ bacteria is needed.
Answer: Methanogenic

Question 21. The hard coal in the mine contains a large amount of_______ in the absorbed state.
Answer: Methane

Question 22. _____________ is also termed as fiery ice.
Answer: Methane Hydrate

Question 23.  _________ gas is obtained from methane hydrate.
Answer: Methane

Chapter 1 Environmental Concern Topic C Proper Uses of Energy and Sustainable Development State Whether True Or False:

Question 1. The calorific value of kerosene is greater than LPG.
Answer: False

Question 2. Semiconductors like Si, Ge, etc., are used in solar cells.
Answer: True

Question 3. Gasoline is a kind of biofuel.
Answer: False

Question 4. Methanogenic bacteria play a major role in biogas production.
Answer: True

Question 5. C02 gas is trapped inside the coal mines.
Answer: False

Question 6. The quality of fuel is determined by its calorific value.
Answer: True

Question 7. Methane hydrate contains 13.3% methane. Bioethanol is a liquid biofuel.
Answer: False

Question 8. The source of solar energy is the nuclear fission reaction that takes place inside the Sun.’
Answer: False

Question 9. The chemical energy stored in biomass is called biomass energy,
Answer: True

Question 10. Solar panels are used in broadcast relay stations.
Answer: True

Question 11. The most widely used definition of sustainable development was produced by the report of the Brundtland Commission.
Answer: True

Question 12. Windmills usually convert electrical energy into mechanical energy.
Answer: False

Question 13. Bioethanol is a liquid biofuel.
Answer: True

Chapter 1 Environmental Concern Miscellaneous Type Questions Match The Columns:

Question 1

Column A	Column B
Chlorofluro-carbon	1. Non-conventional source of energy
Biofuel	2. Ozone layer depletion
Solar energy	3. Fossil fuel
Petroleum	4. Ethanol

Answer:
Chlorofluro-carbon: 2. Ozone layer depletion
Biofuel: 4. Ethanol
Solar energy: 1. Non-conventional source of energy
Petroleum: 3. Fossil fuel

Question 2

Column A	Column B
Greenhouse gas	1. Tidal energy
Renewable source of energy	2. Photovoltaic cell
Natural gas	3. Global warming
Solar energy	4. Methane

Answer:
Greenhouse gas: 3. Global warming
Renewable source of energy: 1. Tidal energy
Natural gas: 4. Methane
Solar energy: 2. Photovoltaic cell

Question 3

Column A	Column B
Turbulent sphere of atmosphere	1. Mesopause
Aurora borealis	2. Exosphere
The coldest region of the atmosphere	3. Troposphere
Hydrogen and helium gases are present	4. Ionosphere

 

Answer:
Turbulent sphere of atmosphere: 3. Troposphere
Aurora borealis: 4. Ionosphere
The coldest region of the atmosphere: 1. Mesopause
Hydrogen and helium gases are present: 2. Exosphere

WBBSE Solutions for Class 10 Physical Science and Environment

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table Synopsis:

Dobereiner’s Law of Triads: In 1817, German scientist Dobereiner first attempted to arrange the elements on the basis of their atomic weights. According to his Law of Triads, if three elements having similar properties are arranged in increasing order of their atomic weights, the atomic weight of the middie element in the triad is found to be the average of the atomic weights of the remaining two elements.

Newland’s Law of Octaves: When the known elements are arranged in increasing order of their atomic weights, the eighth element, starting from a given one, will show resemblance in physical and chemical properties with the first element. This is known as Newland’s Law of Octaves (1861).

Read and Learn Also WBBSE Solutions for Class 10 Physical Science and Environment

Mendeleev’s Periodic Law: In 1869, Mendeleev put forward the periodic law through which he established the relationship between atomic weight of an element to its physical and chemical properties.

The statement of the law was given as:
1. The physical and chemical properties of the elements are periodic functions of their atomic weights.
2. Almost at the same time, German scientist Lothar Meyer came to the same conclusion by observing that physical properties of the elements are changing periodically with their atomic weights.

Wbbse Class 10 Physical Science Solutions

Periodic Table: On the basis of his periodic law, Mendeleev constructed a table by arranging the elements in increasing order of their atomic weights. This table is known as Mendeleev’s periodic table.

Modem Periodic Law: In 1913, British scientist Moseley while observing the effects of X-rays on metals, concluded that atomic number is a more accurate and fundamental property of an element than its atomic weight. Accordingly, Mendeleev’s periodic law was modified. This modified form is known as the modern periodic law which states.

The physical and chemical properties of the elements are the periodic functions of their atomic number. In the periodic table, the horizontal rows of elements are called periods and the vertical columns of elements are called groups. The modified version of Mendeleev’s periodic table has 7 periods and 9 groups.

Wbbse Class 10 Periodic Table Of Elements

Long form of periodic table or Modern periodic table: The periodic table constructed on the basis of the electronic configuration of elements is known as long form of periodic table or the modern periodic table. It is also known as Bohr’s periodic table. It consists of 7 periods and 18 groups.

Elements of different groups are given some special names depending on their properties. For example:

Group	Name	Example
Group 1 elements	Alkali metals	Li, Na. K, Rb
Group 2 elements	Alkaline earth metals	Be, Mg, Ca, Sr
Group 11 elements	Coinage metals	Cu, Ag, Au
Group 15 elements	Pnictogens	N, P, As, Sb
Group 16 elements	Chalcogens	0, S, Se, Te
Group 17 elements	Halogens	F, Cl, Br, 1
Group 18 elements	Noble gases	He, Ne, Ar, Kr

The position of hydrogen in the periodic table is controversial. It shows certain properties that resemble the properties of both group-1 (alkali metals) as well as group- 17 elements (halogens). For these reasons, it can be placed in group 1 as well as in group- 17. Due to this peculiar nature of hydrogen, it is often called the rogue element’.

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table Short And Long Answer Type Questions:

Question 1. State Dobereiner’s law of Triads’? Why was the law discarded?

Answer:

Dobereiner’s law of Triads’:

If the elements of a triad i.e., a group of three elements having similar chemical properties, are arranged in increasing order of their atomic weights, then the atomic weight of the middle element is found to be equal or almost equal to the average of atomic weights of the other two elements.

Example: Li (7), Na (23) and K (39) have similar chemical properties and thus forms a triad. The average of atomic weights of Li and K = 7+39/2 = 23, which is the atomic weight of Na.

The law failed to arrange most of the elements known till then. It was only applicable for few elements. For these reasons, the law was discarded.

Wbbse Class 10 Physical Science Solutions

Question 2. State Mendeleev’s periodic law. What was the basis of Mendeleev’s periodic law?

Answer:

Mendeleev’s periodic law:

1. The physical and chemical properties of the elements are periodic functions of their atomic weights.

2. Mendeleev examined the relationships between the atomic masses of the elements and their physical and chemical properties such as boiling point, melting point, density, atomic mass, the formula of hydrides and oxides of the elements etc. He observed that the physical and chemical properties of the elements change and periodically repeat with the change in atomic masses.

3. Based on these observations, he arranged the elements in the order of increasing atomic masses and formulated the periodic law.

Question 3. Evaluate the contribution of Lothar Meyer in the periodic classification of elements.
Answer:

The contribution of Lothar Meyer in the periodic classification of elements:

German scientist Lothar Meyer (1870) worked extensively on the relationship between physical properties of elements (such as atomic volume, melting point, boiling point etc.) and their atomic weights. He also stated a periodic law which was quite similar to what Mendeleev had proposed.

Wbbse Class 10 Physical Science Solutions

The work of both Mendeleev and Lothar Meyer are acknowledged towards periodic classification of elements. However, Mendeleev got greater recognition in this field as he put forward the theory much earlier than Lothar Meyer. However, the contribution of Lothar Meyer towards the concept of periodicity and periodic classification of elements cannot be neglected.

Wbbse Class 10 Periodic Table Of Elements

Question 4. What do you mean by period and group and groups were present in Mendeleev’s periodic table published in 1871?

Answer: The horizontal arrays of periodic table are termed as periods and the vertical arrays are termed as groups of the periodic table. There were 7 periods (1, 2, 3,…, 7) and 8 groups (1-8) in the periodic table published in 1871.

Question 5. How many elements were mentioned in the main version of Mendeleev’s periodic table published in 1871? Which element was absent in the table and why?

Answer:
1. 63 elements were mentioned in Mendeleev’s periodic table published in 1871.
2. Group ‘0’ (zero) was absent in Mendeleev’s periodic table (1871) as inert gases were not discovered at that time.

Wbbse Class 10 Physical Science Solutions

Question 6. Why were subgroups needed in Mendeleev’s periodic tab?

Answer: In Mendeleev’s periodic table it was found that except for group ‘0’ (zero) and group 8, several elements of different properties are placed in the same group. Each of the groups except 0 and 8 is divided into two subgroups designated as ‘A’ and ‘B’. Properties of the elements of subgroups A and B are altogether different, except their valencies.

However, elements of the same subgroup exhibit more or less similar properties. E.g., the alkali metals of Group l-A are closely alike. But Group IA metals differ remarkably from the coinage metals of Group IB, although they have a common valency of 1.

Question 7. The atomic mass of argon (39.94) is greater than that of potassium (39.1), yet potassium was placed after Argon in Mendeleev’s periodic table—explain why.

Answer: Mendeleev’s main objective was to place elements with similar properties in the same class. That is why these two elements were placed in the group where their properties resemble with other elements although their atomic masses indicate just the reverse places in the periodic table.

Wbbse Class 10 Periodic Table Of Elements

Question 8. How was Mendeleev’s periodic table useful in the periodic classification of elements? Explain with an example.

Answer: The elements present in the same group of the periodic table show similarity in properties.

Wbbse Class 10 Physical Science Solutions

1. So, if the property of one element in a specific group is known we can get an idea about the properties of the other elements present in the group.
2. If we know the properties of sodium of group- IA, the properties of other elements of that group namely potassium, rubidium and caesium can easily be evaluated.

Question 9. How was Mendeleev’s periodic table helpful in describing the structure and electronic configuration of different elements?
Answer: The elements present in the same subgroup of a group in Mendeleev’s periodic table have similar physical and chemical properties. Again, it is known that elements with similar properties have similar electronic configurations of their outermost orbits.

Thus, it can easily be correlated that elements of the same group have similar electronic configuration. Therefore, we can get an idea about the electronic configuration of other elements of a definite group, if the electronic configuration of any one element of the group is known. Similarly, from the knowledge of electronic configuration, we get an idea about the structure of atom.

Question 10. Discuss the limitations of Mendeleev’s periodic table.

Answer:

The limitations of Mendeleev’s periodic table:

Though Mendeleev was almost successful in arranging the elements on the basis of their properties in his periodic table, the periodic table had some serious drawbacks such as:

1. In some cases, in order to place the elements having similar properties in the same group, Mendeleev placed some elements with higher atomic weight before elements having lower atomic weight.
2. The position of hydrogen in Mendeleev’s periodic table was controversial.
3. In some cases, Mendeleev put elements having similar properties in different groups while elements having different properties in the same group.
4. Though the atomic weights of different isotopes are different, they were not given different positions in the periodic table.

Question 11. Mention the observation and inference of Moseley’s experiment on X-ray spectra.

Answer: In 1913, scientist Moseley made extensive research on the X-rays produced when cathode rays are bombarded on plates (target or anticathode) made of different metals.

Observation:
1. He observed that the X-rays produced due to collision of cathode rays with different metals have different frequencies. However, for a definite metal, the frequency of produced X-rays is constant.
2. He also observed when the square root of the frequencies of radiated X-rays was plotted against the atomic numbers of the elements (metals) used as anticathode, a straight line was obtained. However, no such straight line was obtained when atomic masses of the elements were plotted against the square root of the frequencies.

Inference: From the experiment, Moseley inferred that the properties of elements are a function of their atomic numbers and not their atomic masses.

Question 12. Write down the modified or modern periodic law. Mention the total number of periods and groups in the modern edition of Mendeleev’s periodic table.

Answer: The physical and chemical properties of the elements are a periodic function of their atomic numbers. There are 7 periods (1-7) and 9 groups (1, 8 and 0) in the modern edition of Mendeleev’s periodic table.

Question 13. Write down some differences between period and group in the periodic table.

Answer:

Differences between period and group in the periodic table are as follows:

Period	Croup
1. A period is a horizontal row of the periodic table.	1. A group is a vertical column of the periodic table.
2. There are 7 periods in the long form of the periodic table.	2. There are 18 groups in the long form of the periodic table.
3. The elements in a period differ in physical and chemical properties.	3. The elements in each group have some similar properties but not entirely identical properties.
4. Number of electronic shells remains the same for all elements in a period.	4. Number of electronic shells increases gradually from top to bottom in a group.

Question 14. The chemical properties of all the elements in the same period are not the same, but the chemical properties of the elements in the same group are similar— explain.

Answer: The chemical properties of the elements depend on the atomic number and hence on the electronic configuration of the atoms. Since the atomic number and electronic configuration of the elements in a particular period are different, their chemical properties are also different.

On the other hand, although the atomic numbers of different elements belonging to the same group are different, the electronic configurations of the outermost shell are similar. That is why their chemical properties are also similar to each other.

Question 15. Why are the elements arranged in increasing order of their atomic number in the long form of the periodic table, instead of their increasing atomic weight?

Answer: In 1913-14, British scientist Moseley, on the basis of his experiments on X-ray spectra of different atoms, explained that atomic number is a more fundamental property of an element than its atomic weight. It controls the physical and chemical properties of the elements. Again, after the discovery of isotopes, it was well-known that an element may contain atoms of different atomic weights. Thus, it was proved that atomic weight does not control the properties of an element. So, in the long form of the periodic table, elements are arranged in increasing order of their atomic number instead of atomic weight.

Question 16. Mention the differences between modern version of Mendeleev’s periodic table and the long form of the periodic table.

Answer: The major differences between modern version of Mendeleev’s periodic table and long form of periodic table are:

The modern version of Mendeleev’s periodic table	The long form of the periodic table
1. In the modern version of Mendeleev’s periodic table, the elements are arranged on the basis of atomic number.	1. In the long form of the periodic table, the elements are arranged on the basis of the electronic configuration of the atoms.
2. In Mendeleev’s periodic table, there were 7 periods and 9 groups and each group from 1 to 7 was divided into A and B subgroups.	2. In the long form of the periodic table, there are 7 periods and 18 groups and no group is divided into subgroups.

 

Question 17. Which period is termed as the shortest period? Mention the groups where the elements of the shortest period reside.

Answer: The 1st period of the periodic table is termed as the shortest period. It contains only two elements. Among the two elements of the 1st period, Hydrogen belongs to group 1 and helium belongs to group 18 of the periodic table.

Question 18. What are the first and second short periods of the periodic table? Mention the total number of elements in each of these periods. What are the first and last elements in each of these periods and mention the corresponding groups to which they belong?

Answer: Second and third periods of the periodic table are termed as first and second short periods respectively. There are 8 elements in each of the short periods.

Period	Name of the element	Group in which it is placed
First short period	Lithium (first element) Neon (last element)	Group 1 Group 18
Second short period	Sodium (first element) Argon (last element)	Group 1 Group 18

Question 19. What is meant by ‘representative elements’?

Answer:

‘Representative elements’:

The elements of the 2nd and 3rd periods are abundant in nature. Apart from this, the elements of a certain group in these periods show much resemblance in properties in regular intervals. So, these elements are collectively called ‘typical or representative elements’.

[Apart from these elements, 8 elements from 4th period: K(19), Ca(20) and Ga(31) to Kr(36), 8 elements from 5th period: Rb(37), Sr(38) and ln(49) to Xe(54), 8 elements from 6th period: Cs(55), Ba(56) and Th(81) to Rn(86) and Two elements from 7th period: Fr(87) and Ra(88) are also known as representative elements.]

Question 20. What are alkali metals and alkaline earth metals? Why are they called so?

Answer: The elements present in group-IA or group-1 of the periodic table, i.e., lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr) are collectively called alkali metals. The oxides, hydroxides and carbonates of these metals are strongly alkaline in nature. So, they are called alkali metals.

The elements present in group-2A or group-2 of the periodic table, i.e., beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) are collectively called alkaline earth metals.

Compounds of these elements are abundantly found in the earth’s crust and their oxides, hydroxides and carbonates are alkaline in nature. So, they are called alkaline earth metals.

Question 21. Write down the characteristics of alkali metals.

Answer:

Characteristics of alkali metals are:

1. Alkali metals are highly electropositive and they have a tendency to form cations by losing electrons from their valence shell.
2. Because of high reactivity these elements do not occur freely in nature.
3. Valence shell of alkali metals contains 1 electron and thus alkali metals are monovalent.
4. Oxides and hydroxides of alkali metals are strong bases and alkalies respectively.
5. Alkali metals form metallic hydroxides and hydrogen in reaction with water.
6. Alkali metals are good reducing agents. Ascending order of reducing the power of alkali metals is — Li < Na < K < Rb < Cs.

Question 22. Write down the alkaline earth metals.

Answer:

Characteristics of alkaline earth metals are:

1. Alkaline earth metals are highly electropositive and have a tendency to form cations by losing electrons from valence shell.
2. Because of high reactivity, these elements do not occur freely in nature. ‘
3. There are 2 electrons in the outermost shell of the alkaline earth metals and thus they are bivalent.
4. They are strong reducing agent. Ascending order of reducing power of alkaline earth metals is — Be < Mg < Ca < Sr < Ba .
5. Alkaline earth metals are present in, earth crust as oxides.
6. Alkaline earth metals form hydrogen and alkalies in reaction with water under suitable conditions.

\(\begin{aligned}
& \mathrm{Ca}+2 \mathrm{H}_2 \mathrm{O} \frac{\text { room }}{\text { temperature }} \mathrm{Ca}(\mathrm{OH})_2+\mathrm{H}_2 \uparrow \\
& \mathrm{Mg}+2 \mathrm{H}_2 \mathrm{O} \longrightarrow \mathrm{Mg}(\mathrm{OH})_2+\mathrm{H}_2 \uparrow
\end{aligned}\)

 

Question 23. What are bridge elements? Why are they called so?

Answer:

Bridge elements:

1. The inert gases (He, Ne, Ar, Kr, Xe and Rn) present in group-0 or group-18 of the periodic table are known as bridge elements.
2. These elements act as a bridge between strong electronegative halogens of group-17 and strong electropositive alkali metals of group-1 of the next period. So, these elements are called bridge elements.

Question 24. What are transition elements?
Answer:

Transition elements:

The elements from scandium (Sc) to copper (Cu) in the fourth period, yttrium (Y) to silver in the 5th period, lanthanum (La) and hafnium (Hf) to gold (Au) in the 6th period are known as transition elements due to some specific characteristics in their electronic configurations.

Question 25. Write down some characteristics or properties of transition elements.

Answer:
1. Except for mercury (Hg), all transition elements are solid metals.
2. Transition elements generally exhibit variable oxidation states and valencies.
3. They can form both ionic and covalent compounds.
4. Transition elements exhibit a very distinctive property of forming coloured coordination complexes.
5. A number of transition metals such as Cr, Mn, Fe, Co, Ni, Cu etc. and their compounds are used as catalysts.
6. Many transition metals form alloys.

Question 26. Which elements are known as pnictogens? Why are they named so?

Answer: The elements of group V.B. of Mendeleev’s periodic table or group 15 of the long form of the periodic table are known as pnictogens. E.g., nitrogen (N), phosphorous (P), arsenic (As), antimony (Sb), and bismuth (Bi).

The word pnictogen is derived from the Greek word pnicogens meaning ‘suffocating’ or ‘to choke’. The first member of group 15 is nitrogen which is a suffocating gas. That is why members of this group are known as pnictogens.

Question 27. What are chalcogens? Why are they called so?

Answer:

Chalcogens:

1. The elements of group 16 of the periodic table namely O, S, Se, Te and Po are collectively known as chalcogens.
2. The term ‘chalcogen’ means the producer of ores. Most of the metals are found as their oxide or sulphide ores in the earth’s crust. So, these elements are known as chalcogens.

Question 28. What are metalloids? Give example.
Answer:

Metalloids:

There are some elements which have certain characteristics common to both metals and nonmetals. These are called metalloids.
Examples of metalloids are arsenic(As ), antimony (Sb), and bismuth (Bi).

Question 29. What are trans-uranium elements? Mention two important characteristics of these elements?

Answer: The elements placed after uranium (₉₂U) in the 7th period of periodic table are known as transuranium elements.
1. All trans-uranium elements are radioactive.
2. All these elements are artificially synthesised.

Question 30. Isotopes of an element have not been given separate places in the periodic table— this is one of the major drawbacks of Mendeleev’s periodic table. Explain.

Answer: In Mendeleev’s periodic table, the elements were arranged in increasing order of their atomic weights. However, the isotopes of an element have different atomic weights. But, Mendeleev placed all the isotopes of an element in the same group and same period which violates the basic rule of periodic classification. So, it can be considered as a major drawback of Mendeleev’s periodic table.

But in modern periodic table, the elements are arranged in increasing order of their atomic numbers. As all the isotopes of an element have the same atomic number, they should be positioned in the same place in the periodic table according to modern periodic law. Hence, this is no longer a drawback in modern periodic table.

Question 31. Discuss three advantages of modern periodic table over Mendeleev’s periodic table.

Answer: Three advantages of long form of periodic table over Mendeleev’s periodic table are discussed below:

1. In Mendeleev’s periodic table the elements were arranged in increasing order of their atomic weights and later in increasing order of their atomic numbers. But in long form of periodic table, the elements are arranged on the basis of their electronic configurations.

Elements having same electronic configuration in their outermost shell are placed in the same group. As the chemical properties are related to the electronic configuration of the outermost shell, the periodicity of properties of the elements are described more effectively in long form of the periodic table.

2. In Mendeleev’s periodic table, each group except group 8 and group 0, was divided into subgroups. But the elements of different subgroups of a definite group are not very similar in terms of their properties. In the modern periodic table, the concept of subgroups has been eliminated. Elements of different subgroups have been placed in completely different groups. Which is more scientific.

3. In Mendeleev’s periodic table, each period of group 8 had three elements, namely, (Fe, Co, Ni), (Ru, Rh, Pd) and (Os, Ir, Pt). This defied the basic rule of periodic classification. In the long form of periodic table, these elements are placed in different groups.

Question 32. Which elements are halogens? Why are these named so?

Answer:
1. Elements of Gr-17 in the modern periodic table i.e. F, Cl, Br, I, At etc. are termed as halogens.
2. The name ‘halogen’ means ‘salt-producing’. Several salts can be formed by the reaction of metal with these elements. Salts containing F, Cl or I have been found in seawater. That is why these elements are termed as halogens. Some salts of halogen are—NaCI (common salt), CaF₂, NaBr, Kl etc.

Question 33. Write down the characteristics of halogen elements.

Answer:

Characteristics of halogen elements are:

1. Halogens are highly reactive and hence cannot be obtained in a free state in nature. Halogens forms halides with metals and nonmetals.
2. Among halogens, fluorine and chlorine are gaseous, bromine is liquid and iodine is solid. Astatine is the only radioactive halogen.
3. Halogens are highly electronegative (except Iodine). Descending order of electronegativity: F > Cl > Br > I.
4. Number of electrons in the valence shell of halogen elements is 7 and thushalogens are monovalent.
5. Halogens are non-metals and diatomic molecules.
6. Halogen elements are a strong oxidising agents. Descending order of oxidising power or halogens is— F > Cl > Br > I.

Question 34. Which elements form the noble gas family? Among noble gases which one is most chemically active?

Answer:
1. The Group zero of modified Mendeleev’s periodic table or Gr-18 of the long form of periodic table are designated as noble gases, e.g. He, Ne, Ar, Kr, Xe and Rn.
2. Among the noble gases (or inert gases) Xenon (Xe) is the most chemically reactive. State the differences in properties between alkali metals and halogens.

Question 35. State the differences in properties between alkali metals and halogens.

Answer:

The major differences in properties between alkali metals and halogens are as follows:

Property	Alkali metals	Halogens
Physical state	Soft solid substances can be cut into pieces with a knife.	The diatomic molecule can be solid, liquid or gas.
Conductance	Good conductor of heat and electricity.	Do not conduct heat or electricity.
Ionisation energy	In a period, the alkali metal has the least ionisation energy.	In a period, the halogen has the| highest ionisation energy.
Electronegativity	The alkali metal is the least electronegative in a particular period.	The halogen is the most electronegative in a particular period.
Oxidising and reducing property	Strong reducing agents.	Strong oxidizing agents.
Property	Alkali metals	Halogens
Reaction with acids	React violently with acid to produce hydrogen.	Generally do not react with acids.
Formation of  compounds	Form, ionic compounds with non-metals.	Form ionic compounds with metals and covalent compounds with non-metals.

Question 36. Write down the differences between representative elements and transition elements.

Answer:

Differences between and transition elements representative elements are as follows:

Representative elements	Transition elements
1. The elements present in groups 12 and groups 13-17 are known as representative elements.	1. The elements present in group 312 in the long form of the periodic table are known as transition elements.
2. Most of the representative elements show single valency.	2. Transition elements show multiple valencies.
3. The compounds of representative elements are usually colourless.	3. Most of the compounds formed by the transition elements are coloured.
4. Representative elements consist of some metals, nonmetals and metalloids.	4. All transition elements are metallic in nature.

Question 37. Why is hydrogen called an element?

Answer:

While arranging the elements in increasing order of their atomic weights, Mendeleev observed that hydrogen shows similarity with both alkali metals of group-1A as well as with halogens of group-8B. So, hydrogen was placed in both group 1A and group 7B. Mendeleev faced difficulty in ascertaining a definite position for hydrogen and so he named it a ‘rogue element’.

Question 38. Write down Some similarities between hydrogen and alkali metals.

Answer:

Some similarities between hydrogen and alkali metals

1. Like the alkali metals of Gr-1, hydrogen shows mono-valency, high reactivity and electropositive character.
2. Both hydrogen and alkali metals are reducing in nature.
3. Both hydrogen and alkali metals have strong affinity for non-metals.
4. Hydrogen forms H₂O with oxygen. Like hydrogen, alkali metals also form oxides e.g. Na₂0, K₂0 etc.
5. Like halogens, hydrogen reacts with halogens to form halide compounds.

Question 39. Write down some similarities in properties of hydrogen and halogen elements.

Answer:

Some similarities in properties of hydrogen and halogen elements:

1. Like halogens, hydrogen is diatomic, has high ionisation energy and non-metallic character.
2. Like hydrogen, halogens are mostly univalent.
3. Hydrogen combines with metals to form hydrides while halogens combine with metals to form halides.
4. Halogens may be obtained at the anode by electrolysis of molten metallic halides.

Question 40. Mention some limitations of long form of periodic table.

Answer:

Some limitations of long form of periodic table:

The long form of the periodic table successfully eliminated a number of drawbacks of Mendeleev’s periodic table. However, the long form to has a few limitations. Which are discussed below—

1. Even in long form of periodic table, the position of hydrogen is still controversial. Hydrogen cannot be placed in a definite group.
2. Depending on the electronic configuration, He should have been placed in group 2. But it is placed in group 18 with other noble gases.
3. Like Mendeleev’s periodic table, the lanthanides and actinides are placed separately in two different rows below the main periodic table. The elements could not be placed in individual groups in the periodic table.

Question 41. Mention the differences in the properties of alkali metals and hydrogen.
Answer:

The differences in properties of alkali metals and hydrogen are—

Alkali metals	Hydrogen
1. Most alkali metals are usually solid at room temperature.	1. Hydrogen is gaseous at room temperature.
2. Alkali metals are monoatomic.	2. Hydrogen molecules are diatomic.
3. Alkali metals are electropositive metals.	3. Hydrogen is an electropositive non-metal.
4. Alkali metals always form cations, e.g., Na⊕, and K⊕.	4. Hydrogen can form both cation (H⊕), and anion (HΘ).
5. Oxides of alkali metals are basic in nature.	5. Oxide of hydrogen (H2O) is neutral in nature.

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table Very Short Answer Type Questions Choose The Correct Answer:

Question 1. Which of the following sets of elements does not form a triad?
1. Ca, Sr, Ba
2. Li, Na, K
3. O, S, Se
4. Cl, Br, I

Answer: 3. O, S, Se

Question 2. The number of elements in the longest period and the incomplete period are respectively
1. 28,32
2. 32,28
3. 32,32
4. 18,28

Answer: 2. 32,28

Question 3. The elements that found no place in the original periodic table constructed by Mendeleev were
1. Transition metals
2. Alkali metals
3. Noble gases
4. Halogens

Answer: 3. Noble gases

Question 4. In Mendeleev’s periodic table, the arrangement of elements in the order of their increasing atomic weights was not maintained for which pair of elements?
1. C&N
2. S&Cl
3. K&Ar
4. K&Ca

Answer: 3. K&Ar

Question 5. The group having no subgroups in Mendeleev’s periodic table is
1. Group 2
2. Group 8
3. group 7
4. group 5

Answer: 2. Group 8

Question 6. Which group in Mendeleev’s period table contains more than one element in the same position?
1. Group 2
2. Group 3
3. Group 0
4. Group 8

Answer: 4. Group 8

Question 7. The number of elements in each position under group 8 is
1. 4
2. 5
3. 3
4. 2

Answer: 3. 3

Question 8. The pair of elements belonging to pnictogens is
1. N, P
2. P, O
3. F, Cl
4. Na, K

Answer: 1. N, P

Question 9. Moseley observed that the frequency (v) of X-rays produced when metal plates are bombarded with cathode rays is related to atomic number of elements (Z) as
1. v∝Z
2. √v∝ Z
3. v²∝Z
4. v = Z

Answer: 2. √v∝ Z

Question 10. The number of groups in the long form of the periodic table is
1. 9
2. 12
3. 13
4. 20

Answer: 3. 13

Question 11. In long form of periodic table, the elements are arranged on the basis of their
1. Atomic mass
2. Atomic number
3. Electronic configuration
4. Number of neutrons

Answer: 3. Electronic configuration

Question 12. In long form of periodic table, the alkali metals are placed in group
1. 1
2. 2
3. 3
4. 4

Answer: 1. 1

Question 13. In long form of periodic table, the alkaline earth metals are placed in group
1. 1
2. 2
3. 3
4. None of these

Answer: 2. 2

Question 14. In long form of periodic table, the halogens are placed in group
1. 15
2. 16
3. 17
4. 18

Answer: 3. 17

Question 15. In long form of periodic table, the coinage metals are placed in group
1. 11
2. 15
3. 17
4. 1

Answer: 1. 11

Question 16. In long form of periodic table, the chalcogens are placed in group
1. 16
2. 17
3. 13
4. 1

Answer: 4. 1

Question 17. Which of the following is a chalcogen?
1. Se
2. B
3. C
4. Ne

Answer: 1. Se

Question 18. In long form of periodic table, the noble gases are placed in group
1. 17
2. 18
3. 16
4. 15

Answer: 2. 18

Question 19. An example of a transition metal is
1. Ca
2. Mg
3. Fe
4. Al

Answer: 3. Fe

Question 20. In Mendeleev’s periodic table hydrogen can be placed
1. With halogens or with alkali metals
2. With chalcogens or with alkaline earth metals
3. With pnictogens or with coinage metals
4. With pnictogens or with alkaline earth metals

Answer: 1. With halogens or with alkali metals

Question 21. Which of the following is a coinage metal?
1. Iron
2. Silver
3. Aluminium
4. Nickel

Answer: 2. Silver

Question 22. A trans-uranium element is
1. Np
2. Ce
3. Lu
4. La

Answer: 1. Np

Question 23. The outermost shell of second period elements is
1. K-shell
2. L-shell
3. M-shell
4. N-shell

Answer: 2. L-shell

Question 24. Which of the following pair is not correct?
1. Coinage metal : Cu
2. Halogen: I
3. Alkaline earth metal: Fe
4. Inert element : Ne

Answer: 3. Alkaline earth metal: Fe

Question 25. Which one of the following is an alkali earth metal?
1. Rb
2. Ba
3. Sb
4. F

Answer: 2. Ba

Question 26. Total number of elements in the sixth period of modern long form of periodic table is
1. 2
2. 8
3. 18
4. 32

Answer: 4. 32

Question 27. Group 15 elements are called
1. Chalcogen
2. Pnictogan
3. Halogen
4. Aerogen

Answer: 2. Pnictogan

Question 28. To which group of the periodic table does ₁₃X²⁷ belong?
1. 14
2. 13
3. 15
4. 16

Answer: 2. 13

Question 29. Which of the following is not an alkali metal?
1. Rb
2. Cs
3. Ca
4. Na

Answer: 3. Ca

Question 30. Atomic number is more fundamental property of an element than atomic mass— the scientist who proved it is
1. Moseley
2. Lothar Meyer
3. Mendeleev
4. Dalton

Answer: 1. Moseley

Question 31. Which one is an alkali metal?
1. K
2. Mg
3. F
4. Ne

Answer: 1. K

Question 32. Which one of the following is not a transition element?
1. Fe
2. Co
3. Ca
4. Cr

Answer: 3. Ca

Question 33. Ca is
1. An alkali metal
2. An alkaline earth metal
3. Halogen
4. Inert element

Answer: 2. An alkaline earth metal

Question 34. Which one of the following is not a coinage metal?
1. Cu
2. Ag
3. Au
4. Fe

Answer: 4. Fe

Question 35. Number of rare earth elements is
1. 8
2. 12
3. 16
4. 14

Answer: 4. 14

Question 36. Which of the following elements represent the same period?
1. Li, Na, K
2. F, Cl, Al
3. Li, Mg, Ca
4. Cu, Ni, Zn

Answer: 4. Cu, Ni, Zn

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table Answer In Brief:

Question 1. The atomic weight of sodium is the average of the atomic weight of lithium and potassium. What can be inferred from this using the law of triads?
Answer: Using the law of triads, it can be inferred that the chemical properties of lithium, sodium and potassium are similar.

Question 2. Three elements A, B, and C obey Dobereiner’s law of triads. If the atomic weights of A and C are respectively 7 and 39, then find the atomic weight of B.
Answer: According to law of triads, the atomic weight of the middle element of the triad is the average of the atomic weights of the first and third elements.

Hence, atomic weight of B atomic weight of A + atomic weight of B = \(\frac{\text { atomic weight of } A+\text { atomic weight of } B}{2}\) =\(\frac{7+39}{2}\) = 23

Question 3. Newland’s law of octave is not applicable for which type of elements?
Answer: Newland’s law of octave is not applicable for heavy elements.

Question 4. On which basis did Mendeleev arrange the elements in the period table?
Answer: Mendeleev arranged the elements in the periodic table in increasing order of their atomic weights.

Question 5. How many groups in Mendeleev’s periodic table were divided into subgroups A & B?
Answer: In Mendeleev’s periodic table, seven (7) groups were divided into subgroups A & B.

Question 6. Which groups in Mendeleev’s periodic table (corrected version) do not contain any subgroup?
Answer: Zero (0) and Eighth (8) group.

Question 7. How many more elements are there in the second period than the first period of Mendeleev’s periodic table?
Answer: There are 6 more elements present in the 2nd period than the 1st period of Mendeleev’s periodic table.

Question 8. The initial form of Mendeleev’s periodic table does not contain noble gases. Why?
Answer: At the time when Mendeleev first published the periodic table, noble gases were not discovered yet. So, the noble gases did not find any place in the initial form of Mendeleev’s periodic table.

Question 9. What is the position of noble gases in modern version of Mendeleev’s periodic table?
Answer: In modern version of Mendeleev’s periodic table, the noble gases are placed in the ‘zero’ (0) group.

Question 10. The atomic masses of some elements were erroneous when Mendeleev constructed the periodic table. Why?
Answer: Valencies of the elements were incorrectly measured. These inaccurate valencies of elements were responsible for the erroneous values of atomic masses of some elements.

Question 11. Which element was predicted as eka-silicon by Mendeleev?
Answer: The element predicted by Mendeleev as eka-silicon is now known as germanium.

Question 12. Which element was predicted by Mendeleev as eka-aluminium?
Answer: Gallium was predicted by Mendeleev as eka-aluminium.

Question 13. What similarity is observed between Na and Cu regarding their positions in the periodic table?
Answer: Both Na and Cu were placed in group 1. Na was in group 1A and Cu was in group 1B.

Question 14. Which is the most electropositive nonradioactive alkali metal?
Answer: Caesium (Cs) is the most electropositive nonradioactive alkali metal.

Question 15. Which alkali metal is radioactive?
Answer: Francium (Fr) is a radioactive alkali metal.

Question 16. Name an alkali metal which is a liquid at 30°C.
Answer: Caesium is a liquid at 30°C

Question 17. Which alkaline earth metal is the heaviest?
Answer: Radium is the heaviest alkaline earth metal.

Question 18. Which alkaline earth metal is the lightest?
Answer: Beryllium is the lightest alkaline earth metal.

Question 19. Which is the radioactive halogen?
Answer: The radioactive halogen is astatine (At).

Question 20. Which is the lightest halogen?
Answer: Fluorine is the lightest halogen.

Question 21. Name a halogen that shows reducing property.
Answer: Iodine shows reducing property.

Question 22. Which group in the periodic table contains elements in all the three states, namely solid, liquid and gas?
Answer: Group 17 (group-8B) of the periodic table contains elements in all the three states. Fluorine and chlorine are gases, bromine is liquid and iodine is solid at normal temperatures.

Question 23. Which property is similar in all transuranium elements?
Answer: All the trans-uranium elements are radioactive elements.

Question 24. Mention the period in which all the elements are radioactive.
Answer: The elements of the 7th period are all radioactive.

Question 25. Name a radioactive noble gas.
Answer: A radioactive noble gas is radon.

Question 26. What are pnictogens?
Answer: The elements of group 15 in the long form of periodic table are known as pnictogens.

Question 27. Name the noble gas present in 2nd period of the periodic table.
Answer: Neon (Ne) is the noble gas present in 2nd period of the periodic table.

Question 28. Mention the series to Which U belongs.
Answer: Uranium (U) belongs to actinide series.

Question 29. Which period of the periodic table is known as very short period? To which groups do the elements of this period belong?
Answer: The first period of the periodic table is known as very short period. The period consists of two elements— hydrogen and helium. Hydrogen is placed in group-1A (group-1 in the long form of the periodic table). Helium belongs to group 0 (group 18 in the long form of periodic table).

Question 30. What are inert gases? Which one of them is chemically most active?
Answer: The six elements, helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon belonging to group-0 (group-18) of the periodic table are known as inert gases. Xenon is the most reactive among all the inert gases.

Question 31. What are coinage metals?
Answer: Three metals namely, copper (Cu), silver (Ag) and gold (Au) were used to make coins in ancient times as well as in medieval period. So, these metals are also known as coinage metals.

Question 32. How many periods are termed as longest period and mention the periods?
Answer: Two periods of the long form of periodic table are termed as longest periods. They are 6th period and seventh period.

Question 33. How many elements are there in each of the longest periods of the periodic table?
Answer: 32 elements.

Question 34. Mention the position of isotopes of elements in the periodic table.
Answer: Isotopes reside in the same group with the corresponding elements.

Question 35. Which period contains all of the three types of element viz. metals, non-metals and metalloids?
Answer: Fourth period.

Question 36. Why do elements of a particular group resemble in properties? Or, Why do Na and K resemble in properties?
Answer: Due to the presence of a similar type of electron configuration in the valence shell, they show resemblance in properties.

Question 37. Name an alkali metal.
Answer: Example of an alkali metal is sodium (Na).

Question 38. Which gas is produced in reaction with water and alkali metals?
Answer: Hydrogen gas.

Question 39. In which group of the modern periodic table do the alkaline earth metals reside?
Answer: Group 2.

Question 40. Name the halogen element which belongs to the same period where sulphur resides.
Answer: Chlorine (Cl).

Question 41. Name the liquid non-metal.
Answer: Bromine (Br).

Question 42. Which halogen element forms rapidly in reaction with water?
Answer: Fluorine (F).

Question 43. Which group of the modern periodic table belongs to the inert gases?
Answer: Group 18.

Question 44. Name the inert gas of the first period of the periodic table.
Answer: Helium (He).

Question 45. Name the inert gas residing at the 2nd period of the long form of periodic table.
Answer: Neon (Ne).

Question 46. State whether the transition elements are metals or non-metals.
Answer: Transition elements are metals.

Question 47. Which transition element is a liquid metal?
Answer: Mercury (Hg).

Question 48. Give example of noble metal.
Answer: Platinum (Pt)

Question 49. Elements of which group of the periodic table is termed as chalcogens?
Answer: Group-16 elements are termed as chalcogens.

Question 50. Mention the group of the rare earth elements in the periodic table.
Answer: Rare earth elements are present in the 3rd group of the modern periodic table.

Question 51. Name the lanthanide element of lowest atomic number.
Answer: Cerium (Ce) of atomic number 58.

Question 52. Mention two elements of the periodic table which are formed in laboratory.
Answer: Neptunium (Np, Atomic no: 93) and nobelium (No, atomic no: 102)

Question 53. Hydrogen and alkali metals form similar type of oxides. Mention a difference between those oxides.
Answer: Oxide of hydrogen ( H₂O) is a neutral oxide while oxides of alkali metals are basic oxides for e.g. Na₂O, K₂O etc.

Question 54. How many elements are there in shortest period of the periodic table?
Answer: 2

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table Fill In The Blanks:

Question 1. Cl,_____ and I are members of a triad.
Answer: Br

Question 2. Noble gases are placed at extreme _______ of the periodic table.
Answer: Right

Question 3. In the modern periodic table, the elements are arranged on the basis of their ______ instead of ______
Answer: Atomic Numbers, and Atomic weights

Question 4. Moseley from his X-ray experiment came to the conclusion that ______ is the fundamental property of elements.
Answer: Atomic number

Question 5. The _______ period of periodic table is an incomplete period.
Answer: 7th

Question 6. The elements in zero group are chemical ______
Answer: Inert

Question 7. Hydrogen shows similar properties with ________ as well as ______
Answer: Alkali metals, halogens

Question 8. Trans-uranium elements start from the element with atomic number ______
Answer: 93

Question 9. There is no _______in the long form of periodic table.
Answer: Subgroup

Question 10. The elements of A and B subgroups in a given group of Mendeleev’s periodic table show no similarity except in their _______
Answer: Valances

Question 11. The elements sodium, potassium etc., are called alkali metals because the _____ and _______ of these elements are strong bases.
Answer: Oxides, hydroxides

Question 12. Among the halogens, _________ is liquid at normal temperature.
Answer: Bromine

Question 13. The elements Mg, Ca etc., are called alkaline earth metals because these elements are found in the earth’s crust as _______ and their hydroxides are ________ in nature.
Answer: Oxides, Alkaline

Question 14. Apart from ______ all other transition elements are solid metals.
Answer: Mercury

Question 15. There was no mention of _________ in Mendeleev’s old periodic table.
Answer: Inert gases

Question 16. According to the modern form of Mendeleev’s periodic table, halogens belong to the _______ group.
Answer: 7B

Question 17. There are _______ groups in the long form of periodic table.
Answer: 18

Question 18. The number of elements in each of the 2nd and 3rd periods of the long form of periodic table is ________
Answer: 8

Question 19. __________ is a radioactive halogen.
Answer: Astatine

Question 20. Gold, platinum and silver are examples of ______ metal.
Answer: Noble

Periodic Table And Periodicity In The Properties Of Elements Topic A Periodic Table State Whether True Or False:

Question 1. Atomic weight of Na is the average of the atomic weights of Li and K.
Answer: True

Question 2. The law of octaves was found to be satisfactory in case of light elements.
Answer: True

Question 3. The physical and chemical properties of the elements are the periodic functions of their atomic weights.
Answer: False

Question 4. The elements after uranium are called transition elements.
Answer: False

Question 5. Except H, all other elements in group 1 are reactive metals.
Answer: True

Question 6. The initial form of Mendeleev’s periodic table does not contain noble gases.
Answer: True

Question 7. Beryllium is the lightest metal among the alkaline earth metals.
Answer: True

Question 8. All the group-17 elements in the periodic table exist in a gaseous state at normal temperature.
Answer: False

Question 9. Cu, Ag and Au are known as coinage petals.
Answer: True

Question 10. Zn is a transition element.
Answer: False

Question 11. Ne (neon) is the most electronegative element.
Answer: False

Question 12. Law of Triads was given by scientist Lothar Meyer.
Answer: False

Question 13. There was no subgroup in group 8 of Mendeleev’s periodic table.
Answer: True

Question 14. In group 6B of Mendeleev’s periodic table elements of all three states viz. solid, liquid, and gas is present.
Answer: False

Question 15. Halogens are members of Group zero (0) of the modern periodic table.
Answer: False

Question 16. Halogen means sea-salt producer.
Answer: True

Question 17. A number of transition elements in the 4th period is 10.
Answer: True

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements Synopsis:

1. Periodic property: The physical and chemical properties which show regular change while moving across a period from left to right and on going down a group from top to bottom are called periodic properties. Atomic radii, ionisation energy, electronegativity, oxidising and reducing properties etc., are periodic properties.

2. The variation of some physical and chemical properties of elements in a periodic table are as follows:

Property	Variation across a period	Variation down a group
Atomic radii	Decreases in moving from	Increases in moving from top to bottom
Ionisation energy	Increases in moving from left to right	Decreases in moving from top to bottom
Electronegativity	Increases in moving from left to right	Decreases in moving from top to bottom
Oxidizing power	Increases in moving from left to right	Decreases in moving from top to bottom
Reducing power	Decreases in moving from left to right	Increases in moving from top to bottom

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements Short And Long Answer Type Questions:

Question 1. What is meant by ‘periodic properties’?

Answer:

‘Periodic properties’:

With the increase in atomic number, physical and chemical properties of the elements belonging to different periods in the periodic table show a gradual change and as we move along a period, the change is systemic and periodic.

Similarly, the elements present in a group show the periodic variation of similar properties with the increase in atomic number down the group. The physical and chemical properties of elements which show regular and periodic change across a period and down a group are known as periodic properties. Atomic radius, ionisation energy, electronegativity, oxidising and reducing properties etc., are some of the common periodic properties.

Question 2. Define atomic radius.

Answer:

Atomic radius:

The distance from the centre of the nucleus to the outermost shell containing the electrons is called the atomic radius.
Since it is not possible to isolate a single atom for the measurement of its radius, atomic radius is expressed as covalent radius, metallic radius and van der Waal’s radius.

Question 3. Define covalent radius. If the inter-nuclear distance of the two atoms in a hydrogen molecule is 0.74A, then find its covalent radius.

Answer:
1. When two similar atoms are connected to each other by a single bond, one-half of the distance between their nuclei is called the covalent radius.
2. If the internuclear distance of the two atoms in a hydrogen molecule is 0.74 A, the covalent radius of hydrogen will be = 0.74/2 A = 0.37 A.

Question 4. How does atomic radius of the elements change from left to right in a period? Give reason.

Answer:
1. While moving from left to right across a period in the periodic table, the atomic radii or sizes of the atoms progressively decrease.
2. The principal quantum shell remains unchanged In the same period. So, the differentiating electrons enter the same shell. But due to an increase in the number of protons the value of +ve charge of the nucleus is increased. So attractive force of the nucleus for electrons in the outermost shell also increases which in turn gradually decreases the atomic radii.

Question 5. How does the atomic radius change down a group In the periodic table? Give reason.

Answer:
1. On moving down in any group of the periodic table, the atomic sizes rather the atomic radii of the elements increase.
2. On moving down a group, a new electronic shell is added to each succeeding element, though the number of electrons in the outermost shell remains the same. This tends to increase the atomic size.

At the same time, increase in nuclear charge with an increase in atomic number tends to decrease the atomic size. But the effect of the addition of a new electronic shell is so large that it outweighs the contractive effect of the increased nuclear charge. Hence, there occurs a gradual increase in atomic radii on moving down a group in the periodic table.

Question 6. (1) Arrange In decreasing order of atomic size—O, C, F, Li.
(2) Arrange in increasing order of atomic size—Br, F, I, Cl. lYz+VA

Answer:
(1)Given elements are members of the 2nd period of the periodic table. We know, across a period, atomic radius decreases from left to right. As a result, atomic size also decreases in the same fashion. So, the decreasing order of atomic size is Li > C > O > F.

(2)Given elements belong to Gr-17 of the long form of periodic table. We know, down the group atomic radius as well as atomic size gradually increases. So, the increasing order of atomic size is F < Cl < Br < I.

Question 7. What is the first ionisation energy of an element? Arrange the following elements in the increasing order of their first ionisation energy B, C, N, O and F.

Answer: The minimum amount of energy required to remove the most loosely bound electron from the outermost shell of an isolated gaseous atom of an element in its ground state to form a mono-positive ion is called the first ionisation energy of the element.

The ionisation energy of the given elements increases as— B<C<0<N<F.

Question 8. How does ionisation energy of elements change along a period and down a group in the periodic table?

Answer:
Variation along a period: On moving along a period from left to right, the ionisation energy of the elements generally increases, though some elements show exceptions to this trend. In a given period, the alkali metals present at the extreme left have the least ionisation energy while the noble gases at the extreme right have the highest ionisation energy.

Variation down a group: The ionisation energy of elements gradually decreases on moving down a group in the periodic table.

Question 9. (1)On which factors does the ionisation energy of an element depend?
(2) Why ionisation potentials of inert elements and alkali metals are very high and very low respectively?

Answer:
1. Factors influencing ionisation energy are:
(1)Atomic size: Ionisation energy decreases as the atomic size increases and vice-versa.
(2)Magnitude of nuclear charge: Ionisation energy increases with an increase in nuclear charge and vice-versa.
(3)Effect of the electronic configuration of the outermost shell: Ionisation energy increases with an increase in the stability of the electronic configuration.

2. Very high values of ionisation energies of inert gases are due to their highly stable electronic configuration. To remove an electron from this highly stable outer shell, a higher amount of energy is required making their ionisation energy so high.

On the other hand, alkali metals contain only one electron in their valence shell. By the removal of this electron, they can attain the stable electronic configuration of nearest inert gas. This is why the removal of an electron from alkali metals require a lower amount of energy making their ionisation energies low.

Question 10. Will ionisation energies of isotopes of an element be equal to each other? Explain.

Answer: Although the number of neutrons in the isotopes of an element are different, the number of electrons are same. As a consequence, electronic configurations are also the same. That is why the ionisation energies of these isotopes become equal to each other.

Question 11. Define electronegativity. Name the most electronegative and the least electronegative elements.

Answer:

Electronegativity:

Scientist Pauling defined electronegativity of an element as the tendency of an atom of the element to attract the bonded pair of electrons towards itself when it is covalently bonded to another atom, the most electronegative element is fluorine (F) and the least electronegative element is caesium (Cs).

Question 12. How does the electronegativity of the elements change across a period on moving from left to right? Explain with reasons.

Answer: On moving from left to right across a period the nuclear charge of the elements increases and atomic size decreases. As a result, the nuclear force of attraction on the bonded electron pair increases.

Hence, on moving from left to right the electronegativity of elements gradually increases. The alkali metal of a particular period located at the extreme left of the period is the least electronegative while the halogen in group 17 located at the right is the most electronegative element.

Question 13. How does the electronegativity of elements change on moving down a group? Arrange the following elements in the increasing order of their electronegativity. F, Br, Cl, I.

Answer: On moving down a group, the atomic size of the elements increases. As a result, the nuclear force of attraction on the outermost electrons decreases even though the nuclear charge increases. Hence, on moving down a group electronegativity of the elements gradually decreases.

The increasing order of electronegativity of the given elements: I < Br < Cl < F

Question 14. (1) Arrange the following elements in order of electronegativity—P, Cl, Si, S.
(2)Order of electronegativity of halogens— K Br < Cl < F. Explain.
(3)Name the elements with the highest and lowest electronegativity.

Answer:
(1)Given elements belong to the 3rd period in the periodic table, we know across a period from left to right electronegativity gradually increases. Therefore the increasing order of electronegativity of the given elements- Si < P < S < Cl.

(2)Halogens belong to group 17 of the periodic table. Now down a- group, electronegativity gradually decreases from top to bottom. Therefore we have the increasing order of electronegativity as given—I < Br < Cl < P.

(3) Fluorine (F) is the most electronegative element and Caesium (Cs) is the least electronegative element.

Question 15. Discuss how the oxidising and reducing power of elements change across a period and down a group in the periodic table.

Answer:

Variation across a period: On moving from left to right across a period, the atomic number as well as a number of protons gradually increases by one unit. At the same time, number of electrons also increases in the same outermost subshell. So, on moving from left to right nuclear attractive force on the outermost electrons gradually increases and the tendency to lose the electrons decreases.

At the same time, the tendency to accept electrons increases. As a result, the elements at the left part of the periodic table show a greater tendency to release electrons while the elements at the right part of the periodic table show a greater tendency to accept electrons. Hence, on moving from left to right across a period, the reducing power of the elements decreases whereas oxidising power increases. Alkali metals in group-1 of each period are the strongest reducing agents while the halogens in group-17 of each period are the strongest oxidising agents.

Vauation down a group: On moving down a group, atomic size of the elements increases and hence attractive force between the nucleus and the outermost electrons decreases. So, on moving down a group, the tendency of the elements to lose electrons increases. Hence, on moving down a group, reducing power of elements increases whereas oxidising power decreases.

Question 16. Discuss the variation of oxidising and reducing properties of the elements across a period. Arrange the following elements as per the given direction
(1)Na, Al, Si, Mg, P, Cl, S (decreasing order of reducing power).
(2)I, Cl, F, Br (increasing order of oxidising power).

Answer: On moving from left to right across a period, the oxidising power of the elements gradually increases while the reducing power gradually decreases.

(1)Decreasing order of reducing the power of the given elements: Na > Mg>AI>Si > P > S > Cl
(2)Increasing order of oxidising power of the non-radioactive halogen is: I < Br < Cl < F

Question 17. Arrange the elements of group-1A and group-2A in the correct order with respect to the following properties—
(1)Atomic size,
(2)Ionisation energy or ionisation potential,
(3) Reducing power.

Answer: For group-1A elements,
(1)Increasing order of atomic size: H < Li < Na < K < Rb < Cs < Fr
(2)Decreasing order of ionisation energy: H > Li > Na > K > Rb > Cs > Fr
(3)Increasing order of reducing power: H < Li < Na < K < Rb < Cs < Fr

For group-2A element,
(1)Increasing order of atomic size: Be < Mg < Ca < Sr < Ba < Ra
(2)Decreasing order of ionisation energy: Be > Mg > Ca > Sr > Br > Ra
(3)Increasing order of reducing power: Be < Mg < Ca < Sr < Ba < Ra

Question 18. Arrange the halogens in the correct order with respect to the following properties
(1) Atomic size,
(2) Ionisation energy or ionisation potential,
(3) Oxidising power.

Answer: Increasing order of atomic size: F < Cl < Br < I < At
Decreasing order of ionisation energy: F > Cl > Br > I > At
Decreasing order of oxidising power: F > Cl > Br > I > At.

Question 19. Some elements of a definite -period are given in the order in which they appear in the periodic table-Li, Be, B, C, Q, F, Me
1. State the period to which the elements belong.
2. Which element of that period is absent here? Mention its position.
3. Which element among these shows III catenation?
4. Arrange F, Be and N in increasing order of electronegativity.
5. Which element among these is a halogen?
6. Which one is an alkaline earth metal?

Answer:
1. All these elements belong to 2nd period.
2. The element absent here is nitrogen. In the periodic table, it is present between C and 0.
3. C shows catenation property.
4. Increasing order of electronegativity— Be < N < F.
5. F is a halogen.
6. Be is an alkaline earth metal.

Question 20. The names of the elements of a group in the periodic table are as follows: (1) boron, (2) aluminium, (3) gallium, (4) indium, (5) thallium. Answer the following questions:
1. Which of these elements shows the maximum metallic character?
2. Which element has the highest electronegativity?
3. If the electronic configuration of aluminium is K.= 2, L = 8, M = 3, then how many valence electrons are there in thallium?
4. The atomic number of boron is 5. Write the formula of the compound formed when boron reacts with chlorine.
5. Which of these elements has the least ionisation energy?
6. How many valence electrons are there in the elements present at the right side of this group?

Answer: The answers are given below:
1. The maximum metallic character is shown by thallium.
2. Boron has the highest electronegativity.
3. Thallium also has 3 valence electrons.
4. The formula of the compound formed when boron reacts with chlorine is  BCI₃.
5. Thallium has the least ionisation energy.
6. There are 4 valence electrons in the elements present at the right side of this group.

Question 21. The atomic number of three elements &V>B and C are (Z-2), Zand (Z+l) respectively. B is an inert gas. 0 Which of these elements is the most electronegative?
1. Which element has the highest ionisation energy?
2. State the nature of the compound formed by A and C and write its formula.

Answer: The atomic number of B is Z and it is an inert gas. Therefore, it has 8 electrons in its outermost shell. Hence, element A (atomic number -I- 2) has (8 – 2) = 6 electrons in its valence shell and element C (atomic number = Z+1) has 1 electron in its valence shell, i.e., C is an alkali metal. Therefore, we can conclude that

1. A is the most electronegative element.
2. B is an inert gas. Hence, it possesses the highest ionisation energy.
3. The number of electrons in the valence shell of A is 6. Hence, its valency is 8-6 = 2. Similarly, the valency of C is 1 as it has one electron in its outermost shell

Therefore, the formula of the compound formed between A and C will be C₂A. C is an alkali metal and A is an electronegative element. Hence, the nature of the compound formed by A and C will be electrovalent or ionic.

Question 22. The atomic number of the elements A, B and C are respectively 3,11 and 19.
1. State the position of these elements in Mendeleev’s periodic table as well as the long form of the periodic table.
2. Which one is the most reducing in nature?
3. State whether B will form a covalent or an ionic compound with chlorine.

Answer:
1. The position of the elements in the periodic table is given below:

Element	Electronic configuration				Position in Mendeleev’s periodic table		Position in long form of the periodic table
	K	L	M	N	Period	Croup	Period	Group
3A	2	1			2nd	IA	2nd	1
11B	2	8	1		3rd	IA	3rd	1
19C	2	8	8		4th	IA	4th	1

2. C is the strongest reducing agent.
3. The number of valence electrons in B is 1. Hence, it is an alkali metal. On the other hand, chlorine is a strongly electronegative element and a non-metal. So, B will form an ionic compound with chlorine.

Question 23. Mention the position of the following elements in the long form of the modern periodic table, 11^Na, 9^F,13^AI

Answer:

Element	Electronic configuration			Position in the long form of the periodic table
	K	L	M	Period	Group
11Na	2	8	1	3	1
13AI	2	8	3	3	13
9F	2	7		2	17

Question 24. Assume the symbol of two elements are A and B and their atomic numbers are 7 and 20 respectively. Predict their positions in Mendeleev’s periodic table.
Answer:

Element	Electronic configuration				Position in Mendeleev’s periodic table
	K	L	M	N	Period	Group
A(7)	2	5			2	5B
B(20)	2	8	8	2	4	2A

Question 25. Among Cl (17), Na (11), Mg (12), Ca (20), S (16) and F (9)
1. Which metallic elements belong to Group 2?
2. Which non-metallic elements belong to 3rd period of Mendeleev’s periodic table?

Answer:

Element	Atomic No.	Electronic configuration				Position in the periodic table		Metal or nonmetal
		K	L	W	N	Period	Group
Cl	17	2	8	7		3	7B	non-metal
Na	11	2	8	1		3	1A	metal
Mg	12	2	8	2		3	2A	metal
Ca	20	2	8	8	2	4	2A	metal
S	16	2	8	6		3	6B	non-metal
F	9	2	7			2	6B	non-metal

1. Mg (12) and Ca are the metals which belong to group 2 in Mendeleev’s periodic table.
2. S (16) and Cl (17) are the non-metals which belong to the 3rd period of Mendeleev’s periodic table.

Question 26. The number of electrons and neutrons in an atom of an element are 17 and 18 respectively. Predict .the position of the element in the Mendeleevas periodic table.

Answer: Number of electrons = 17
Electronic configuration =2,8,7
The elements will belong to 3rd period and Group 8 B in Mendeleev’s periodic table.

Question 27. If the electronic configuration of an element is 2, 8, 8, 1, then answer the following questions
1. Write down the number of periods and groups.
2. Write down the electronic configuration of another element of this group.

Answer:
1. The element belongs to the 4th period and group- 1 in the long form of modern periodic table.
2. Electronic configuration of another element of group-1 is 2, 8,1.

Question 28. A, B, and C are three different elements. B is an inert gas. Their atomic numbers are (Z -1), Z and (Z+l) respectively
1. Except B which one is a non-metal?
2. Which one is definitely a metal?
3. Predict the group-wise position of ‘A’ in the periodic table and the valency of A.

Answer:
1. ‘A’ is definitely a non-metal.
2. ‘C’ is definitely metal.
3. Element ‘A’ belongs to Group 7B in Mendeleev’s periodic table or Group 17 in the long form of the periodic table. Being an element of Gr 17, the valency of A is 1.

Question 29. The atomic numbers of the three elements A, B, and C are 6,8, and 10 respectively. C is an inert gas.
1. Which one is most electronegative?
2. Atomic size is least for which element?
3. In which group does element B reside?

Answer:
The electronic configuration of ₆A is K(2), L(A).
The electronic configuration of ₈B is K(2), L(6).
Electronic configuration of ₁₀C is K(2),L(8).

1. It is obvious from the electronic configuration that element B is the most electronegative.
2. Atomic size is least for the element B.
3. The element B belongs to the Gr-16 of the modern periodic table. It is a member of the 2nd period.

Question 30. The number of electrons in the outermost shell of elements X and Y are 2 and 7 respectively.
1. In which group of periodic table should the elements be placed?
2. Which one is an oxidising agent and which one is a reducing agent in case of forming compounds?

Answer:
1. ‘X’ should be placed in Gr-2 of the modern periodic table as the number of electrons in the valence shell is 2.
‘T’ should be placed in Gr-17 of the modern periodic table as the number of electrons, in the valence shell is 7.

2. ‘Y’ will act as an oxidising agent as it has a tendency to accept electrons.’X’ will act as reducing agent as it has a tendency to lose electrons.

Question 31. The atomic numbers of three elements A, B and C are 9, 10 and 11 respectively.
1. Which element among these is the most electronegative?
2. Metallic character is the highest for which one?
3. What will be the nature of the compound formed by the combination of A and C?

Answer:

Element	Atomic No.	Electronic configuration				Position in the periodic table
		K	L	M	N	Period	Group
A	9	2	7			2	17
B	10	2	8			2	18
C	11	2	8	1		3	1

1. ‘A’ is a most electronegative element.
2. ‘C’ has the highest metallic character.
3. The compound formed by the combination of A and C will be ionic in nature.

Question 32. The atomic numbers of three elements A, B, and C are 3,11 and 19 respectively.
1. Mention the group where the elements are placed in Mendeleev’s periodic table.
2. Metallic character is the highest for which one?
3. What is the nature of the compound formed in the reaction of B and Cl₂?

Answer:

Element	Atomic No.	Electronic configuration				Position in Mendeleev’s table
		K	L	M	N	Period	Group
A	3	2	1			2	IA
B	11	2	8	1		3	IA
C	19	2	8	8	1	4	IA

1. The elements belong to Gr-1A of Mendeleev’s periodic table.
2. Element ‘C has the highest metallic character.
3. The compound formed by the reaction of B and Cl₂ will be ionic in nature.

Question 33. The atomic numbers of the four elements A, B, C, and D are 3, 9/11 and 17 respectively. Which one is the most electropositive element and which one is the most electronegative element?

Answer:

Element	Atomic No	Electronic configuration			Position in the periodic table
		K	L	M	Period	Group
A	3	2	1		2	1
B	9	2	7		2	17
C	11	2	8	1	3	1
D	17	2	8	7	3	17

‘C’ is the most electropositive and ‘B’ is the most electronegative element among these elements.

Question 34. The atomic numbers of Elements A, B, C and D are 3, 9,11 and 19 respectively.
1. Which one has the highest reducing property? ‘
2. Electronegativity of which one is the highest?

Answer:

Element	Atomic No.	Electronic configuration				Position in the periodic table
		K	L	M	N	Period	Group
A	3	2	1			2	1
B	9	2	7			2	17
C	11	2	8	1		3	1
D	19	2	8	8	1		1

 

1. Reducing the property of D is the highest.
2. Electronegativity of B is the highest.

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements Very Short Answer Type Questions Choose The Correct Answer:

Question 1. Moving across a period from left to right, the atomic radii of elements
1. Gradually increases
2. Gradually decreases
3. Remains unchanged
4. Initially increases then decreases

Answer: 2. Gradually decreases

Question 2. On going down a group, the atomic radii of elements
1. Gradually decreases
2. Gradually increases
3. Remains unchanged
4. Initially increases then decreases

Answer: 2. Gradually increases

Question 3. In a given period, with decreasing atomic size of elements, the ionisation energy
1. Increases
2. Decreases
3. Remains unchanged
4. Either increases or decreases

Answer: 1. Increases

Question 4. In a given period, with an increase in nuclear charge of the atom of an element, the ionisation energy
1. Decreases
2. Increases
3. Remains unchanged
4. Either increases or decreases

Answer: 2. Increases

Question 5. The unit of electronegativity is
1. eV
2. erg
3. dyne
4. It is a unitless quantity

Answer: 4. It is a unitless quantity

Question 6. In a given period, the electronegativity is minimum for the elements of
1. Group-1
2. Group-2
3. Group-14
4. Group-17

Answer: 1. Group-1

Question 7. In a given period, the electronegativity is maximum for the elements of the group
1. 16
2. 17
3. 18
4. 14

Answer: 2. 17

Question 8. In a given period, with an increase in the atomic size of an element, its electronegativity will
1. Increase
2. Decrease
3. Increases or decrease
4. Remain unchanged

Answer: 2. Decrease

Question 9. For the elements of 2nd period, the correct order of atomic radii is
1. B > C > N > 0
2. B > N > C > 0
3. B > 0 > N > C
4. N > 0 > B > C

Answer: 1. B > C > N > 0

Question 10. The correct order of atomic radii for alkali metals is
1. Na < Rb < K < Cs
2. Na < K < Rb < Cs
3. K < Na < Cs < Rb
4. Na < Cs < K < Rb

Answer: 2. Na < K < Rb < Cs

Question 11. The least electronegative non-radioactive element is
1. Na
2. K
3. Rb
4. Cs

Answer: 4. Cs

Question 12. For the elements of 2nd period, the correct order of ionisation energy is
1. B < C < N >0
2. B < C < O < N
3. C < B < N > 0
4. N > C < B < 0

Answer: 1. B < C < N >0

Question 13. The correct order of ionisation energy for group-1 elements is
1. Li > K < Rb > Na
2. Li > Na > Rb > K
3. Li > Na > K > Rb
4. Na < Li > Rb > K

Answer: 3. Li > Na > K > Rb

Question 14. The correct order of electronegativity for the following element is
1. Li < Be < C < B
2. Li <Be < B < C
3. Li < B > Be > C
4. Be > Li < B < C

Answer: 2. Li <Be < B < C

Question 15. The correct order of reducing power of the elements in 3rd period is
1. Al < Si > P > S
2. Al > Si > S > P
3. Al > Si > P > S
4. Si < Al > S < P

Answer: 3. Al > Si > P > S

Question 16. The correct order of oxidising power of the halogens is
1. F > Br > Cl > I
2. Cl > F > Br > 1
3. Cl < F < Br < I
4. F > Cl > Br > I

Answer: 4. F > Cl > Br > I

Question 17. Most electronegative element is
1. 0
2. CL
3. F
4. Ne

Answer: 3. F

Question 18. Which one is more electronegative?
1. Na
2. K
3. Rb
4. Cs

Answer: 1. Na

Question 19. Which one of the following has the highest reducing ability?
1. K
2. Na
3. Li
4. Cs

Answer: 4. Cs

Question 20. Which halogen element has the least electronegativity?
1. F
2. Cl
3. Br
4. I

Answer: 4. I

Question 21. Which one is not a periodic property?
1. Oxidising property
2. Reducing property
3. Radioactivity
4. Atomic radius

Answer: 3. Radioactivity

Question 22. Which one of the following has the highest atomic radius?
1. K
2. H
3. Li
4. Na

Answer: 1. K

Question 23. The most oxidising element in second period of the periodic table is
1. F
2. Cl
3. Be
4. O

Answer: 1. F

Question 24. Which of the following alkali metals has the highest atomic radius?
1. Li
2. Na
3. Rb
4. Cs

Answer: 4. Cs

Question 25. Which of the following indicates the atomic number of a Gr-!A element?
1. 13
2. 17
3. 11
4. 10

Answer: 3. 11

Question 26. Which one of the following has the least electronegativity?
1. Ca
2. Mg
3. Sr
4. Ba

Answer: 4. Ba

Question 27. The correct order of reducing power is
1. Na < Al < Si < Mg
2. Na < Mg < Al < Si
3. Si < Al < Na < Mg
4. Si < Al < Mg < Na

Answer: 4. Si < Al < Mg < Na

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements Answer In Brief:

Question 1. What is meant by ‘periodicity of elements’?
Answer:

‘Periodicity of elements’:

The recurrence of properties of the elements at definite intervals when arranged in increasing order of their atomic numbers is called the periodicity of elements.

Question 2. In the case of the 4th period, after how many elements are properties repeated?
Answer: After 18 elements, in the 19th element properties are repeated.

Question 3. What is meant by the term ‘atomic radius? Classify atomic radius.
Answer: It is the distance between the centre of the nucleus and the outermost orbit of an atom. Atomic radius is classified as:
1. Covalent radius,
2. Metallic radius and
3. Van der Waals’ radius.

Question 4. Define metallic radius.
Answer:

Metallic radius:

Half of the distance between the nuclei of two adjacent metal atoms in a metallic crystal is called the metallic radius.

Question 5. What is van der Waals’ radius?
Answer: Half of the distance between the nuclei of two non-bonded neighbouring atoms of two adjacent molecules of the same substance in solid state is known as van der Waals’ radius.

Question 6. Arrange van der Waals’ radius, metallic radius and covalent radius in decreasing order of their values.
Answer: The decreasing order of the three atomic radii is as follows: van der Waals’ radius > metallic radius > covalent radius.

Question 7. The elements of which group in a given period of the periodic table have the largest atomic radii?
Answer: The elements of group 1 in a given period of the periodic table have the largest atomic radii.

Question 8. Whose atomic radius is smaller—Li or F?
Answer: Atomic radius of F is smaller.

Question 9. Which of the following has the smaller atom — Mg and Al?
Answer: Atoms of Al are smaller than that of Mg.

Question 10. Arrange Be, Na, Ar, S in order of atomic radius.
Answer: Ar > Na > S > Be .

Question 11. Arrange Al, Mg, Cl, Na, P, S and Si in decreasing order of their atomic radii.
Answer: Decreasing order of atomic radii: Nai > Mg > Al > Si > P > S > Cl.

Question 12. How does ionisation energy change across a period?
Answer: From left to right, across a period, the ionisation energy of elements gradually increases.

Question 13. The elements of which group in a given period of the periodic table have the lowest ionisation energy?
Answer: The elements of group 1 in a given period of the periodic table have the lowest ionisation energy.

Question 14. Name the element with highest ionisation energy.
Answer: Helium (He)’.

Question 15. Name the element with the lowest ionisation energy.
Answer: Caesium.

Question 16. Arrange the following elements in increasing order of ionisation energy—Li, Rb, K and Na.
Answer: Rb < K < Na < Li.

Question 17. What is the relation between atomic radii and electronegativity of the elements in a period?
Answer: Smaller the atomic radii of elements in a period, greater will be the electronegativity.

Question 18. Arrange F, Cl, Br and I in decreasing order of electronegativity.
Answer: Decreasing order of electronegativity: F > Cl > Br > I.

Question 19. How does the electronegativity of the Gr-1 elements change from top to bottom?
Answer: From top to bottom of group 1, electronegativity gradually decreases.

Question 20. Mention the position of the most electronegative element in the periodic table.
Answer: The most electronegative element F belongs to the 2nd period and Gr-17 in the periodic table.

Question 21. Arrange Na, Mg, Al, Cl and P is increasing order of electronegativity.
Answer: Na < Mg < Al < P < Cl.

Question 22. Arrange the following elements in increasing order of electronegativity—Na, O, N, F.
Answer: Na < N < 0 < F.

Question 23. Arrange Li, Be, B, N in decreasing order of electronegativity.
Answer: N > B > Be > Li.

Question 24. What is the relation between ionisation energy and electronegativity?
Answer: The ionisation energy of an element increases with the increase in electronegativity.

Question 25. What is the relation between ionisation energy and the reducing property of an element?
Answer: Lower the ionisation energy of an element, the stronger will be its reducing power.

Question 26. What is the relation between electronegativity and oxidising property of an element?
Answer: Higher the electronegativity of an element, the stronger will be its oxidising power.

Question 27. An element A is positioned in group 13 of the periodic table. What is the formula of its chloride?
Answer: The valency of all elements of group 13 is 3. Hence, the formula of the chloride of element A will be ACI₃.

Question 28. Which elements among the halogens exhibit highest and lowest metallic property?
Answer: Among the halogens, iodine exhibits highest metallic properties while fluorine exhibits the lowest metallic properties.

Question 29. Arrange Si, Mg, P, Al, Na, S and Cl in increasing order of their oxidising power.
Answer: Increasing order of oxidising power: Na < Mg < Al < Si < P < S < Cl.

Question 30. Arrange Na/ K, Rb and Cs in increasing order of electro-positivity.
Answer: Increasing order of electro-positivity: Na < K < Rb < Cs.

Question 31. Elements of which group in the periodic table has the highest oxidising power?
Answer: Elements of group 17 of the periodic table have the highest oxidising power among the elements of the corresponding period.

Question 32. Name the strongest oxidising agent of 2nd period.
Answer: Fluorine (F) is the strongest oxidising agent of 2nd period.

Question 33. Name the element with highest reducing ability of Gr-1 in the long form of the periodic table.
Answer: Lithium (Li).

Question 34. Arrange the halogens in increasing order of reducing power.
Answer: F < Cl < Br < I.

Question 35. Arrange in decreasing order of reducing property—K, Na, Li, Rb.
Answer: Rb < K < Na> Li.

Question 36. Arrange Cl, Br, I, F in increasing order of oxidising property.
Answer: I < Br < Cl < F.

Question 37. Arrange in decreasing order of oxidising property—Al, O, F, Cl.
Answer: F > O > Cl > Al.

Question 38. Among Na, K, O and N which one is the strongest oxidising agent and strongest reducing agent?
Answer: Strongest oxidising agent is O and strongest reducing agent is K.

Question 39. Give an example of a halogen that has reducing property.
Answer: Iodine.

Question 40. Arrange the following elements in increasing order of metallic character- 11 Na, K, Rb, Cs and 0 Br, F, Cl, I.
Answer: Na < K < Rb < Cs and 0 F < Cl < Br < I

Question 41. Arrange the following in increasing order of metallic character—Mg, Na, Al, S, Si, Cl, P.
Answer: Cl < S < P < Si < Al < Mg < Na.

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements Fill In The Blanks:

Question 1. _________is not a periodic property of elements.
Answer: Radioactivity

Question 2. In a given period the ______ have the smallest atomic radii.
Answer: Halogens

Question 3. The atomic radii of noble gases in a given period is ________ than the atomic radii of halogen atoms.
Answer: Greater

Question 4. The minimum energy required to remove the most loosely bound electron in the outermost orbit of the atom of an element is known as ________ energy of the element.
Answer: Ionisation

Question 5. The electronegativity of elements is measured in ______ scale.
Answer: Pauling

Question 6. The element _________ is the strongest oxidising agent.
Answer: Fluorine

Question 7. The higher the _________ of an element, the stronger will be its oxidising power.
Answer: Electronegativity

Question 8. Noble gases being ________ their covalent radii cannot be measured.
Answer: Mono-atomic

Question 9. The greater the tendency of an element to accept an electron greater will be its property and the greater the tendency of an element to release an electron, the greater will be its _______ property.
Answer: Oxidising, reducing

Question 10. The atomic radius of Al is ________ than SI.
Answer: Greater

Question 11. The ionisation potential of oxygen is _______ than that of nitrogen.
Answer: Less

Question 12. The ________ power of Na is less than that of Si.
Answer: Oxidising

Question 13. The reducing power of caesium is _______ than that of rubidium.
Answer: More

Question 14. The magnetic property of elements is not a ________ property.
Answer: Periodic

Question 15. Across a period from left to right atomic radius of elements gradually _______
Answer: Decreases

Question 16. Among all elements, atomic radius is the least in the case of _______
Answer: Hydrogen

Question 17. Among Al, S, Si and P, element with the highest electronegativity is _________
Answer: S

Question 18. The most reactive non-metal is _______
Answer: Fluorine

Periodic Table and Periodicity in the Properties of Elements Topic B Chemical Periodicity Of Elements State Whether True Or False:

Question 1. A covalent radius is always greater than a metallic radius.
Answer: False

Question 2. The correct order of atomic size is— F < Cl < Br < I < At
Answer: False

Question 3. The greater the size of an atom of an element, the greater will be its electronegativity.
Answer: False

Question 4. In the measurement of the electronegativity of different elements by the Pauling scale, the electronegativity of hydrogen is taken as 2.1.
Answer: True

Question 5. The increasing order of electropositivity is Na < K < Rb < Cs
Answer: True

Question 6. Mendeleev emphasized mainly on chemical and physical properties of elements while constructing the periodic table.
Answer: True

Question 7. Down a group atomic size of an element gradually decreases.
Answer: False

Question 8. Ionisation energy decreases with an increase in atomic size.
Answer: True

Question 9. Down a group metallic character decreases and oxidising power increases.
Answer: False

Question 10. Across a period, from left to right oxidising property decreases.
Answer: False

Question 11. Halogen elements are strong reducing agents.
Answer: False

Question 12. Order of reducing power— F > Cl > Br > I.
Answer: False

Question 13. If the number of electrons in the valence shell is 3, then the position of element in the long form of periodic table will be in group 13.
Answer: True

Periodic Table and Periodicity in the Properties of Elements Miscellaneous Type Questions Match The Columns:

Question 1.

Column A	Column B
Transition element	1. He
Coinage metal	2. H2
Inert gas	3. Fe
Rogue element	4. Ag

Answer:
Transition element:  3. Fe
Coinage metal: 4. Ag
Inert gas:  1. He
Rogue element: 2. H₂

Question 2.

Column A	Column B
Alkaline earth metal	1. Te
Chalcogen	2. Rb
Representative element	3. Xe
Bridge element	4. Ra

Answer:
Alkaline earth metal: 4. Ra
Chalcogen: 1. Te
Representative element: 2. Rb
Bridge element: 3. Xe

Question 3.

Column A	Column B
Highest electronegativity	1. Fe
Highest reducing power among group-1 elements	2. F
Variable oxidation state	3. He
Highest  ionisation energy	4. Cs

Answer:
Highest electronegativity:  2. F
Highest reducing power among group-1 elements: 4. Cs
Variable oxidation state: 1. Fe
Highest  ionisation energy: 3. He

WBBSE Solutions for Class 10 Physical Science and Environment

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds Synopsis:

The force of attraction which holds two of more constituents (atoms, ions, etc.) together in different chemical species is known as a chemical bond. It is formed either by losing or gaining electrons or by sharing pairs of electrons.

Octet rule: To attain chemical stability, an atom tends to acquire stable electronic configuration of the nearest noble gas, i.e., tends to attain 8 electrons in its outermost shell either by gaining or losing electrons or by sharing pairs of electrons. This tendency to acquire 8 electrons in the outermost shell of the atoms is called octet rule. There are many compounds where the constituent atoms do not follow the octet rule.

Wbbse Class 10 Physical Science Notes

Duplet rule: The elements near helium (H, Li, Be ) in the periodic table try to attain the stable electronic configuration of He by gaining, losing or sharing electrons in their outermost shells. This tendency to acquire 2 electrons in the outermost shell of atoms is called duplet rule.

Read and Learn Also WBBSE Solutions for Class 10 Physical Science and Environment

Electrovalency: To achieve stable electronic configuration of the nearest noble gas, some atoms completely give up one or more electrons from their outermost shell to form stable cations while some other atoms gain these electrons to form stable anions.

These two oppositely charged ions then combine together through the electrostatic force of attraction to form chemical compounds. The capacity of the elements for such chemical combination is known as electrovalency or ionic valency.

The bond formed between the atoms due to electrovalence is called an electrovalent bond or ionic bond and the compound formed in this way is called electrovalent or ionic compound.

Characteristics of ionic compounds: Ionic compounds do not exist as discrete molecules. They are solid crystalline substances with high melting and boiling points. They do not conduct electricity in solid state but are good conductors in a molten state or in an aqueous solution. Ionic compounds are soluble in polar solvents and react very fast in solution.

The term ‘molecular mass’ is not applicable for ionic compounds as they do not exist in molecular form. Instead, the formula mass of these compounds is determined. For covalent compounds, both molecular and formula mass are the same.

Wbbse Class 10 Physical Science Notes

In NaCI crystal, a large number of Na⁺ and Cl^– ions remain closely packed by a strong electrostatic force of attraction. This leads to the formation of stable three-dimensional crystal lattices.

Electron structure of NaCI:

Electron structure of CaO:

Electron structure of MgCI:

Examples of electrovalent compounds include calcium oxide, magnesium chloride, sodium fluoride, lithium hydride, etc. In it, the outermost shell of each constituent atom contains one electron.

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds Short And Long Answer Type Questions:

Question 1. Why do atoms of different elements react With each other?

Answer: The electronic configuration of the noble gases are very stable. To attain the electronic configuration of the nearest noble gases, atoms of different elements react with each other.

Question 2. Define chemical bond.

Answer:

Chemical bond:

The force of attraction which holds two or more constituents (atoms, ions etc.) together, different chemical species is .called chemical bond. It is formed either by losing or gaining electrons or by sharing pairs of electrons by atoms.

Wbbse Class 10 Physical Science Solutions

Question 3. What is the octet rule?

Answer:

Octet rule:

To attain chemical stability, an atom tends to acquire stable electronic configuration of the nearest noble gas either by gaining or losing electrons or by sharing pairs of electrons of its outermost shell. This tendency to acquire 8 electrons in the outermost shell of an atom is called the octet rule.

Question 4. state the limitations of octet rule.

Answer:

Octet rule has certain limitations. These are:

1. In some compounds, the central atom has more or less than 8 electrons, yet the molecules are highly stable. The Octet rule fails to explain this.
2. Octet rule is based on the stability and inertness of noble gases. But some noble gases specifically Xe form compounds with oxygen and fluorine.
3. This rule cannot explain the relative stability of different molecules.

Question 5. The octet rule is not always followed in case of ionic compounds. Give examples.

Answer:

The octet rule is not always followed in case of ionic compounds

The stability of ionic compounds cannot be always explained by octet rule. In case of transition metals, it is observed that octet rule is not followed. For example, in FeCl₂, Fe²⁺ ion has 14 electrons in its valence shell, and in FeCI₃, the Fe³⁺ ion has 13 electrons in its valence shell.

In both cases, octet rule is violated. Again, in LiH, h u+ ion and H~ ion have 2 electrons in their valence shells. Hence, octet rule is not valid for all ionic compounds.

Question 6. Show with an example that compounds can be formed without fulfillment of octet of the corresponding ions.

Answer: The stability of all ionic compounds can not be explained with the concept of sharing 8 electrons in the valence shell of the constituting ions, or the octet rule, e.g., in LiH the constituting ions Li^⊕ and H^Θ both have 2 electrons in their valence shell.

Moreover, in case of elements showing variable electrovalency, especially in transition metals octes rule is not obeyed in many instances. E.g., in FeCl₂ or FeCI₃, octet rule is not obeyed in Fe²⁺ and Fe³⁺.

Question 7. What is duplet rule?

Answer:

Duplet rule:

To attain chemical stability an atom of an element nearer to He (H, Li, Be) tends to acquire stable electronic configuration like He either by gaining or losing electrons or by sharing electron pairs of its outermost shell. This tendency to acquire 2 electrons in the outermost shell of an atom is called duplet rule.

Wbbse Class 10 Physical Science Solutions

Question 8. Which type of chemical bond is formed by complete transfer of electron from one atom to another atom? Explain with an example.

Answer:
1. Ionic bond or electrovalent bond is formed by complete transfer of electron from one atom to another atom.

2. Magnesium (electronic configuration: 2, 8, 2) loses 2 valence electrons to form an Mg²⁺ ion and attains the stable electronic configuration of nearest inert gas. On the otherhand oxygen (electronic configuration: 2, 6) accepts these two electrons to form O^2- ion and attains the stable electronic configuration of the nearest inert gas.

Therefore when Mg and O atoms combine, the former transfers its valence electron to the latter resulting in the formation of Mg²⁺ and O^2- respectively. The two ions are then held together by strong electrostatic forces of attraction to form MgO.

Question 9. Define electrovalency.

Answer:

Electrovalency:

To achieve the stable electronic configuration of nearest noble gas, some atoms give up one or more electrons from the outermost shell completely to form stable cations while some other atoms gain these electrons to form stable anions. These two types of oppositely charged ions combine together through electrostatic force of attraction to form chemical compounds. The capacity of the elements for such chemical combination is known as electrovalency.

Question 10. How is the electrovalency of an atom measured? Explain with example.

Answer:
1. The number of electrons gained or released by an atom during the formation of an ionic compound is measured as the electrovalency of that atom.
2. For example, during formation of NaCl, a Na – atom releases one electron to form a Na⁺  ion while a Cl – atom accepts that electron and converts it to Cl^– ion. Hence, the electrovalency of both sodium and chlorine in NaCl is 1.

Question 11. Ionic bonds are not real bonds- Justify the statement.

Answer:

Ionic bonds are not real bonds:

In covalent compounds, two atoms share one or more pairs of electrons with each other to form bonds having definite directions. But, in electrovalent compounds, the bond between two atoms is formed due to the strong electrostatic force of attraction between the cations and the anions. Here, no electron pair is shared by the involved ions. That is why ionic bonds are not considered as real bond.

Question 12. State two important characteristics of ionic bond.

Answer:

An ionic bond shows the following characteristics:

1. An ionic bond is formed between atoms of an electropositive and an electronegative element.
2. Ionic bond is in fact an electrostatic force of attraction between oppositely charged ions. So, sometimes, it is not considered, as a real bond.

Wbbse Class 10 Physical Science Solutions

Question 13. Write down some properties of ionic compound.

Answer:

Some properties of ionic compound:

1. Ionic compounds form crystals of definite geometrical shape.
2. The melting and boiling points .of ionic compounds are generally very high.
3. Ionic bond is non-directional in nature.
4. Ionic compounds do not exhibit isomerism.
5. Ionic compounds do not conduct electricity in the solid state, but in the molten state or in solution in a suitable solvent (like water) conduct electricity.
6. Ionic compounds generally dissolve in polar solvents.

Question 14. Why any ionic compound possesses high boiling point and melting point?

Answer: In ionic compounds, the oppositely charged ions are held together by strong electrostatic force of attraction. The mutual attraction between large number of oppositely charged ions results in the formation of large crystal lattices. Separation of ions from these crystal lattices requires large amount of energy. Hence, ionic compounds possess high melting and boiling points.

Question 15. Discuss the physical state of ionic compounds.

Answer:

The physical state of ionic compounds:

At normal temperature and pressure all ionic compounds are crystalline solids. This is because the oppositely charged ions are held together by strong electrostatic force of attraction which results in the formation of regular three-dimensional network called crystal lattice. Ionic compounds do not exist as discrete molecules.

Question 16. Why are ionic compounds hard in nature?

Answer: In ionic compounds, large number of oppositely charged ions are held together by strong electrostatic force of attraction which results in the formation of large crystal lattices. The ions remain closely packed in the crystal and the intermolecular space between them is very small. The hardness of an ionic compound is due to the aggregation of large number of ions in a very small space.

Question 17. Ionic compounds conduct electricity in molten state or in aqueous solutions but are non-conductors in the solid state. Explain with reason.

Answer: In solid state, the cations and anions in an ionic compound are held together by a strong electrostatic force of attraction. Hence, the ions are immobile and the compound fails to conduct electricity in solid state.

However, when the compound is in a molten state or is dissolved in water, the crystal lattice breaks, and the ions are separated. Due to presence of such mobile ions, ionic compounds conduct electricity in a molten state or in an aqueous solution.

Wbbse Class 10 Physical Science Solutions

Question 18. Discuss with example the solubility of ionic compounds.

Answer: Ionic compounds are generally soluble in polar solvents (such as water), but are insoluble in nonpolar solvents (such as CH₂, CCI₄, benzene, etc). For example, NaCI is soluble in water but insoluble in benzene.

Question 19. Why does the reactions involving ionic compounds occur at a faster rate?

Answer: Reactions of ionic compounds generally take place in solutions. In solution, the ionic compounds dissociate to form ions that take part in the reaction. Due to the participation of the ions, the reactions occur at a faster rate.

Question 20. Explain why ionic compounds are soluble in polar solvents.

Answer: When an ionic compound is dissolved in a polar solvent, the negatively charged end of the polar solvent attracts the cations present in the crystal lattice of the solid. Similarly the positive end of the solvent attracts the anions.

This attraction by the solvent molecules decreases the force of attraction between cations and anions in the crystal lattice. Hence the dissociated ions are separated and are surrounded by solvent molecules in solution to get stabilised. This is why ionic compounds are soluble in polar solvents.

Question 21. Why are ionic compounds insoluble in non-polar solvents?

Answer: The molecules of a non-polar solvent cannot attract the ions of an ionic compound effectively. Hence, ionic compounds do not generally dissociate in non-polar solvent. That is why ionic compounds are generally insoluble in non-polar solvents.

Question 22. A compound has high boiling and melting point. It does not conduct electricity in solid state but is a good conductor in solution. Which type of bond is present in the compound? What is the nature of the constituent particles of the compound?

Answer:
1. The compound is an ionic or electrovalent compound. So, ionic bond is present in it.
2. The compound is composed of cations and anions.

Question 23. Describe the formation of NaCI.

Answer:

The formation of NaCl

Electronic configuration of Na-atom is K = 2, L = 8, M = 1, and that of chlorine is K= 2, L = 8, M = 7. When sodium reacts with chlorine, 1 electron from the outermost shell of a sodium atom is transferred to a chlorine atom. As a result, Na⁺ and Cl^–  ions are formed. These ions combine by the electrostatic force of attraction to from NaCl.

Question 24. Why the bond of sodium chloride can not be expressed as Na — Cl?

Answer: Na-atom loses its valence electron to form Na⁺ cation and  Cl^– atom accepts it to form Cl^– anion. Na⁺and Cl^–  ions are then held together by strong electrostatic forces of attraction to form the electrovalent compound NaCI. Here no electron pair is formed and this force of attraction has no definite direction. So it can be said that the bond of sodium chloride can not be expressed as Na—Cl.

Question 25. Describe the nature of bond present in CaO.

Answer:

The nature of bond present in CaO:

Electronic configuration of Ca-atom is K= 2, L = 8, M = 8, N = 2 and that of oxygen is K = 2, L = 6 . When calcium reacts with oxygen, Ca-atom releases 2 electrons from its outermost shell to form Ca²⁺ ion and acquire the stable electronic configuration of the nearest noble gas Ar [K = 2, L = 8, M = 8).

Similarly, oxygen atom accepts those electrons to form O^2- ion and attains the electronic configuration of neon [K = 2, L= 8). Ca²⁺ ions and O^2- ions then combine by electrostatic force of attraction to form crystals of CaO. Hence the bond in CaO is ionic.

Wbbse Class 10 Physical Science Solutions

Question 26. Describe the formation of MgCI₂.

Answer:

The formation of MgCI₂

Electronic configuration of Mg-atom is K = 2, L = 8, M = 2, and that of chlorine is K = 2, L = 8, M = 1. When magnesium reacts with chlorine, an Mg-atom releases 2 electrons from its outermost shell to form an Mg²⁺ ion and acquire the electronic configuration of the nearest noble gas Ne (K = 2, L = 8).

On the other hand, two Cl-atoms accept one electron each to form two Cl^–  ions, and each Cl^– ion attains the electronic configuration of Ar [K= 2, L = 8, M = 8). Mg²⁺ ions and two Cl^– ions combine by the electrostatic force of attraction to form MgCI₂. In the MgCI₂ crystal, Mg²⁺ and Cl^– ions are present in 1: 2 ratio.

Question 27. Describe how NaF is formed?

Answer: Electronic configuration of the Na-atom is K-2, L – 8, M = 1, and that of the F-atom is K = 2, L = 7. When Na reacts with F, a Na-atom releases 1 electron from its outermost shell to form a Na⁺ ion and acquire the electronic configuration of the nearest noble gas Ne {K = 2, L = 8). An F -atom accepts this electron to form F“ ion and attains the electronic configuration of Ne [K = 2, L = 8). These ions combine by the electrostatic force of attraction to form NaF.

Question 28. Describe the nature of bond present in LiH.

Answer:

The nature of bond present in LiH:

Electronic configuration of Li-atom is K= 2, L = 1, and that of hydrogen is K = 1, When Li reacts with H, 1 electron is transferred from the outermost shell of Li to the outermost shell of hydrogen. Due to loss of an electron, Li is converted to Li⁺ ion while hydrogen becomes H^– ion by accepting the electron.

Both of them attain the electronic configuration of their nearest noble gas helium (K = 2). Oppositely charged Li⁺ and H^– ions combine through electrostatic force of attraction to form ionic bond.

Question 29. Three elements X, Y, and 2 have atomic numbers 6, 9, and 12 respectively.
1. Which elements among them will form an anion and cation?
2. Which type of bond will be formed between Y and Z when they chemically combine with each other?

Answer:

Element	X	Y	z
Atomic number	6	9	12
Electronic configuration	2,4	2,7	2,8,2

1. Y has only one electron short of the stable electronic configuration of its nearest noble gas. Hence, it will easily form an anion by accepting an electron. Z has two electrons in its outermost orbit. So, it can easily release these two electrons to form a Z²⁺ cation.

2. Y is an electronegative element and Z is an electropositive element. Hence, the bond formed between Y and Z will be ionic in nature. The electrovalency of Z is 2 and that of Y is 1. Hence the formula of the compound will be  ZY₂.

Wbbse Class 10 Physical Science Solutions

Question 30. Complete the following table:

Element	Na	P	C
Formula of chloride
Nature of bond
The physical state of chloride

Answer:

Element	Na	P	C
Formula    of chloride	NaCI	PCI3,PCI5	CCI4
Nature of bond	Ionic	Covalent, covalent	Covalent
The physical state of chloride	Solid	Liquid, liquid	Liquid

Question 31. Electronic configuration of two elements M and N are 2, 8, 1, and 2,7 respectively. By which type of valency will the atoms combine to form compound? What will be the probable formula of the compound?

Answer:
1. It is obvious from the electronic configurations that element M has a tendency to lose an electron from its valence shell to attain the stable electronic configuration and form M⁺ cation. On the other hand element, N has the tendency of accepting the electron in its valence shell to attain the stable electronic configuration and form N^– anion.

Therefore M⁺ and N^– ions will held together by strong electrostatic force of attraction and form compound by electrovalency.

2. The probable formula of the compound will be MN.

Question 32. Electronic configuration of two elements A and B are K(2) 1(7) and K(2) L(8) M(2) respectively. By which type of valency they will combine to form the compound? What will be the formula of the compound.

Answer:
1. Element B will lose 2 electrons to attain the stable electronic configuration and form a B cation while two atoms of elements A will accept one electron each for attaining the stable electronic configuration and form two A^– anions.

These ions will be hold together by strong electrostatic force of attraction to form the compound. Therefore by electrovalency or ionic valency, they will combine to form compound.

2. Formula of the compound will be BA₂.

Question 33. Which one between NaCI and CHCL₃ will form a white precipitate by reacting with AgNO₃?

Answer: NaCI is an ionic compound. Hence it dissociates in aqueous solution to form Na⁺ and Cl^– ions. Due to the presence of free chloride ions in aqueous solution, it reacts with AgNO₃ to produce white precipitate of AgCl.

But CHCL₃ does not produce chloride ions in aqueous solution. Hence, it will not produce any precipitate when reacted with AgNO₃.

Question 34. An element X combines with oxygen to form a compound X₂O. If the molten oxide conducts electricity, then
1. In which group of the periodic table does the element X belong to?
2. what will be the nature and formula of the chloride if X combines with chlorine?

Answer: The molten oxide conducts electricity. Hence, the compound X₂O is an ionic compound. As O is electronegative, X must be an electropositive element.

1. X should be present in . group IA of Mendeleev’s periodic table.
2. Chlorine is electronegative and X is an electropositive element. Hence, ionic bond will be formed between them. The formula of the compound will be XCI.

Question 35. Sodium is a highly reactive metal which reacts with water violently. Chlorine is also a highly reactive non-metal having strong oxidising property. But the compound NaCI is safely used as common salt. Explain with reason.

Answer: Ions are generally more stable than free atoms. During chemical combination, Na-atom loses one electron to form Na⁺ ion while Cl – atom easily accepts that electron to form Cl^– ion. Both Na⁺ and Cl^– ions have stable electronic configurations of their nearest noble gases.

In case of NaCI, a large number of Na⁺ and Cl^– ions aggregate together to form crystals of NaCI. So, free atoms of sodium and chlorine are not present in NaCI and hence the compound is not as much reactive as free sodium and chlorine atoms. Hence, we can safely use NaCI as common salt.

Question 36. When an aqueous solution of NaCI is added to a solution of AgNO₃ the reaction takes place at a very fast rate. Explain why.

Answer: In the reaction between NaCI and AgNO₃, AgCI is precipitated out. It is insoluble in water as well as in acid. Both NaCI and AgNO₃ are ionic compounds. In aqueous solution, both the compounds completely dissociate into ions. These ions react with each other at a very faster rate in the solution. Hence ionic reactions take place rapidly.

\(\begin{gathered}
\mathrm{AgNO}_3(a q) \rightarrow \mathrm{Ag}^{+}(a q)+\mathrm{NO}_3^{-}(a q) \\
\mathrm{NaCl}(a q) \rightarrow \mathrm{Na}^{+}(a q)+\mathrm{Cl}^{-}(a q) \\
\mathrm{NaCl}(a q)+\mathrm{AgNO}_3(a q) \rightarrow \mathrm{AgCl}^2+\mathrm{NaNO}_3(a q)
\end{gathered}\)

Question 37. Atomic number of two elements A and B are respectively 19 and 17. Which type of compound will be formed between A and B? Justify your answer.

Answer: Electronic configuration of A is K =2, L = 8, M = 8, N = 1, and electronic configuration of B is K = 2, Z. = 8, M = 7.

From the electronic configurations, it can easily be understood that A has a strong tendency to release a single electron from its outermost shell to acquire the stable electronic configuration of Ar. Hence A will easily form A⁺ ion.

On the other hand, B has only one electron short from the stable electronic configuration of the nearest noble gas. So, it will show strong tendency to accept an electron to form B^– ion. These ions will bind by electrostatic force of attraction to form an ionic compound AB.

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds Very Short Answer Type Questions Choose The Correct Answer:

Question 1. At normal temperatures, ionic compounds exist as
1. Solids
2. Liquids
3. Gases
4. Crystalline solids

Answer: 4. Crystalline solids

Question 2. The solubility of ionic compounds are higher in
1. Water
2. Alcohol
3. Benzene
4. Ether

Answer: 1. Water

Question 3. In aqueous solution, the crystals of ionic solid dissociates to release
1. Molecules
2. Atoms
3. Cations
4. Both cations and anions

Answer: 4. Both cations and anions

Question 4. Which of the following has the highest boiling point?
1. Sugar
2. Naphthalene
3. Wax
4. NaCI

Answer: 4. NaCI

Question 5. Which of the following, though ionic, is insoluble in water?
1. CuSO₄
2. CH₃COOH
3. NaNO₃
4. CaCO₃

Answer: 4. CaCO₃

Question 6. An example of a polar solvent is
1. Benzene
2. Ether
3. Chloroform
4. Water

Answer: 4. Water

Question 7. Which of the following is not characteristic of an ionic compound?
1. High melting point
2. Electrical conductivity in fused state
3. Consist of oppositely charged ions
4. Constituent ions form crystal lattice fused state

Answer: 4. Constituent ions form crystal lattice fused state

Question 8. The bond formed between the elements of group-1 (except H) and group-17 of the periodic table is
1. Ionic
2. Covalent
3. Coordinate
4. Polar covalent

Answer: 1. Ionic

Question 9. The idea of ionic bond was given by
1. Dalton
2. Lewis
3. Kossel
4. Avogadro

Answer: 3. Kossel

Question 10. The term ‘formula mass’ is applicable for
1. CH₄
2. CaCl₂
3. H₂O
4. CH₃COOH

Answer: 2. CaCl₂

Question 11. In which of the following compounds, the metal ion has an incomplete octet?
1. LiH
2. NaCI
3. CaO
4. MgCI₂

Answer: 1. LiH

Question 12. The electrovalency of Fe in ferrous compounds is
1. 3
2. 1
3. 2
4. 4

Answer: 3. 2

Question 13. An element has 2 electrons in its outermost shell. Which of following statements is correct for the element?
1. The element will form cation
2. The element will form anion
3. The element is capable of forming both cation and anion
4. The element will not form ions

Answer: 1. The element will form cation

Question 14. An element has 6 electrons in its outermost shell. The element will
1. Form cation
2. Form anion
3. Form both cation and anion
4. Not form ions

Answer: 2. Form anion

Question 15. Between Na and Na+, which one has a filled octet?
1. Na
2. Na⁺ 
3. Both Na and Na⁺ 
4. None of these

Answer: 2. Na⁺ 

Question 16. Which one among O, O^– and O^2- most stable?
1. O
2. O^–
3. O^2-
4. All are of equal stability

Answer: 3. O^2-

Question 17. The electronic configuration of Li+ resembles which of the following noble gas?
1. He
2. Ne
3. Ar
4. Kr

Answer: 1. He

Question 18. The electronic theory of valency was proposed by
1. Lewis and Kossel
2. Born and Haber
3. Proust
4. Avogadro

Answer: 1. Lewis and Kossel

Question 19. The nature of bond present in a compound consisting of molecules will be
1. Ionic
2. Electrovalent
3. Covalent
4. Metallic

Answer: 3. Covalent

Question 20. In liquid state, a compound comprised of molecules
1. Is ionic
2. Is electrolytic
3. Conducts electricity
4. Does not conduct electricity

Answer: 4. Does not conduct electricity

Question 21. An ionic compound used in our daily life is
1. Common salt
2. Water
3. Sugar
4. Kerosene

Answer: 1. Common salt

Question 22. The atomic number of elements A, B, and C are respectively (Z- 2), Z, and (Z + 1). If B is a noble gas, the formula of compound formed by A and C will be
1. CA₂
2. CA
3. C₂A₃
4. C₂A

Answer: 4. C₂A

Question 23. Which of the following electrovalent compounds is least soluble in water?
1. NaCI
2. AgCI
3. NaBr
4. Nal

Answer: 2. AgCI

Question 24. The non-directional bond is
1. Covalent bond
2. Coordinate bond
3. Electrovalent bond
4. All of these

Answer: 3. Electrovalent bond

Question 25. Which one of the following does not conduct electricity?
1. Molten NaOH
2. Molten KOH
3. Solid NaCI
4. Molten NaCI

Answer: 3. Solid NaCI

Question 26. Which one is an ionic compound?
1. HCI
2. CH₄
3. MgCI₂
4. NH₃

Answer: 3. MgCI₂

Question 27. If the atomic number of two elements X and Y be 11 and 17 respectively, then the formula of the compound formed with them
1. XY
2. X₂Y
3. XY₂
4. X₂Y₃

Answer: 1. XY

Question 28. In the compound LiH, both/Li’ and ‘H’ atom have attained the electronic configuration of the inert gas as in
1. Ar
2. He
3. Ne
4. Kr

Answer: 2. He

Question 29. The values of electronegativity of A and B are 1.2 and 3.5 respectively. What will be the nature of chemical bond formed between them?
1. Covalent
2. Coordinate
3. Metallic bond
4. Ionic bond

Answer: 3. Metallic bond

Question 30. Which of these does not exist in molecular form?
1. CO₂
2. CH₄
3. NaCl
4. SO₂

Answer: 3. NaCl

Question 31. Solid form of which compound is made up of ions?
1. Sodium chloride
2. Hydrogen chloride
3. Naphthalene
4. Glucose

Answer: 1. Sodium chloride

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds Answer In Brief:

Question 1. What happens when electricity is passed through an aqueous solution of an ionic compound?
Answer: Electrolysis takes place when electricity is passed through an aqueous solution of an ionic compound, resulting in the formation of new substances at the anode and cathode.

Question 2. Mention the nature of force that operates between constituents of an ionic compound and water molecules in an aqueous solution.
Answer: Electrostatic force of attraction operates between constituents of an ionic compound and water molecules in aqueous solution.

Question 3. If an electron is transferred from one atom to another what will be the nature of valency?
Answer: Transfer of an electron from one atom to another result in electrovalency.

Question 4. Between magnesium chloride and naphthalene, which one conducts electricity in a molten state?
Answer: Magnesium chloride, being an ionic compound, conducts electricity in molten state.

Question 5. In MgCI₂, the electronic configuration of the constituent ions resembles which noble gases?
Answer: In MgCI₂, electronic configuration of magnesium and chlorine ions resembles the electronic configuration of neon and argon respectively.

Question 6. What are valence electrons?
Answer: The electrons present in the outermost shell of an atom of an element takes part in bond formation. These electrons are known as valence electrons.

Question 7. Who discovered noble gases?
Answer: Scientists Ralley and Ramsay discovered noble gases.

Question 8. Find the formula mass of sodium chloride.
Answer: The formula mass of sodium chloride is (23 + 35.5) = 58.5

Question 9. The atomic number of an element is 12. What type of bond will it form with chlorine?
Answer: It will form electrovalent bond with chlorine.

Question 10. The atomic number of an element is 9. What type of bond will it form with sodium?
Answer: It will form ionic bond with sodium.

Question 11. Name the atom with which the electronic configuration of hydride ion is identical.
Answer: The electronic configuration of the hydride ion is identical with that of a helium atom.

Question 12. What is the similarity between O^2- ion, and Ne-atom?
Answer: Both O^2- ion and Ne-atom contain 10 electrons and hence they have identical electronic configurations, i.e., they are isoelectronic species.

Question 13. Name an electropositive element that forms both ionic as well as covalent compounds.
Answer: Hydrogen forms both ionic as well as covalent compounds.

Question 14. Write the formula of the compound formed between elements X and Y placed in group 2 and group 17 respectively in periodic table.
Answer: The formula of the compound will be XY₂.

Question 15. What is an electrovalent bond?
Answer: The bond formed between the atoms due to electrovalence is called an electrovalent bond or ionic bond.

Question 16. Name an ionic compound in which the electrovalency of both the anion and cation is 3.
Answer: Aluminium nitride (AIN)

Question 17. A metallic element M forms an ionic oxide MO. How many electrons are present in the valence shell of the M- atom?
Answer: The valency of the metal M is 2. Hence it contains 2 electrons in its valence shell.

Question 18. Give example of an ionic compound where the octet rule is not applicable.
Answer: Lithium hydride (LiH).

Question 19. Give example of a stable cation in which octet is not fulfilled.
Answer: Li⁺.

Question 20. Write down the number of valence electrons of alkali metals and halogens.
Answer: Number of valence electrons in alkali metals and halogens are 1 and 7 respectively.

Question 21. In which type of compound molecule does not exist?
Answer: Ionic compound.

Question 22. How manyelectrons are present in the valence shell of Ca?
Answer: 2 electrons.

Question 23. Which type of chemical bond is present in CaO?
Answer: Ionic bond.

Question 24. Which type of bond is present in CaCl_2.
Answer: Ionic bond.

Question 25. Atomic number of an element is 12. which type of compound will be formed by its reaction with chlorine?
Answer: Ionic compound will be formed.

Question 26. Atomic number of an element is 9. Which type of compound will be formed by its reaction with sodium?
Answer: Ionic compound will be formed.

Question 27. Which type of bond will be formed between A (2, 8, 2) and B (2, 8,7)?
Answer: Ionic bond will be formed between A and B.

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds Fill In The Blanks:

 

Question 1. Ionic compounds are generally ________ and their crystals have a definite _____ shape.
Answer: Solid, geometrical

Question 2. The crystals of ionic compounds are formed by a large number of ______ and ________
Answer: Cations, anions

Question 3. Ionic compounds do not conduct electricity in ______ state.
Answer: Solid

Question 4. Ionic compounds conduct electricity in ________ state or in _______ solution.
Answer: Fused, aqueous

Question 5. Ionic compounds are soluble in _______ Solvents such as ________
Answer: Polar, water

Question 6. Ionic compounds are generally __________ in organic solvents.
Answer: Insoluble

Question 7. An element which forms a cation to achieve the electronic configuration of neon is _______
Answer: Sodium

Question 8. In ______ bonds, the particles are held together by electrostatic force of attraction.
Answer: Ionic

Question 9. An atom gains _______ by either accepting or losing electron(s).
Answer: Stability

Question 10. An ionic bond is formed between a _______ and a _______element.
Answer: Metallic,non-metallic

Question 11. Two elements form ionic bond with each other when the difference in electronegativity between the elements is _______
Answer: High

Question 12. An ionic compound is formed when a stable three-dimensional _______ is formed.
Answer: Crystal

Question 13. Electrovalency of Cu in cupric salts is _______
Answer: 2

Question 14. Greater the magnitude of the _______ energy of an ionic compound, the greater will be the ________ of its crystal.
Answer: Lattice, stability

Question 15. In ionic compounds, the ionisation potential of the element forming cation should be ________
Answer: Low

Question 16. In ionic compounds, the electronegativity of the elementforming anion should be _______
Answer: High

Question 17. In CaO, the charge of calcium ion is _______
Answer: 2+

Question 18. In LiH, the cation is _______ and the anion is ________
Answer: Li⁺, H^–  

Question 19. _______ bond is not considered as a real bond.
Answer: Ionic

Question 20. In 1894, Scientist _______ discovered existence of Argon gas.
Answer: Lord Raley

Question 21. Except He, all the inert gases contain ________ electrons in their valence shell.
Answer: 8

Question 22. In _______ compound bond means the electrostatic force of attraction.
Answer: Ionic

Question 23. Ionic bonds are formed between _______ and _______ elements.
Answer: Metalic,non-metallic

Ionic and Covalent Bonds Topic A Electronic Concept of Valency, Ionic Bonding, and Properties of Ionic Compounds State Whether True Or False:

Question 1. Li, Be and H obeys the octet rule.
Answer: False

Question 2. Sodium chloride All the inert gas elements are chemically inert.
Answer: False

Question 3. Duplet rule is only applicable to H, Li and Be.
Answer: True

Question 4. In ionic compounds, there is no existence of independent molecules.
Answer: True

Question 5. In electrovalency, atoms complete their octet or duplet by sharing electron pairs.
Answer: False

Question 6. Ionic compounds undergo very fast reactions in solution.
Answer: True

Question 7. Crystal lattice is found in ionic compounds.
Answer: True

Question 8. Molecular mass can be calculated for ionic compounds.
Answer: False

Question 9. BN (boron nitride) is an ionic compound.
Answer: False

Question 10. Ionic bonds are real bonds as they can form crystal lattice.
Answer: False

Question 11. An example of a compound with incomplete octet is BF₃.
Answer: True

Question 12. Ionic compounds are hard as the intermolecular spaces between the constituent particles are very small.
Answer: True

Question 13. Formation of ionic bond is favoured by high electron gain enthalpy of the atom that forms the anion.
Answer: True

Question 14. Ionic compounds are good conductors of electricity in solid state as the crystal lattice of ionic compounds contain free electrons.
Answer: False

Question 15. Ionic bonds are directional and hence they can show isomerism.
Answer: False

Question 16. Distilled water can conduct electricity.
Answer: False

Question 17. Individual molecules of ionic compounds do not exist.
Answer: True

Question 18. Isomerism isnotobserved in ioniccompounds.
Answer: True

Question 19. Na and Cl atoms are combined to each other by ionic valency.

Answer: True

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds Synopsis:

 

Covalency: In order to achieve the stable electronic configuration of the nearest noble gas, equal number of electron(s) from the outermost shells of two atoms form one or more electron pairs which are evenly shared between the two atoms. The capacity of the elements for this type of chemical combination is called covalency. The bond formed between the atoms due to covalency is called a covalent bond and the compound formed in this way is called a covalent compound.

Types of covalent bonds: Covalent bonds of three types. The atoms involved in the formation of covalent bonds may share 1, 2, or 3 electron pairs equally with each other to form a single bond, double bond, and triple bond respectively.

Characteristics of covalent compounds: Covalent compounds exist as discrete molecules. These compounds may be solid, liquid or gas and have low melting and boiling points. They usually do not conduct electricity. Covalent compounds are generally soluble in non-polar solvents and the rate of the reactions involving covalent molecules are very slow.

Lewis-dot diagram is a simple representation of a covalent molecule. Here, each electron is represented by a dot, and each bond is represented by a pair of dots in between the two atoms involved in bond formation. In the Lewis-dot diagram only the bonded electron pairs are shown.

H₂, N₂, O₂, Cl₂ etc., are examples of covalent element while NH₃, CO₂, CH₄, HF etc, are examples of covalent compounds. Single bond is present in the molecules of H₂, F₂, HF, HCI, H₂O, NH₃, CH_4, etc. Molecules like CO₂, C₂H₄, O₂, etc., contain double bond while triple bond is present in areN₂, C₂, H_2, etc.

The Lewis-dot diagram of O₂ molecule does not agree with the experimental observations regarding the structure of O₂. Hence, O₂ is not represented by Lewis-dot structure. In fact no such Lewis-dot structure of O₂ is possible which can describe the experimental observations.

 

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds Short And Long Answer Type Questions:

Question 1. What is covalency?

Answer:

Covalency:

In order to achieve the electronic configuration of the nearest noble gas, equal number of electron(s) from the outermost shells of two atoms form one or more electron pairs which are evenly shared by the two atoms. The capacity of the elements for this type of chemical combination is called covalency.

Question 2. What are covalent bonds and covalent compounds?

Answer: The bond formed between the atoms by the mutual sharing of electron pairs is called covalent bond and the compound containing covalent bonds is called a covalent compound.

Question 3. Classify covalent molecules.

Answer:

Covalent molecules are of two types. These are:

1. Covalent elementary molecules: The atoms of the same element are covalently bonded with each other to form covalent elementary molecules. H₂, N₂, O₂ F_2, etc., are examples of this type.
2. Covalent compound molecule: The atoms of different elements are covalently bonded with each other to form covalent compound molecules. NH₃, CO₂, HCl, and CH₄ are examples of this type.

Question 4. How is a covalency covalent molecule?

Answer: During the formation of a covalent molecule, the number of electron pairs formed by an atom of an element is its covalency. For example, in methane (CH₄) C-atom forms four electron pairs with 4 hydrogen atoms. Hence, covalency of carbon in methane is 4.

Question 5. Glucose is a solid but methane is a gas though both are covalent compounds. Explain.

Answer:

Glucose is a solid but methane is a gas though both are covalent compounds

The force of attraction between the molecules of a covalent compound is very weak and the attractive force increases with an increase in molecular mass.

The molecular mass of glucose is much higher than that of methane and hence the force of attraction between glucose molecules is stronger than that between methane molecules. Hence, glucose is solid at room temperature but methane Is a gas.

Question 6. Explain why naphthalene is insoluble in water but readily dissolves in benzene.

Answer: Covalent molecules are generally insoluble in polar solvents and soluble in non-polar solvents. Water is a polar solvent while benzene is a nonpolar solvent. Thus, naphthalene being a covalent compound dissolves in benzene but is insoluble in water.

Question 7. An aqueous solution of sodium chloride conducts electricity, but aqueous solution of sugar or glucose does not conduct electricity. Explain why.

Answer:

An aqueous solution of sodium chloride conducts electricity, but aqueous solution of sugar or glucose does not conduct electricity.

Sodium chloride (NaCI) being an ionic compound dissociates in aqueous solution to form Na⁺ and Cl^– ions. These ions are responsible for conducting electricity and hence NaCI is a good conductor of electricity in aqueous solution. On the other hand, sugar or glucose being covalent compound does not dissociate in solution and hence they do not conduct electricity.

Question 8. Though HCI is a covalent compound, its aqueous solution conducts electricity. Why?

Answer:

Though HCI is a covalent compound, its aqueous solution conducts electricity.

The electronegativity of chlorine is greater than that of hydrogen. As a result, the electron pair shared between H and Cl-atom in HCI molecule shifts more towards Cl-atom. Due to this, a partial positive charge develops on H-atom and a partial negative charge develops on Chatom.

Water is also a polar molecule in which H-atom is partially positively charged and O-atom is partially negatively charged. Due to force of attraction between the opposite charges of HCI and water molecules, HCI ionises to form H₃O⁺ and Cl^– ions. As a result, of the formation of these ions in solution, HCI conducts electricity.

Question 9. Why are the melting and boiling point of covalent compounds generally low?

Answer: The intermolecular force acting between the molecules of covalent compounds is very weak. Hence, less energy is required to separate the molecules. So, covalent compounds generally have low melting and boiling point.

Question 10. The reactions involving covalent molecules are slow. Explain with reason.

Answer:

The reactions involving covalent molecules are slow.

In covalent molecules, the atoms are bonded by strong covalent bonds. The cleavage of these bonds requires sufficient energy, time and sometimes catalysts. That is why, reactions involving covalent molecules take place slowly. For example— fermentation of glucose to ethanol takes almost 3 days to complete.

Question 11.  Cl₂ molecules are formed but Ne₂ molecules do not exist. Explain why.

Answer:

Cl₂ molecules are formed but Ne₂ molecules do not exist.

Cl-atom has 7 electrons in its outermost shell. Two chlorine atoms contribute one electron each to form a commonly shared pair and in this way, both of them complete their octets.

The sharing of the electron pair results in the formation of Cl₂ molecule. On the other hand, Ne-atoms already have filled octet and hence do not show any tendency to combine with each other to form Ne₂ molecule.

Question 12. What is polar covalent bond?

Answer:

Polar covalent bond?:

When a covalent bond is formed between two atoms of different elements, due to the difference in electronegativity between the atoms the shared electron pair is attracted more towards the atom with greater electronegativity.

As a result the more electronegative atom becomes partially negatively charged. On the other hand, the less electronegative element develops a partial positive charge. This charge difference between the atoms causes polarity in the covalent bond. This type of bond is called polar covalent bond.

Question 13. Discuss about the electrical conductivity of covalent compound.

Answer:

The electrical conductivity of covalent compound:

Covalent compounds do not possess cations or anions and so cannot usually conduct electricity in the fused or dissolved state. However, some covalent compounds when dissolved in a polar solvent form ions by reacting with the solvent molecules and conducting electricity.

For example, hydrogen chloride dissolves in water and forms hydronium ions and chloride ions by reacting with water molecule. As a consequence aqueous solution of hydrogen chloride can conduct electricity.

\(\mathrm{HCl}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{Cl}^{-}\)

Question 14. Mention which of the following compounds conduct electricity and which can not—Cl₂, Al₂O₃, an aqueous solution of HO, NaBr.

Answer: Al₂O₃, aqueous solution of HCI and NaBr can conduct electricity and Cl₂ can not conduct electricity.

Question 15. Classify the following compound as ionic and covalent compound—CCI₄, Na₂SO₄, CaCI₂, CaO, CHCI₃, C₂H₂, NaH, LiH, MgO, NaBry NaCI, Cl₂, O₂, H₂O, KCI, N₂, CO₂, Na₂O, NH₃, Mg(NO₃)₂, K₂S, C₂H₆, HCI, NaF, MgCI₂, CH₄.

Answer:
Ionic compound: Na₂SO₄, CaCl₂, CaO, NaH, LiH, MgO, NaBr, NaCI, KCI, Na₂O, Mg(NO₃)₂, K₂S, NaF, MgCI₂.
Covalent compound: CCI₄, CHCI₃, C₂H₂, Cl₂, O₂, H₂O, N₂, CO₂, NH₃, C₂H₆, HCl, CH₄.

Question 16. Compare naphthalene and sodium chloride by the following properties—boiling point and solubility in water.

Answer:

Comparison between naphthalene and sodium chloride are:

Property	Naphthalene	Sodium chloride
Boiling point	Since naphthalene is a covalent compound its boiling point is comparatively low(80.2°C).	Being an ionic compound sodium chloride has a comparatively higher boiling point (801°C).
Solubility in water	Naphthalene is insoluble in water.	Sodium chloride is soluble in water.

Question 17. Differentiate between ionic and covalent compounds on the basis of the following properties:
1. Hardness and brittleness and
2. Melting and boiling point.

Answer:

The differences are listed below:

Property	Ionic compound	Covalent compound
1. Hardness and brittleness	The crystals of ionic compounds are hard and brittle in nature.	Most of the solid covalent compounds are very soft. However, some covalent compounds are able to form giant crystal lattices having a high degree of hardness.
2. Melting    and boiling point	Ionic compounds generally have high melting and boiling point.	Except for some solid covalent compounds ( SiO2 etc.), the melting and boiling point of most covalent compounds are low.

Question 18. Differentiate between ionic and covalent compounds on the basis of the following properties:
1. Electric conductance and
2. Solubility.

Answer:

The differences are listed below:

Property	Ionic compound	Covalent compound
1. Electrical conductance	Ionic compounds are non-conductors of electricity in solid state but in aqueous solution or in molten state they conduct electricity.	Pure covalent compounds are usually non-conductors of electricity. (Exception— aqueous solution of HCI conducts electricity).
2. Solubility	Ionic compounds are generally soluble in polar solvents but insoluble in non-polar solvents.	Covalent compounds are generally soluble in non-polar solvents but insoluble in polar solvents. However, covalent compounds such as HCl, sucrose, glucose, ethanol are soluble in water.

Question 19. Differentiate between ionic and covalent compounds on the basis of the following properties:
1. Nature of bond and
2. Electrolytic nature.

Answer:

The differences are listed below:

Property	Ionic compound	Covalent compound
1. Nature of bond	Ionic compounds are formed by the electrostatic force of attraction acting between two oppositely charged ions. There is no existence of discrete molecules.	Covalent compounds are formed by sharing of equal number of electron pairs between the two atoms. Covalent compounds form discrete molecules.
2. Electrolytic nature	Almost all ionic compounds are electrolytes in a molten state or when dissolved in a suitable solvent.	Most covalent compounds are nonelectrolytes. However, some polar covalent compounds (e.g, HCI) ionise in polar solvents and hence act as electrolytes.

Question 20. State three important differences between ionic and covalent bond.

Answer:

The three major differences between ionic and covalent bonds are discussed below:

Characteristic	Ionic bond	Covalent bond
1. Process of bond formation	An ionic bond is formed by transfer of electron(s) from the atom of an electropositive element to the atom of an electronegative element.	The atoms involved in covalent bond formation share equal number of electrons to form electron pair(s) and both of them use these electron pair(s) equally to form a covalent bond.
2. Difference in electronegativity	The difference in electronegativity of the participating atoms is very high.	The participating atoms have comparable electronegativity.
3. Nature of bond	It is actually an electrostatic force of attraction.	The force of attraction arises due to sharing of electron pairs is known as a covalent bond.

Question 21. Explain with an example how the covalent bond is formed according to Lewis’s concept.

Answer:

G.N. Lewis (1916) proposed an electronic model, according to which, when both the atoms taking part in a chemical combination are short of electrons than the stable electronic configuration of the nearest noble gas, they can share their electrons in order to complete their octets (duplet in the case of H) to form a covalent bond.

During formation of a covalent bond, the two combining atoms contribute equal number of electrons for sharing. The shared electrons become a common property of both atoms and are responsible for holding the atoms together. Since such a combination of atoms does not involve the transfer of electrons form one atom to another, the atoms are electrically neutral.

Example: Of the two F-atoms (electronic configuration 2, 7), one electron from each F atoms are shared to form an electron pair. This electron pair is then used by both atom to attain the stable configuration like Ne-atom. (electronic configuration: 2, 8). As a consequence, F₂ molecule is formed.

Question 22. Describe the formation of hydrogen molecule using Lewis-dot structure.

Answer:

Hydrogen molecule is diatomic. Two hydrogen atoms combine with each other by single bond to form  H₂ molecule. Hydrogen atom contains a single electron in its K orbit. During the formation of  H₂ molecule, two hydrogen atoms share 1 electron each to form an electron pair, and this electron pair equally.

As a result, both the atoms attain the electronic configuration of the nearest noble gas He. Due to sharing of an electron pair, a single bond is formed between two H -atoms to form an H₂ molecule.

Question 23. Justify the Lewis-dot structure of oxygen.

Answer:

The Lewis-dot structure of oxygen

Two O-atoms combine with each other by forming covalent double bond (O=O) in which both the atoms attain the stable electronic configuration of Ne. But the Lewis-dot structure of oxygen molecule fails, to explain all the experimental observations regarding the structure of O₂ molecule.

Actually, no such Lewisdot structure of O₂ molecule can describe all the properties of oxygen.

 

Question 24. Describe the formation of F₂ molecule with the help of Lewis-dot diagram. Or, Describe the formation of homonuclear diatomic molecule.

Answer: Two F-atoms combine chemically by forming a single covalent bond to form an F₂ molecule. An F -atom contains 2 electrons in its K orbit and in its L orbit. During the formation of F₂ molecule, each F-atom shares one electron from its valence shell equally with the other to form an electron pair. As a result, both the atoms attain the stable electronic configuration of the nearest noble gas Ne. This results in the formation of stable F₂ molecule.

 

Question 25. Explain how an HF molecule is formed with the help of Lewis-dot diagram.

Answer: HF molecule is formed by the chemical combination of an H-atom and an F-atom. H-atom has 1 electron in its valence shell. On the other hand, F-atom has 7 electrons in its valence shell.

During chemical combination, an H-atom and an F-atom share one electron each from their valence shell to form an electron pair which is equally shared between them. As a result, both the atoms attain the stable electronic configuration of their nearest noble gases. This results in the formation of HF molecule.

Question 26. Explain how an HCI molecule is formed with the help of Lewis-dot diagram. Or, Describe the formation of a heteronuclear diatomic molecule.

Answer: An HCI molecule is formed by the chemical combination of an H-atom and a Cl-atom. H-atom has 1 electron in its valence shell. On the other hand, Cl-atom has 7 electrons in its valence shell.

During chemical combination, an H-atom and a Cl-atom share one electron each from their valence shell to form an electron pair which is equally shared between them. As a result, both the atoms attain the stable electronic configuration of their nearest noble gases. This results in the formation of HCI molecule.

Question 27. Describe the formation of molecule with the help of Lewis-dot diagram.

Answer:

The formation of molecule with the help of Lewis-dot diagram:

In H₂O molecule, an O-atom is covalently bonded to two H-atoms. H-atom has 1 electron in its valence shell. On the other hand, O-atom has 6 electrons in its valence shell. During chemical combination, an oxygen atom shares its two valence electrons to separately pair with each electron of the two H-atoms.

As a result, O-atom attains the stable electronic configuration of Ne and both the hydrogen atoms attain the electronic configuration of He. Thus, two single bonds are formed between O-atom and two H-atoms. The other two pairs of electrons in the valence shell of oxygen atom do not take part in bond formation and remain as lone pairs of electrons.

Question 28. Describe the formation of NH₃ molecule with the heip of Lewis-tlot diagram.

Answer:

The formation of NH₃ molecule with the heip of Lewis-tlot diagra

In NH₃ molecule, N-atom is covalently bonded with three H-atoms. H-atom has 1 electron in its valence shell. On the other hand, N-atom has 5 electrons in its valence shell. During chemical combination, a nitrogen atom shares its three valence electrons to separately pair with each electron of the three H-atoms.

As a result, N-atom attains the stable electronic configuration of Ne and all the three hydrogen atoms attain stable electronic configuration of He. Thus, three single bonds are formed between N-atom and three H-atoms. One pair of electron in the valence shell of nitrogen does not take part in bond formation and remains as lone pair of electron.

Question 29. Describe the formation of CH₄ molecule with the help of Lewis-dot diagram.

Answer:

The formation of CH₄ molecule with the help of Lewis-dot diagram:

In a CH₄ molecule, a C-atom (2,4) is covalently bonded with four H-atoms. H-atom has 1 electron in its valence shell. On the other hand, C-atom has 4 electrons in its valence shell. During the chemical combination, a carbon atom shares its four valence electrons to separately pair with each electron of the four H-atoms.

As a result, C- atom attains the stable electronic configuration of Ne and all the four hydrogen atoms attain the electronic configuration of He and become stable. Thus, four single bonds are formed between C- atom and four H-atoms. This results in the formation of CH₄ molecule.

Question 30. Describe how a CO₂ molecule is formed with the help of Lewis-dot diagram.

Wbbse Class 10 Properties Of Covalent Compounds

Answer: During the formation of CO₂ molecule, two electrons out of the 4 valence electrons of C-atom are shared with 2 valence electrons of an oxygen atom.

As a result, a double bond is formed between C-atom and its O-atom. C-atoms uses its remaining two valence electrons to form a double bond with another O-atom similarly. Thus, in a CO₂ molecule the C-atom forms two double bonds with two oxygen atoms.

Question 31. Describe the formation of a C₂H₄ molecule with the help of Lewis-dot diagram.

C2h4 Covalent Bond Diagram

Answer: Ethylene molecule is formed by the combination of two C-atoms with four H-atoms. Each of the two C-atoms share two valence electron pairs with each other to form a double bond (C — C). Each of the C-atoms uses its two remaining valence electrons to form two single bonds with two hydrogen atoms. Thus, the molecule contains four C—H bonds apartfrom a C = C bond.

Question 32. Describe the formation of a N₂ molecule with the help of Lewis-dot diagram.

Answer:

The formation of a N₂ molecule with the help of Lewis-dot diagram

During the formation of N₂ molecule by the combination of two N-atoms (2,5), each N-atom shares its three valence electrons with three valence electrons of the other N-atom. Thus, three electron pairs are formed, which are equally shared by both N -atoms. Due to sharing of three electron pairs, a triple bond is formed between two N-atoms, and N₂ molecule is formed.

C2h4 Covalent Bond Diagram

Question 33. Describe how a C₂H₂ molecule is formed with the help of Lewis-dot diagram.

Answer: An acetylene molecule is formed by the combination of two C-atoms with two H-atoms. Each of the two C-atoms share three electron pairs with each other to form a triple covalent bond (C = C). Each of the C-atoms uses its remaining valence electron to form single bond with a hydrogen atom, Thus, C₂H₂ molecule contains two C—H bonds apart from a C=-C bond.

Question 34. Draw the Lewis dot structure of CCI₄.

Answer:

 

Question 35. Through what sort of chemical bonding does chloride combines with sodium and hydrogen to form sodium chloride and hydrogen chloride- respectively? Draw corresponding electron dot structure in both cases.

Wbbse Class 10 Properties Of Covalent Compounds

Answer: Chlorine combines with sodium through ionic bonding or electrovalent bonding and with hydrogen through covalent bonding. Electron dot structure of sodium chloride Electron dot structure of hydrogen chloride:

Electron dot structure of hydrogen chloride:

Question 36. Atomic number of the elements X, Y, and Z are respectively 9, 10, and 11.
1. Which one of them will form ionic bond only and which one will form both ionic and covalent bonds? Answer bn the basis of electronic configuration only.
2. Draw the Lewis-dot diagram of the compounds formed by the combination of these elements.

Answer:
1. Electronic configuration of X is K = 2, L = 7 . The electronic configuration of Y is AT = 2, L = 8. The electronic configuration of Z is K= 2, L = 8, and M = 1.

C2h4 Covalent Bond Diagram

X has 7 electrons in its outermost orbit. So, it will form both ionic and covalent bonds. The complete octet of Vindicates that it is chemically inert. Z has only 1 electron in its outermost shell. Hence, it can form ionic bonds only.

2. Atoms of X and Z will combine with each other to form a compound ZX.

Two atoms of element X combine with each other to form X₂ molecule.

Question 37. Number of valence electrons in the atoms of the elements A, B, C, and D are respectively 1, 4, 6, and 7. Form two ionic and two covalent compounds with these elements and draw their structures.

Answer:

Element	A	B	C	D
Number of valence electrons	1	4	6	7

Ionic compounds formed from these elements are:

Covalent compounds formed from the these elements are:

Question 38. Atomic numbers of elements A, B and C are 3,11, and 19 respectively. Which type of compound, ionic or covalent, will be formed by the reaction of B with chlorine?

Answer:

Element	Atomic No.	Electronic configuration
A	3	2	1
B	11	2	8	1
C	19	2	8	8	1

The atom of element B contains 1 electron in its valence shell which may be donated to attain the stable electronic configuration. On the other hand Cl (electronic configuration: 2, 8, 7) atom has the tendency to accept electron to attain the stable electronic configuration.

Hence atom of B will lose 1 electron to form the B^⊕ ion and Cl^Θ ion. These two ions will hold together, to form BCI, by a strong electrostatic force of attraction. Hence ionic compound will be formed by the reaction of B with chlorine.

Question 39. Atomic number of the element X is 7. If two atom of X combines to form molecule, then which type of bond-ionic or covalent will be formed? Draw the electron dot structure of the molecule.

Answer: Electronic configuration of X is 2, 5. Each of the two combining atoms of X will contribute 3 electrons form their valence shell to form 3 common shared pair of electron and in this way both of them complete their octets. Hence covalent triple bond will be formed during the formation of X₂ molecule.

Question 40. Assume symbols of two elements are A and B and their atomic numbers are 7 and 20 respectively. If two atom of A combine to form a molecule then which type of bonding occurs during bond formation. Draw the electron dot structure of the molecule.

Answer:
1. During the formation of molecule of A by the combination of two atoms of it will combine through covalent bonding.
2. Electronic configuration of A is 2, 5.

Now electron dot structure of the molecule A_2.

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds Very Short Answer Type Questions Choose The Correct Answer:

Question 1. With respect to covalent compounds, the melting points of ionic compounds are generally
1. Lower
2. Higher
3. Equal
4. Lower or higher

Answer: 2. Higher

Question 2. Which of the following is insoluble in water?
1. MgCI₂
2. Na₂SO₄
3. CHCI₃
4. HCL

Answer: 3. CHCI₃

Wbbse Class 10 Properties Of Covalent Compounds

Question 3. A covalent compound used in our daily life is
1. Washing soda
2. Alum
3. Water
4. Chalk

Answer: 3. Water

Question 4. Both ionic and covalent bonds can not be formed by
1. H
2. O
3. Na
4. Cl

Answer: 3. Na

Question 5. Which of the following compounds form both ionic and covalent bonds?
1. NH3
2. CCI₄
3. CO₂
4. CaCO₃

Answer: 4. CaCO₃

Physical Science Class 10 West Bengal Board

Question 6. Which of the following, though covalent, but possesses high melting point?
1. C₂H₆
2. SiO₂
3. C₄H₁₀
4. CO₂

Answer: 2. SiO₂

Question 7. The concept of covalent bonding was given by
1. Lewis
2. Arrhenius
3. Kelvin
4. Avogadro

Answer: 1. Lewis

Question 8. The Lewis structure does not justify the experimental result for
1. N₂
2. NH₃
3. CH₄
4. O₂

Answer: 4. O₂

Question 9. A double bond is present in
1. H₂O
2. NH₃
3. CO₂
4. CH₄

Answer: 3. CO₂

Physical Science Class 10 West Bengal Board

Question 10. A triple bond is present in’
1. H₂
2. O₂
3. N₂
4. F₂

Answer:  3. N₂

Question 11. Elements forming covalent bonds are placed in periodic table at the
1. Extreme left
2. Extreme right
3. Middle of the table
4. The bottom

Answer: 2. Extreme right

Question 12. Maximum number of covalent bonds that can be formed between two atoms is
1. 1
2. 2
3. 3
4. 4

Answer: 3. 3

Question 13. In which of the following compounds, the central atom has an expanded octet?
1. CH₄
2. PCI₅
3. NH₃
4. NCl₃

Answer: 2. PCI₅

Question 14. In which of the following compounds, the central atom has an incomplete octet?
1. HCN
2. BF₃
3. CO₂
4. H₂O

Answer: 2. BF₃

Physical Science Class 10 West Bengal Board

Question 15. Covalent bond is not formed between
1. C&CI
2. B&F
3. Cl & Cl
4. Na & F

Answer: 4. Na & F

Question 16. A triple bond is absent in
1. C₂H₂
2. N₂
3. HCN
4. C₂H₄

Answer: 4. C₂H₄

Question 17. The nature of the bonds in HCI(g) .
1. Covalent
2. Ionic
3. Coordinate
4. Ionic and covalent

Answer: 1. Covalent

Wbbse Class 10 Properties Of Covalent Compounds

Question 18. In which of the following compounds ionic, covalent and coordinate—all the three types of bond are present?
1. KCI
2. H₂O
3. NH₄CI
4. NaCN

Answer: 3. NH₄CI

Question 19. Which of the following combines with hydrogen to form covalent compounds?
1. Ca
2. Na
3. C
4. Li

Answer: 3. C

Question 20. Which of the following is a covalent compound?
1. CaO
2. MgO
3. NaCI
4. CH₄

Answer: 4. CH₄

Question 21. Which of the following compound central atom contains lone pair of electron?
1. CH₄
2. NH₃
3. C₂H₂
4. C₂H₄

Answer: 2. NH₃

Physical Science Class 10 West Bengal Board

Question 22. A compound with covalent triple bond is
1. H₂O
2. CO₂
3. C₂H₂
4. CH₄

Answer: 3. C₂H₂

Question 23. Compound containing covalent double bond is
1. OF₂
2. N₂
3. C₂H₄
4. C₂H₂

Answer: 3. C₂H₄

Wbbse Class 10 Properties Of Covalent Compounds

Question 24. Which one of the following does not conduct electricity in molten state or in solution?
1. NaCI
2. NaF
3. CCI₄
4. KCI

Answer: 3. CCI₄

Question 25. True bond can be found in
1. Ionic bond
2. Covalent bond
3. Nuclear bond
4. None of these

Answer: 2. Covalent bond

Question 26. Number of covalent bonds in nitrogen molecule is
1. 1
2. 2
3. 3
4. 4

Answer: 3. 3

Physical Science Class 10 West Bengal Board

Question 27. In which of the following compound covalent bonds are present?
1. Hydrogen chloride
2. Sodium chloride
3. Lithium hydroxide
4. Calcium oxide

Answer: 1. Hydrogen chloride

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds Answer In Brief:

Question 1. Name a non-polar solvent.
Answer: Benzene is a non-polar solvent.

Question 2. Between water and benzene, which one dissolves kerosene?
Answer: Kerosene, a covalent organic substance, dissolves in a non-polar solvent like benzene.

Question 3. Name a crystalline substance that contains covalent bonds.
Answer: Diamond is a crystalline solid in which C-C covalent bonds are present.

Question 4. Name a covalent solid whose hardness is more than any ionic compound.
Answer: Diamond, a covalent compound, is the hardest substance. Its hardness is more than any ionic compound.

Question 5. How many lone pairs of electrons are present in nitrogen molecules?
Answer: Nitrogen molecule contains two lone pairs of electrons (N = N).

Wbbse Class 10 Properties Of Covalent Compounds

Question 6. Define crystal lattice.
Answer: A crystal lattice is defined as a regular three-dimensional network of constituent particles (atoms, molecules, or ions) of the crystal.

Physical Science Class 10 West Bengal Board

Question 7. The atomic number of an element is 6. What type of bond will it form with the element having atomic number 17?
Answer: Both elements are non-metals. So, they will form covalent bonds.

Physical Science Class 10 West Bengal Board

Question 8. Name an electronegative element that forms both ionic and covalent compounds.
Answer: Chlorine forms both ionic and covalent compounds.

Question 9. What happens when crystals of sugar are heated?
Answer: When crystals of sugar are heated, it releases molecules of water and black particles of carbon are left behind as residue.

Question 10. Name a non-volatile covalent compound.
Answer: Sugar or cane sugar (C₁₂H₂₂O₁₁) is a nonvolatile covalent compound.

Question 11. Give example of a solid and a liquid covalent compound.
Answer: A solid covalent compound is a glucose (C₆H₁₂O₆) and a liquid covalent compound is water (H₂O).

Question 12. Between water and ethanol, which one is more volatile?
Answer: Ethanol is more volatile than water.

Question 13. Is it possible to differentiate between solid naphthalene and solid common salt on the basis of their electrical conductivity?
Answer: In solid state, both naphthalene and common salt are bad conductors of electricity. Hence, it is not possible to differentiate between them on the basis of their electrical conductivity.

Physical Science Class 10 West Bengal Board

Question 14. How many covalent bonds can be formed between two atoms?
Answer: Two atoms can form a maximum of 3 covalent bonds between them.

Question 15. Show the structure of  C₂H₂ by dash formula.
Answer: H—C=C—H.

Physical Science Class 10 West Bengal Board

Question 16. Give an example of a single bonded covalent molecule which is liquid at normal temperature.
Answer: Water is a single bonded covalent molecule which is liquid at normal temperature.

Question 17. Name a gaseous hydrocarbon that contains only single bonds.
Answer: Methane (CH₄) is a gaseous hydrocarbon that contains only single bonds.

Wbbse Class 10 Properties Of Covalent Compounds

Question 18. Name a gaseous hydrocarbon that contains double bond.
Answer: Ethylene (C₂H₄) is a gaseous hydrocarbon that contains a double bond.

Question 19. Give an example of a molecule which contains an atom having a lone pair of electrons.
Answer: In ammonia (NH₃) molecule, the nitrogen atom contains a lone pair of electrons.

Question 20. The electronic configuration of N-atom is 2(K), 5(L). How many electrons in the valence shell of nitrogen do not take part in the formation of N₂ molecule?
Answer: 2 electrons from each N-atom do not take part in the formation of N₂ molecule.

Question 21. Name a gaseous covalent compound which ionises in aqueous solution and conducts electricity.
Answer: Hydrogen chloride (HCI) is a gaseous covalent compound which ionises in an aqueous solution and conducts electricity.

Question 22. What is Lewis-dot diagram?
Answer: Lewis-dot diagram is a simple representation of a covalent molecule. Here, each electron is represented by a dot, and each bond is represented by a pair of dots in between two atoms involved in bond formation.

Physical Science Class 10 West Bengal Board

Question 23. What is the linear formula of covalent compounds?
Answer: Linear formula of a covalent compound is a simple representation in which each covalent bond is represented by a line or dash between the atoms involved in bond formation.

Question 24. How many types of covalent bonding are there?
Answer: There are three types of covalent bond single bond, double bond, and triple bond.

Question 25. Isomerism can be found in which type of compound?
Answer: Covalent compound.

Wbbse Class 10 Properties Of Covalent Compounds

Question 26. Give example of a solid covalent which is good conductor of electricity.
Answer: Graphite.

Question 27. Give example of a single bonded covalent compound which is liquid at room temperature.
Answer: Water.

Question 28. Give example of two covalent compounds whose boiling and melting points are very high.
Answer: Boron carbide ( B₄C) and Silica ( SiO₂).

Question 29. Which one between ionic or covalent compounds are generally electrolytes?
Answer: Ionic compounds are generally electrolytes.

Question 30. Which type of chemical bond is present in hydrogen chloride molecule?
Answre: Covalent bonds are present in pure hydrogen chloride molecule.

Question 31. Draw the Lewis dot structure of F₂ molecule.
Answer:

Question 32. Which type of bond is present in HF?
Answer: Covalent bonding is present in HF.

Question 33. By which type of bonding do carbon and hydrogen atoms combine to form methane?
Answer: Covalent single bond.

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds Fill In The Blanks:

Question 1. An example of a solid covalent crystalline compound is ______
Answer: Sugar

Question 2. The total number of valence electrons in the atoms of the molecule of HCN is _______
Answer: 10

Question 3. Ethanol is _______ in water.
Answer: Soluble

Question 4. Chloroform and water are _______ each other.
Answer: Immiscible

Question 5. An aqueous solution of glucose is a _______ of electricity.
Answer: Non-conductor

Wbbse Class 10 Properties Of Covalent Compounds

Question 6. _______is a covalent substance that conducts electricity in a solid state.
Answer: Graphite

Question 7. Only _______ electrons are shown in the Lewis-dot diagram.
Answer: Bonded

Question 8. ______ is an example of a covalent compound which is widely used as a solvent.
Answer: Water

Question 9. Sugar does not ______ in aqueous solution.
Answer: Ionise

Question 10. Covalent bonds are formed due to the formation of _______
Answer: Electron-pairs

Question 11. The covalent bond in HCI is _______ in nature.
Answer: Polar

Question 12. The number of lone pairs of electrons in O atom of H₂O molecule is _______
Answer: 2

Question 13. _________ is a covalent compound in which all the 4 bonds are identical.
Answer: Methane

Question 14. The number of bond pairs of electrons in CO₂ is _______
Answer: 4

Wbbse Class 10 Properties Of Covalent Compounds

Question 15. Among CH₄, C₂H₄ and C₂H₂, triple bond is present in _______
Answer: C₂H₂

Question 16. Covalent bond has a definite _______
Answer: Direction

Question 17. In NH₃ molecule, N-atom has ______ lone pair of electrons.
Answer: 1

Question 18. The structure of covalent compounds are represented by ______ diagram.
Answer: Lewis-dot

Question 19. The number of single bonds present in C₂H₄ is a _______
Answer: 4

Question 20. Number of covalent bonds present i nitrogen molecule is _____
Answer: Three

Ionic and Covalent Bonds Topic B Covalent Bonding And Properties Of Covalent Compounds State Whether True Or False:

Question 1. Lewis-dot diagram cannot explain’ the experimental observation regarding the structure of O₂.
Answer: True

Question 2. Boiling and melting point of covalent compounds are high due to strong force of attraction between the covalent molecules.
Answer: False

Question 3. van der Waal’s force of attraction between two covalent molecules increases as the molecular mass increases.
Answer: True

Question 4. A coordinate bond can be found in H₃O⁺ ion.
Answer: True

Question 5. Covalent compounds can exist as discrete molecules.
Answer: True

Question 6. Covalent compounds can be soft solid at normal temperature.
Answer: True

Question 7. Some transition metal elements can form four bonds between two atoms.
Answer: True

Question 8. Formula mass and molecular mass are not the same for covalent compounds.
Answer: False

Question 9. NaH is a covalent compound.
Answer: False

Question 10. Aqueous solution glucose can conduct electricity.
Answer: False

Question 11. Carbon tetrachloride is a polar solvent.
Answer: False

Question 12. Maximum two covalent bonds can be formed between two atoms.
Answer: False

Question 13. There are 5 covalent bonds present in ethyne molecule.
Answer: True

Ionic and Covalent Bonds Miscellaneous Type Questions Match The Columns:

 

Question 1.

Column A	Column B
CaCO3	1. Ionic and soluble in water
NaCI	2. Contains coordinate bond
NH4+	3. Covalent but conducts electricity in solution
HCI	4. Ionic but insoluble in water

Answer:
CaCO₃: 4. Ionic but insoluble in water:
NaCI: 1. Ionic and soluble in water
NH₄+: 2. Contains coordinate bond
HCI: 3. Covalent but conducts electricity in solution

Question 2.

Column A	Column B
The Duplet rule is applicable for	1. Na
Electrovalency	2. Be
Both ionic and covalent bonds are not formed by	3. PCI5
Expanded octet	4. Electrostatic force of attraction

 

Answer:
The Duplet rule is applicable for: 2. Be
Electrovalency: 4. Electrostatic force of attraction
Both ionic and covalent bonds are not formed by:  1. Na
Expanded octet:  3. PCI₅ 

WBBSE Solutions for Class 10 Physical Science and Environment