Green Science 10 Final PDF (2076)

Magnesium + Chloride Magnesium chloride Calcium carbonate heat Calcium oxide + Carbon dioxide Zinc + Sulphuric acid Zinc sulphate + Hydrogen Potassium Chlorate heat Potassium chloride + Oxygen Formula Equation A formula equation or chemical equation is the chemical reaction expressed by writing symbols and molecular formulae of reactants and products. It is more informative than a word equation. Some examples of formula equations are given below: 2Na + Cl2 2NaCl 2H2 + O2 Mg + Cl2 2H2O CaCO3 MgCl2 Zn + H2SO4 CaO + CO2 2KClO3 ZnSO4 + H2 2KCl + 3O2 Reactants and Products Reactants are chemical substances which take part in a chemical reaction. They are written on the left hand side of the arrow. Products are the chemical substances which are produced after a chemical reaction. They are written on the right hand side of the arrow. Examples ZnCl2 + H2 Do You Know Zn + HCl (Products) (Reactants) Reactants are written on the left hand side HCl + NaOH NaCl + H2O and products are written on the right hand (Reactants) (Products) side of the arrow. Unbalanced Chemical Equation The chemical equation in which the total number of atoms of each element in reactants and products are not equal is called an unbalanced chemical equation. Examples H2O H2 + O2 ZnCl2 + H2 Zn + HCl Na2SO4 + H2O NaOH + H2SO4 GREEN Science (Chemistry) Book-10 151

Balanced Chemical Equation A balanced chemical equation can be defined as the chemical equation written by balancing the total number of atoms of each element in reactants and products. Examples 2H2O 2H2 + O2 ZnCl2 + H2 Zn + 2HCl Na2SO4 + 2H2O 2NaOH + H2SO4 Methods of Writing Balanced Chemical Equation 1. First of all, a chemical change is written correctly in the form of word equation. For example: Sodium + Chlorine Sodium chloride 2. A word equation is written correctly in the form of formula equation. For example: Na + Cl2 NaCl 3. The number of atoms of each element are balanced by using suitable coefficient without changing the molecular formulae of reactants and products. For example: 2Na + Cl2 2 NaCl Information obtained from a Balanced Chemical Equation We can get the following information from a balanced chemical equation 1. The names of reactants and products 2. Symbols and molecular formulae of reactants and products 3. Total number of atoms or molecules of reactants and products 4. Type of chemical reaction 5. Atomic weight and molecular weight of reactants and products Types of Chemical Reaction There are different types of chemical reaction on the basis of process of formation of products from the reactants. Some of the common types of chemical reactions are as follows: 1. Combination reaction 2. Decomposition reaction 3. Displacement reaction 4. Acid- base reaction 1. Combination reaction When carbon (C) burns in air (oxygen), it forms carbon dioxide (CO2). In the chemical reaction, carbon (C) combines with oxygen (O2) and forms a product carbon dioxide 152 GREEN Science (Chemistry) Book-10

(CO2). This type of chemical reaction is called combination reaction. The chemical reaction in which two or more reactants combine to form a single product is called combination reaction. It is also called addition or synthesis reaction. Various factors like heat, light, pressure, etc. are responsible for combination reaction. Some examples of combination reaction are given below: 1. Carbon + Oxygen burn Carbon dioxide C + O2 CO2 [Combination] 2. Potassium + Chlorine Potassium chloride 2K + Cl2 2KCl [Combination] 3. Sodium + Chlorine Sodium chloride 2Na + Cl2 2NaCl 4. Aluminium + Nitrogen Aluminium nitride 2 Al + N2 2AlN 5. Nitrogen + Hydrogen Ammonia N2 + 3H2 2NH3 6. Iron + Sulphur Iron (Ferrous) sulphide Fe + S FeS 7. Iron + Oxygen Ferric (iron) oxide 4Fe + 3O2 2Fe2O3 8. Calcium carbonate + Carbon dioxide + Water Calcium bicarbonate CaCO3 + CO2 + H2O Ca(HCO3)2 9. Sodium + Oxygen Sodium Oxide 4Na + O2 2Na2O 10. Hydrogen + Oxygen Water 2H2 + O2 2H2O 11. Iron + Sulphur Ferrous Sulphide Fe + S FeS GREEN Science (Chemistry) Book-10 153

2. Decomposition reaction When calcium carbonate (CaCO3) is heated, Do You Know it breaks down into two products, viz. calcium oxide (CaO) and carbon dioxide Decomposition reaction is also called (CO2). Here, a single reactant decomposes dissociation or analysis reaction. into two products due to action of heat. When lead nitrate [Pb(NO3)2] is heated, it decomposes into three products, viz. Lead oxide (PbO), Nitrogen dioxide (NO2) and Oxygen (O2). These are the examples of decomposition reaction. So, the chemical reaction in which a single reactant decomposed into two or more products is called decomposition reaction. Various factors like heat, light, electricity, catalyst, etc. are responsible for decomposition reaction. 1. Water electricity Hydrogen + Oxygen 2H2O 2H2 + O2 2. Calcium carbonate heat C alcium oxide + Carbon dioxide CaCO3 D CaO + CO2 3. Ammonium hydroxide heat Ammonia + Water NH4OH NH3 + H2O 4. Copper carbonate heat Copper oxide + Carbon dioxide CuCO3 D CuO + CO2 5. Potassium chlorate heat Potassium chloride + Oxygen 2KClO3 D 2KCl + 3O2 6. Hydrogen peroxide MnO2 W ater + Oxygen 2H2O2 MnO2 2H2O + O2 7. Silver nitrate heat Silver + Nitrogen dioxide + Oxygen 2AgNO3 D 2Ag + 2NO2 + O2 8. Copper nitrate heat Copper + Nitrogen dioxide + Oxygen 2Cu (NO3)2 D 2Cu + 2NO2 + O2 9. Lead nitrate Lead monoxide + Nitrogen dioxide + Oxygen 2Pb (NO3)2 2PbO + 4NO2 + O2 10. Silver oxide Silver + Oxygen 2Ag2O 4Ag + O2 154 GREEN Science (Chemistry) Book-10

3. Displacement reaction When zinc (Zn) reacts with dilute sulphuric Do You Know acid (H2SO4), zinc displaces Hydrogen of the acid, i.e. Sulphuric acid and forms Displacement reaction is also called Zinc sulphate (ZnSO4) and Hydrogen replacement reaction. gas (H2). This type of chemical reaction is called displacement reaction. The chemical reaction in which an atom or radical of a compound is displaced by another element to form new products is called displacement reaction. Displacement reaction is of two types, viz. single displacement reaction and double displacement reaction. i. Single displacement reaction It is the displacement reaction in which one atom or a radical is displaced by another element. Some examples of single displacement reaction are as follows: 1. Potassium + Hydrochloric acid Potassium chloride+ Hydrogen 2K + 2HCl 2KCl + H2 2. Magnesium + Hydrochloric acid Magnesium chloride + Hydrogen Mg + 2HCl MgCl2 + H2 3. Calcium + Sulphuric acid Calcium sulphate + Hydrogen Ca + H2SO4 CaSO4 + H2 4. Iron + Copper sulphate Iron sulphate + Copper Fe + CuSO4 FeSO4 + Cu 5. Magnesium + Zinc chloride Magnesium chloride + Zinc Mg + ZnCl2 MgCl2 + Zn ii. Double displacement reaction It is the chemical reaction in which an atom or radical of a compound is mutually displaced by a radical or an atom of another compound. Some examples of double displacement reaction are as follows: 1. Sodium chloride + Silver nitrate Sodium nitrate + Silver chloride NaCl + AgNO3 NaNO3 + AgCl or, Na+ Cl– + Ag+ NO3– 2. Magnesium chloride + Silver nitrate Magnesium nitrate + Silver chloride MgCl2 + 2Ag NO3 Mg(NO3)2 + 2AgCl GREEN Science (Chemistry) Book-10 155

3. Sodium hydroxide + Ferrous chloride Sodium chloride + Ferrous hydroxide 2NaOH + FeCl2 2NaCl + Fe(OH)2 4. Calcium chloride + Silver nitrate Calcium nitrate + Silver chloride CaCl2 + 2Ag NO3 Ca(NO3)2 + 2AgCl 5. Sodium sulphate + Lead nitrate Lead sulphate + Sodium nitrate Na2SO4 + Pb (NO3)2 PbSO4 + 2NaNO3 6. Magnesium nitride + Water Magnesium oxide + Ammonia Mg3N2 + H2O 3MgO + 2NH3 7. Mercuric chloride + Potassium iodide Potassium chloride + Mercuric iodide HgCl2 + 2KI 2KCl + HgI2 4. Acid - base reaction When Hydrochloric acid (HCl) reacts Do You Know with Sodium hydroxide (NaOH), it forms Sodium chloride (NaCl) and Acid-base reaction is also called neutralization water (H2O). It is an example of an acid- reaction. However, all acid-base reactions are base reaction. The chemical reaction not neutralization reaction. which takes place between an acid and a base to form salt and water is called acid-base reaction. Some examples of acid-base reaction are given below: Acid + Base Salt + Water 1. Hydrochloric acid + Potassium hydroxide Potassium chloride + Water HCl + KOH KCl + H2O 2. Nitric acid + Potassium hydroxide Potassium nitrate + Water HNO3 + KOH KNO3 + H2O 3. Sulphuric acid + Sodium hydroxide Sodium sulphate + Water H2SO4 + 2NaOH Na2SO4 + 2H2O 4. Sulphuric acid + Calcium oxide Calcium sulphate + Water H2SO4 + CaO CaSO4 + H2O 5. Sulphuric acid + Ferrous oxide Ferrous sulphate + Water H2SO4 + FeO FeSO4 + H2O 6. Acetic acid + Sodium hydroxide Sodium acetate + Water CH3COOH + NaOH CH3COONa + H2O 156 GREEN Science (Chemistry) Book-10

7. Hydrochloric acid + Ammonium hydroxide Ammonium chloride + Water HCl + NH4OH NH4Cl + H2O In acid-base reaction, both acid and base lose their properties and form two neutral substances, viz. salt and water. Therefore, acid-base reaction is also called neutralization reaction. Factors that bring out chemical reaction Various conditions are required to bring out chemical reaction. Some of them are discussed below: 1. Simple contact Some chemical reactions take place when the reactants are brought in contact. For example, when sodium is brought in contact with chlorine, chemical reaction takes place. As a result, sodium chloride is formed. Na + Cl2 2NaCl 2. Heat Some chemical reactions take place when reactants are heated. Heat energy increases the kinetic energy of the molecules of reactants which brings the reacting molecules in close contact. As a result, chemical reaction takes place. For example, CaCO3 CaO + CO2 2KClO3 2KCl + 3O2 3. Light Some chemical reactions take place when the reactants are exposed to light. Light energy makes reactant molecules active which brings out chemical change. For example, H2 + Cl2 Sunlight 2HCl 4. Electricity Some chemical reactions take place when electricity is passed through solution state or fused state of reactants. Electricity helps the ions move towards oppositely charged electrodes and chemical reaction takes place. For example, when water is electrolyzed, it decomposes into hydrogen and oxygen. 2H2O 2H2 + O2 5. Pressure Some chemical reactions take place when the reactants are kept under certain pressure. Pressure brings the reacting molecules closer and chemical reaction takes place. For example, 3H2 + N2 200 – 900 atm (pressure) 2NH3 500°C/Fe, Mo GREEN Science (Chemistry) Book-10 157

6. Solution state Some chemical reactions take place when reactants are mixed in solution state. For example, when Sodium Chloride and Silver Nitrate are mixed in solid state, chemical reaction does not occur. But they react in solution state. NaCl + AgNO3 AgCl + NaNO3 7. Catalyst Some chemical reactions take place only in the presence of catalyst. For example, Hydrogen peroxide (H2O2) decomposes in the presence of catalyst Manganese dioxide (MnO2). 2H2O2 MnO2 2H2 + O2 (Catalyst) Factors Affecting the Rate of Chemical Reaction The rate of different chemical reaction is different. Some chemical reaction occur very fast whereas some chemical reaction occur very slowly. Some reactants undergo chemical reaction when they come in contact whereas some chemical reactions do not take place without supplying heat, light, electricity, catalyst, etc. The rate of a chemical Do You Know reaction depends on the concentration physical nature and chemical nature of The rate of chemical reaction is defined as the reactants. Various factors increase or positive change in concentration of a reactant decrease the rate of a chemical reaction. or a product per unit time. Some of the major factors that affect the rate of chemical reaction are given below: 1. Temperature 2. Light 3. Surface area 4. Pressure 5. Catalyst 1. Temperature The rate of chemical reaction increases on increasing the temperature of reactants. Similarly, the rate of chemical reaction decreases on decreasing the temperature of reactants. More heat is supplied to the reacting molecules while increasing the temperature. It provides more kinetic energy to reacting molecules and frequency of collision of these molecules increases to give more products. Increase in temperature increases the rate of dissociation and recombination of reacting molecules. As a result, the rate of chemical reaction increases. Many chemical reactions do not occur without heating the reactants to a certain temperature. Example 2KClO3 360°C 2KCl + 3O2 158 GREEN Science (Chemistry) Book-10

Experiment 1 Objective : To demonstrate that the rate of chemical reaction increases on increasing the temperature of reactants. Materials required : Beakers, dilute sulphuric acid, water, potassium permanganate, potassium thiosulphate, oxalic acid (aq), spirit lamp, stand Procedure • Take two beakers and keep one crystal of oxalic acid in each of them. • Now, add about 10ml of dilute sulphuric acid in each beaker. • Add 5 ml of potassium permanganate solution in each beaker. • Stir the solution in each beaker using a glass rod and observe the solution carefully. • You can see that the solution in both beakers appear pink. • Now, heat one of the beakers upto 60° – 80° C with the help of spirit lamp. • Observe carefully, in which beaker does the pink colour disappear faster? Observation The pink colour of the solution in the beaker disappears which is heated after some time. But the pink colour of the solution does not disappear in the beaker which is not heated. It shows that chemical reaction occurs faster at high temperature. Conclusion From this activity, it can be concluded that the rate of chemical reaction increases on increasing the temperature of reactants. 2. Light Many chemical reactions take place in the presence of sunlight. The rate of a chemical reaction increases in the presence of light. Some examples of chemical reactions that take place in the presence of light (sunlight) are as follows. i. Photosynthesis in green plants 6CO2 + 6H2O Sunlight C6H12O6 + 6O2 ii. Formation of hydrochloric acid H2 + Cl2 Sunlight 2HCl iii. Dissociation of silver bromide 2AgBr Sunlight 2Ag + Br2 iv. Chlorination of methane CH4 + Cl2 CH3Cl + HCl v. Formation of Ozone from oxygen 3O2 UV rays 2O3 GREEN Science (Chemistry) Book-10 159

3. Surface area The rate of a chemical reaction increases if the contact area of reacting molecules is more and vice-versa. If the contact area is more, many reacting molecules come in contact. As a result, the rate of chemical reaction increases. The area of contact can be increased by either of the following methods. i. By breaking down reactants into small pieces ii. By using the reactants in powdered from iii. By using a common solvent Experiment 2 Objective : To demonstrate that the rate of chemical reaction increase on increasing the surface area of reacting molecules. Materials required : Zinc powder, Zinc granules, dilute Hydrochloric acid, beaker (2), measuring cylinder, top pan balance, watch glass (2), glass rod Procedure • Take a measuring cylinder and measure 25 ml of dilute hydrochloric acid. • Keep 25/25 ml of dilute hydrochloric acid in two beakers. • Take a top pan balance and measure 2.5 gram of zinc granules and 2.5 gram of zinc powder. Keep zinc power and zinc granules in a separate watch glass. • Now, pour the zinc powder into a beaker and zinc granules into another beaker simultaneously. Stir the mixture by using a glass rod. • Observe the chemical reaction that occurs in the beakers carefully. In which beaker does the reaction complete faster? Observation The bubbles of gas (hydrogen) evolve earlier in the beaker having zinc powder than in the beaker having zinc granules. Similarly, chemical reaction completes faster in the beaker having zinc powder. It shows that chemical reaction takes place faster in powdered from than the granules, because the area of contact is more in powdered form than that of the granules. From this activity, it can be concluded that the rate of a chemical reaction increases on increasing the area of contact of the reacting molecules. 4. Pressure The rate of chemical reaction of gases molecules depends on the pressure of the reactants. The rate of chemical reaction increases on increasing the pressure of the reacting molecules and vice-versa. Some gases react only in the high pressure of the reacting molecules and vice-versa. Some gases react only in the high pressure. Increase in pressure brings the molecules of reacting gases closer and the rate of chemical reaction increases. Some examples of the chemical reactions that take place due to application of pressure are as follows. i. Formation of ammonia gas by Haber's process N2 + 3H2 2NH3 160 GREEN Science (Chemistry) Book-10

ii. When pressure is applied on the mixture of sulphur and potassium chloride, it explodes. iii. Fire crackers get exploded on applying pressure. 5. Catalyst The chemical substance which is used to increase or decrease the rate of a chemical reaction is called a catalyst. A catalyst remains chemically unchanged throughout the chemical reaction but its presence may increase or decrease the rate of a chemical reaction. The catalyst which increases the rate of a chemical reaction is called positive catalyst. It is used to speed up the rate of chemical reaction. A positive catalyst decreases or lowers the energy required to break down the chemical bond of the molecules of reactants. As a result, the rate of chemical reaction increases. Examples of positive catalyst i. Manganese dioxide (MnO2) acts as a positive catalyst during decomposition of Hydrogen peroxide. (H2O2) 2H2O2 MnO2 2H2O + O2 ii. Manganese dioxide (MnO2) acts as a positive catalyst during decomposition of Potassium chlorate (KClO3). It means that MnO2 speeds up the decomposition of KClO3. 2KClO3 MnO2 2KCl + 2O2 The catalyst which decreases the rate of a chemical reaction is called a negative catalyst. It is used to slow down the rate of a chemical reaction. Example negative catalyst Glycerol [C3H5(OH)3] acts as a negative catalyst during decomposition of Hydrogen peroxide (H2O2). It means that glycerol slows down th decomposition of hydrogen peroxide. 2H2O2 Glycerol 2H2O + O2 Characteristics of catalyst 1. The mass of a catalyst does not change till the end of the chemical reaction. 2. A catalyst remains chemically unchanged throughout the chemical reaction. 3. A catalyst does not initiate a chemical reaction but increases or deceases the rate of chemical reaction. Endothermic reaction and Exothermic reaction Most chemical reactions occur due to change in heat. Some chemical reactions absorb heat whereas some chemical reactions evolve heat during the chemical reaction. On this basis, there are two types of chemical reaction. They are : Endothermic reaction and exothermic reaction. GREEN Science (Chemistry) Book-10 161

The chemical reaction that absorbs heat during the chemical change is called endothermic reaction. Examples N2 + O2 + Heat 2NO CaCO3 + Heat CaO + CO2 2NaCl + Heat 2Na + Cl2 NH4Cl + NaNO2 + Heat NaCl + 2H2O + N2 2KClO3 + Heat 2KCl + 3O2 The chemical reaction that evolves heat during the chemical change is called exothermic reaction. Examples CO2 + Heat C + O2 CH4 + Heat C + 2H2 ZnCl2 + H2 + Heat Zn + 2HCl Ca(OH)2 + Heat CO2 + 2H2O + Heat CaO + H2O CH4 + 2O2 Difference between Endothermic and Exothermic reaction Endothermic Reaction Exothermic Reaction 1. Heat is absorbed during chemical 1. Heat is evolved during chemical reaction. reaction. For example, For example, 2KClO3 + Heat 2KCl + 3O2 C + O2 CO2 + Heat Key Concepts 1. The combination, decomposition or replacement that occurs in the molecules of matter during a chemical change is called chemical reaction. 2. The chemical bond present in the molecules of reactants breaks due to heat, light, electricity, etc. during a chemical change. 3. Losing, gaining or sharing of electrons by an atom to gain stable electronic configuration is the major cause of chemical reaction. 4. The chemical equation in which the total number of atoms of each element in reactants and products are not equal is called an unbalanced chemical equation. 5. A balanced chemical equation can be defined as the chemical equation written by balancing the total number of atoms of each element in reactants and products. 6. The chemical reaction in which two or more reactants combine to form a single product is called combination reaction. It is also called addition or synthesis reaction. 162 GREEN Science (Chemistry) Book-10

7. The chemical reaction in which a single reactant decomposed into two or more products is called decomposition reaction. 8. The chemical reaction in which an atom or radical of a compound is displaced by another element to form new products is called displacement reaction. 9. Double displacement reaction is the chemical reaction in which an atom or radical of a compound is mutually displaced by a radical or an atom of another compound. 10. The chemical reaction which takes place between an acid and a base to form salt and water is called acid-base reaction. 11. In acid-base reaction, both acid and base lose their properties and form two neutral substances, viz. salt and water. Therefore, acid-base reaction is also called neutralization reaction. 12. The rate of different chemical reaction is different. Some chemical reaction occur very fast whereas some chemical reaction occur very slowly. 13. Various factors increase or decrease the rate of a chemical reaction. Some of the major factors that affect the rate of chemical reaction are given below: a. Temperature b. Light c. Surface area d. Pressure e. Catalyst 14. The rate of chemical reaction is defined as the positive change in concentration of a reactant or a product per unit time. 15. The chemical substance which is used to increase or decrease the rate of a chemical reaction is called a catalyst. 16. The chemical reaction that absorbs heat during the chemical change is called endothermic reaction. 17. The chemical reaction that evolves heat during the chemical change is called exothermic reaction. Sequential General Exercise 1 1. Choose the best answer from the given alternatives. a. Magnesium burns in air and forms magnesium oxide. What type of chemical reaction is this? Combination reaction Decomposition reaction Displacement reaction Acid base reaction b. Which of the given chemical reactions is a displacement reaction. Electrolysis of water Reaction between Zinc and Hydrochloric acid Reaction between Sodium and Chlorine Reaction between Magnesium and Sulphuric acid GREEN Science (Chemistry) Book-10 163

c. Which of the following reaction is a neutralization reaction? Reaction between Sulphuric acid and Sodium hydroxide Reaction between Magnesium and Oxygen Reaction between Iron and Copper sulphate Reaction between Potassium and Chlorine d. Which of the given factors is essential for the chemical reaction during photosynthesis? Heat Sunlight Light Electricity e. Which of the following is a positive catalyst? Glycerol Sodium chloride Manganese dioxide Calcium carbonate 2. Answer the following questions. a. Define chemical reaction with any three examples. b. What are reactants and products? c. What is a word equation? Give any three examples. d. What is a chemical equation? Write any three examples. e. Name the four types of chemical reaction. f. Define combination reaction with any three examples. g. Define decomposition reaction with any three examples. h. What is meant by displacement reaction? Name its types. i. Write any three examples of decomposition reaction. j. What is single displacement reaction? Give any two examples. k. What is double displacement reaction? Write any two examples. l. What is acid-base reaction? Write any two examples. m. What is meant by rate of chemical reaction? Name any three factors that affect the rate of chemical reaction. n. How is rate of chemical reaction affected by increase or decrease in temperature? Describe in brief. o. What is the effect of light on the rate of chemical reaction? p. Write down the effect of increase or decrease in surface area of reactants on the rate of chemical reaction. q. What is the effect of pressure on the rate of chemical reaction? 164 GREEN Science (Chemistry) Book-10

r. What is a catalyst? Write its types. s. Define positive and negative catalyst with any one example of each. t. Define exothermic and endothermic reaction with any two examples of each. 3. Differentiate between: a. Reactants and Products b. Word equation and Chemical equation c. Combination reaction and Decomposition reaction d. Positive catalyst and Negative catalyst e. Endothermic reaction and Exothermic reaction 4. Give reason: a. The chemical reaction between Hydrogen and Oxygen is called combination reaction. b. The chemical reaction between Zinc and dilute Hydrochloric acid is called displacement reaction. c. The chemical reaction between sulphuric acid and sodium hydroxide is called neutralization reaction. d. The rate of chemical reaction increases on increasing the temperature of reactants. e. Manganese dioxide is called a positive catalyst. f. Positive catalyst increases the rate of chemical reaction. g. Increase in pressure increases the rate of chemical reaction on gaseous reactants. 5. Convert following unbalanced chemical equations into balanced chemical equations. a. Mg + N2 Mg3N2 b. HCl + K2O KCl + H2O c. K + O2 K2O d. Fe + CuSO4 FeSO4 + Cu e. HNO3 + Ca(OH)2 Ca(NO3)2 + H2O f. Au + Cl2 AuCl3 g. H2SO4 + NaOH Na2SO4 + H2O h. CaCO3 CaO + CO2 i. AgNO3 Ag + NO2 + O2 j. Na2SO4 + Pb (NO3)2 PbSO4 + NaNO3 6. Write down the given word equations in the form of balanced chemical equations. a. Nitrogen + Hydrogen Ammonia b. Hydrogen + Oxygen Water c. Aluminum + Nitrogen Aluminium nitride GREEN Science (Chemistry) Book-10 165

d. Magnesium + Chlorine Magnesium chloride e. Calcium oxide + Water Calcium hydroxide f. Calcium carbonate Calcium oxide + Carbon dioxide g. Zinc + Hydrochloric acid Zinc chloride + Hydrogen h. Copper + Oxygen Copper oxide i. Potassium chlorate Potassium chloride + Oxygen j. Nitric acid + Calcium hydroxide Calcium nitrate + Water 7. Describe an activity to demonstrate that the rate of a chemical reaction increases on increasing the temperature of reactants. 8. Explain an activity to demonstrate that the rate of a chemical reaction increases on increasing the area of contact of reactants. Grid-based Exercise 2 Group ‘A’ (Knowledge Type Questions) (1 Mark Each) 1. What is a chemical reaction? Give one example. 2. Define formula equation with one example. 3. State any four factors that bring out a chemical change. 4. What is a precipitate ? Give one example. 5. Define negative catalyst with one example. 6. Define exothermic reaction with one example. 7. What is combination reaction? Give one example. 8. Define decomposition reaction with one example. 9. What is acid-base reaction? Give one example. 10. What is meant by the rate of chemical reaction? 11. Define endothermic reaction with one example. 12. What is displacement reaction? Write with examples. 13. Write a balanced chemical equation of decomposition reaction which is carried out by the catalyst. 14. Write two information which can be obtained from the balanced chemical equation. 15. Write an example of single displacement chemical reaction. Group ‘B’ (Understanding Type Questions) (2 Marks Each) 16. Write any two differences between reactants and products. 17. Acid-base reaction is also called neutralization reaction, why? 18. The rate of a chemical reaction increases on increasing temperature, why? 19. Write any two differences between synthesis reaction and dissociation reaction. 20. The rate of a chemical reaction increases on powdering the reactants, why? 21. Write any two limitations of balanced chemical equation. 166 GREEN Science (Chemistry) Book-10

22. Which type of chemical equation is given below? Define it. Fe + CuSO4 → FeSO4 + Cu 23. How does heat enhance the rate of chemical reaction? Write in short. 24. The rate of a chemical reaction increases on increasing the concentration of reactants, why? 25. Write any two differences between exothermic chemical reaction and endothermic chemical reaction. Group ‘C’ (Application Type Questions) (3 Marks Each) 26. In what condition do sodium chloride and silver nitrate react? Write the balanced chemical equation of that reaction. 27. Change the given word equation into balanced chemical equation. What type of chemical reaction is it ? What is the role of MnO2 in this reaction ? Potassium Chlorate MnO2 Potassium Chloride + Oxygen 28. Write the balanced chemical equation for the following word equations: i. Iron + Oxygen → Ferric Oxide ii. Potassium Chlorate Heat Potassium Chloride + Oxygen 29. Give one example of exothermic chemical reaction. Write any two applications of catalyst. 30. Describe an experiment to demonstrate that the rate of chemical reaction increases on increasing the surface area of reactants. Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each) 31. Change the given word equations into formula equation. Also, write down the type of the chemical equation. i. Calcium bicarbonate → Calcium carbonate+Water+Carbon dioxide ii. Aluminium + Hydrochloric acid → Aluminium chloride + Hydrogen 32. How does the concentration of sodium thiosulphate affect the rate of chemical reaction in between the hydrochloric acid and sodium thiosulphate? Write the chemical equation of the reaction between more active metal and more active non-metal. What type of chemical reaction is it ? 33. Write a balanced chemical equation of decomposition reaction which is carried out by the catalyst. Which type of chemical equation is given below? Define it. AgNO3 + CaCI2 → AgCI + Ca(NO3)2 34. Describe an experiment to demonstrate that the rate of chemical reaction increases on increasing the temperature of reactants. 35. Chemical reaction takes place when iron dust is added into Copper sulphate solution but no reaction takes place when Copper dust is added into Ferrous sulphate solution. Why ? Describe in brief the effect of physical state of reactants in the rate of chemical reaction. GREEN Science (Chemistry) Book-10 167

UNIT Acid, Base and Salt 9 Weighting Distribution Theory : 7 Practical: 2 Before You Begin Matter can be defined as anything having mass and volume. All matter have mass and they occupy space. For example, air, soil, water, milk, stone, brick, wood, smoke, cloud, petrol, kerosene, iron, gold, plastic, etc. Sound, light, shadow, heat, etc. do not have mass and volume. So they are not matter. Matter can be soluble or insoluble, transparent or opaque and good conductor or bad conductor of heat and electricity. Matter exist in three different states, viz. solid, liquid and gas. Same matter can exist in three different states. For example, water can exist in all three states, viz. solid (ice), liquid (water) and gas (vapour). Learning Objectives Syllabus After completing the study of this unit, students will be able to: i. introduce acid, base and salt with examples. • Acid-Introduction and types ii. explain the properties and uses of acid, base and salt. • Physical and chemical properties of acids iii. state the use of acid, base and salt in our daily life. • Uses of acids iv. write neutralization reaction between acid and base and balance neutralization reactions. • Base and alkali-Introduction • Physical and chemical properties of bases/alkalis • Uses of bases • Salt: Introduction • Properties and uses of salts • Indicators, pH and pH scale • Neutralization reaction Glossary: A dictionary of scientific/technical terms acid : the substances that produces H+ ions when dissolved in water base : metal oxide or metal hydroxide alkali : the base that dissolves in water salt : the substance formed by the reaction between an acid and a base 168 GREEN Science (Chemistry) Book-10

Acids Do You Know When we drink lemon juice, we feel sour due to the presence of acid (i.e. citric acid) Most acids are sour in taste. But it is in lemon juice. Similarly, pickles containing dangerous to touch and taste acids in vinegar are also sour in taste. It shows that laboratory as acids burn our skin and acids possess sour taste. The word 'acid' tongue. refers to a sour substance. However, it does not mean that all acids are sour. Most acids are sour. The word 'acid' has been derived from the Latin word 'acidus' which means sour in taste. Citric acid, lactic acid, carbonic acid, hydrochloric acid, sulphuric acid and nitric acid are the examples of acids. The foods having acids are sour in taste. Fruits like lemon, orange, etc. taste sour due to presence of citric acid. Grape fruit tastes sour due to the presence of tartaric acid. Similarly, sour milk contains lactic acid, vinegar contains acetic acid and vitamin C contains ascorbic acid. The chemical substances which give hydrogen (H+) ions when dissolved in water are called acids. For example: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid (H2CO3), etc. HCl +H2O H+ + Cl– H2SO4 +H2O 2H+ + SO4– – HNO3 +H2O H+ + NO3– H2CO3 +H2O 2H+ + CO3– – Fig. 9.1 Sulphuric acid (H2SO4) Nitric acid (HNO3) Hydrochloric acid (HCl) There are two types of acids on the basis of source or chemical nature. They are (i) Inorganic acids and (ii) Organic acids. i. Inorganic acids The acids which are obtained from minerals are called inorganic acids. Acids like Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid (H2CO3), etc are inorganic acids. Inorganic acids are also obtained from minerals. So they are called mineral acids. They are commonly used in laboratories and industries. Inorganic acids are strong in nature. GREEN Science (Chemistry) Book-10 169

ii. Organic acids The acids which are obtained from living organisms (plants and animals) are called organic acids. Citric acid, Acetic acid, Tartaric acid, Formic acid, Lactic acid, etc. are examples of organic acids. These acids are weak in nature. Citric acid is found in lemon. Ascorbic acid and tartaric acid are found in fruits and vegetables. Similarly, formic acid is found in ant bite. There are two types of acids on the basis of strength. They are: (i) Strong acids and (ii) Weak acids. i. Strong acids The acids which produce a high concentration of hydrogen (H+) ions when dissolved in water are called strong acids. They are more corrosive in nature. Examples: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), etc. These acids produce high concentration of hydrogen ions (H+) when dissolved in water. They have low pH value and are good conductors of electricity. ii. Weak acids The acids which produce a low concentration of hydrogen (H+) ions when dissolved in water are called weak acids. They are less corrosive in nature. Examples: Carbonic acid (H2CO3), Acetic acid (CH3COOH), Formic acid (HCOOH), etc. These acids produce a low concentration of H+ ions when dissolved in water. They have high pH value. They are poor conductors of electricity. Physical properties of acids Do You Know 1. Most acids are sour in taste. Most acids are sour in taste but acids like 2. They change blue litmus paper into red Boric acid, Stearic acid and Salicylic acid and methyl orange into red. are not sour. 3. They are corrosive in nature. 4. They do not change the colour of phenolphthalein. 5. They burn our skin. Chemical properties of Acids 1. Acids reacts with active metals like Zn, Mg, Na, etc. and form hydrogen gas. Dilute acid + Metal Salt + Hydrogen gas 2HCl + Zn ZnCl2 + H2 H2SO4 + Zn ZnSO4 + H2 2HCl + Mg MgCl2 + H2 H2SO4 + Mg MgSO4 + H2 2HNO3 + Zn Zn(NO3)2 + H2 170 GREEN Science (Chemistry) Book-10

2HCl + Na 2NaCl + H2 6HNO3 + 2Al 2Al(NO3)3 + 3H2 2. Acids react with bases and form salt and water. Acid + Base Salt + Water H2SO4 + KOH K2SO4 + H2O HCl + NaOH NaCl + H2O 2HCl + CaO CaCl2 + H2O H2SO4 + 2NaOH Na2SO4 + 2H2O H2SO4 +CaO CaSO4 + H2O HNO3 + KOH KNO3 + H2O 2HNO3 + Ca(OH)2 Ca(NO3)2 + 2H2O HCl + KOH KCl + H2O 3. Acids dissolve in water and produce hydrogen ions. HCl +H2O H+ + Cl– H2SO4 +H2O 2H+ + SO4– – H+ + NO3– HNO3 +H2O H+ + CH3COO – CH3COOH +H2O 4. Acids react with carbonates and form salt, water and carbon dioxide gas. Acid + Carbonates Salt + Water + Carbon dioxide 2HCl + Na2CO3 2NaCl + H2O + CO2 2HCl + MgCO3 MgCl2 + H2O + CO2 H2SO4 + MgCO3 MgSO4 + H2O + CO2 5. Acids react with bicarbonates and form salt, water and carbon dioxide gas. Acid + Bicarbonates Salt + Water + Carbon dioxide HCl + KHCO3 KCl + H2O + CO2 2HCl + CaCO3 CaCl2 + H2O + CO2 H2SO4 + 2NaHCO3 Na2SO4 + 2H2O + 2CO2 H2SO4 +Mg(HCO3)2 MgSO4 + 2H2O + 2CO2 H2SO4 + Ca(HCO3)2 CaSO4 + 2H2O + 2CO2 2HCl + Ca(HCO3)2 CaCl2 + 2H2O + 2CO2 2HNO3 + Na2CO3 2NaNO3 + H2O + CO2 GREEN Science (Chemistry) Book-10 171

Uses of Acids 1. Hydrochloric acid, sulphuric acid and nitric acid are used in science laboratories and industries. 2. Sulphuric acid is used for making chemical fertilizers, drugs and detergents. 3. Hydrochloric acid is used in tanning and printing industries. 4. Nitric acid is used for making plastics, dyes and explosives. 5. Carbonic acid is used in soft drinks like coca-cola, soda water, beer, etc. 6. Acetic acid (vinegar) is used in pickles. 7. Carbolic acid (phenol) is used to kill germs. 8. Citric acid is used as a source of vitamin C. Some acids of daily use and their sources are given below. S.N. Acids Sources 1. Lemon, tomato 2. Citric acid Grape fruit 3. Tartaric acid Red ant 4. Formic acid Milk, curd 5. Lactic acid Sour fruits 6. Ascorbic acid Chari amilo Oxalic acid Activity 1 Take a test tube and keep 10 ml of dilute hydrochloric acid. Now, keep, a piece of magnesium. Take another test tube and keep 10 ml of acetic acid in it. Keep a piece of magnesium. What do you observe? Write down the conclusion of this activity. Bases Do You Know Bases are metallic oxides and metallic The bases that dissolve in water are called hydroxides which react with acids and alkalis. But bases like CuO, HgO, BaO, produce salt and water. For example, PbO, etc. do not dissolve in water. So, Sodium oxide (Na2O), Calcium oxide all alkalis are bases but all bases are not (CaO), Magnesium oxide (MgO), Sodium alkalis. hydroxide (NaOH), Calcium hydroxide [Ca(OH)2], etc. The bases that dissolve in water and produce hydroxyl (OH–) ions are called alkalis. Na2O + H2O 2NaOH (Alkali) NaOH +H2O Na+ +OH– 172 GREEN Science (Chemistry) Book-10

K2O + H2O 2KOH KOH +H2O K+ + OH– Sodium hydroxide (NaOH), Magnesium hydroxide [Mg(OH)2], Potassium hydroxide (KOH), Calcium hydroxide [Ca(OH)2], etc. are examples of alkalis. Bases are bitter in taste. Fig. 9.2 Magnesium hydroxide Calcium hydroxide Sodium hydroxide Differences between Acids and Bases Acids Bases 1. Acids produce hydrogen (H+) ions 1. Bases produce hydroxyl (OH–) ions when dissolved in water. when dissolved in water. 2. They turn blue litmus paper into red. 2. They turn red litmus paper into blue. Types of bases On the basis of strength, there are two types of bases, viz. (i) Strong bases and (ii) Weak bases. a. Strong bases The bases that produce a high concentration of hydroxyl (OH–) ions in water are called strong bases. Examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Magnesium hydroxide [Mg(OH)2], etc. Strong bases have a high pH value. Their rate of decomposition is more than that of weak bases. b. Weak bases Do You Know The bases that produce a low concentration Strong bases/alkalis burn our skin, So, of hydroxyl (OH–) ions in water are called we should not touch bases/alkalis in a weak bases. Examples: Ammonium laboratory. hydroxide (NH4OH), Copper hydroxide ([Cu)OH)2], Ferric hydroxide [Fe(OH)3], etc. Ammonium hydroxide (NH4OH) is Weak bases have a low pH value. Their rate called a weak base because it produces of decomposition is less than that of strong a low concentration of hydroxyl ions bases or alkalis. (OH–) when dissolved in water. GREEN Science (Chemistry) Book-10 173

Physical Properties of Bases/Alkalis 1. Most bases are bitter in taste. 2. Their solutions have a soapy touch. 3. They turn red litmus into blue, methyl orange into yellow and phenolphthalein into pink. 4. Strong alkalis dissolve oil and grease. 5. They burn our skin. So, we should not touch strong alkalis in science laboratories. Chemical Properties of Bases/Alkalis 1. Bases react with acids and form salt and water. Base/Alkali + Acid Salt + Water 2NaOH + H2SO4 Na2SO4 + H2O NaOH + HCl KOH + HCl NaCl + H2O 2NaOH + H2SO4 KOH + HNO3 KCl + H2O MgO + 2HCl Ca(OH)2 + 2HNO3 Na2SO4 + H2O Ca(OH)2 + 2HCl KNO3 + H2O MgCl2 + H2O Ca(NO3)2 + 2H2O CaCl2 + 2H2O 2. Bases/Alkalis react with carbon dioxide and form carbonate and water. Base/Alkali + Carbon dioxide Carbonate + Water LiOH + CO2 Li2CO3 + H2O 2NaOH + CO2 Na2CO3 + H2O Mg(OH)2 + CO2 MgCO3 + H2O 2KOH + CO2 K2CO3 + H2O Ca(OH)2 + CO2 CaCO3 + H2O 3. Alkalis dissolve in water and produce hydroxyl (OH–) ions. +H2O NaOH Na+ + OH– Mg(OH)2 +H2O Mg++ + 2OH– KOH +H2O K+ + OH– Ca(OH)2 +H2O Ca++ + 2OH– NH4OH +H2O NH4+ + OH– 174 GREEN Science (Chemistry) Book-10

4. Alkalis react with Ammonium salts and produce salt, water and ammonia gas. Alkali + Ammonium salt Salt + Water + Ammonia gas Ca(OH)2 + 2NH4Cl CaCl2 + 2H2O + 2NH3 NaOH + NH4Cl NaCl + 2H2O + NH3 Ca(OH)2 + (NH4)2CO3 CaCO3 + 2H2O + 2NH3 Mg(OH)2 + (NH4)2CO3 MgCO3 + 2H2O + 2NH3 5. Alkalis separate insoluble metal hydroxides when they are kept in heavy metal salts. Alkali + Heavy metal salts Salt + Insoluble hydroxide 2NaOH + CuSO4 Na2SO4 + Cu(OH)2 3NH4OH + FeCl3 3NH4Cl + Fe(OH)3 2KOH + ZnCl2 2KCl + Zn(OH)2 Uses of Bases/Alkalis 1. Sodium hydroxide (NaOH) or Caustic Soda is used for making soft soap, paper, etc. It is also used to purify petroleum products. 2. Potassium hydroxide (KOH) or Caustic potash is used for making soft soap, chemical fertilizers, etc. 3. Aluminium hydroxide [Al(OH)3] is used to reduce hyperacidity. 4. Calcium hydroxide [Ca(OH)2] or lime water is used in laboratory and to reduce hardness of water. It is also used to reduce acidity of soil and to prepare bleaching powder. 5. Ammonium hydroxide (NH4OH) is used for making nitric acid, chemical fertilizers, dyes and plastics. 6. Calcium oxide (CaO) or quick lime is used for softening hard water, for making cement and in purification of sugar. 7. Magnesium hydroxide [Mg(OH)2] is used as an antacid to reduce hyperacidity. Activity 2 Take a magnesium ribbon and burn it. The magnesium burns in air with a very bright flame and forms a white powder, i.e. Magnesium oxide. – Collect the Magnesium oxide in a beaker and add a few drops water. The magnesium oxide combines with water and forms an alkali, i.e., Magnesium hydroxide [Mg (OH)2]. Mg (s) + O2 (g) 2MgO (s) MgO(s) + H2O (l) Mg (OH)2 (aq) (Alkali) In this way, we can produce an alkali from magnesium ribbon. GREEN Science (Chemistry) Book-10 175

Salt A salt is a chemical substance formed by the chemical reaction between an acid and a base. Examples: (i) Sodium chloride (NaCl), (ii) Potassium chloride (KCl), (iii) Sodium sulphate (Na2SO4), (iv) Copper sulphate (CuSO4), (v) Ammonium chloride (NH4Cl), etc. Fig. 9.3 Copper sulphate Ammonium chloride Sodium chloride Salt can also be defined as a substance formed by partial or complete replacement of hydrogen atom by a metal or ammonium radical. KOH + H2SO4 KHSO4 + H2O (Partial displacement) 2KOH + H2SO4 K2SO4 + H2O (Complete displacement) Most salts are neutral but some are acidic and some are basic in nature. The process in which an acid reacts with a base and forms two neutral substances, i.e. salt and water is called neutralization reaction. A salt consists of two types of radicals. They are acid radical or non-metallic radical and basic radical or metallic radical. The radical obtained from an acid is called an acid radical or non- metallic radical. Similarly, the radical obtained from a base is called basic or metallic radical. In salt NaCl, Na+ is a basic radical as it comes from the base i.e. Na2Oand Cl– is an acid radical as it comes from acid, i.e. HCl. There are different types of salt on the basis of chemical nature and method of formation. The different types of salts are given below: 1. Neutral salt 2. Acidic salt 3. Basic salt 4. Acid salt 5. Base salt 6. Hydrated salt 1. Neutral salt or Normal salt Neutral salt is the salt which is formed by the chemical reaction between a strong acid and a strong alkali or a weak acid and a weak base. Neutral salts are formed by complete displacement of hydrogen from acids. Examples of neutral salts or normal salts: i. Sodium chloride (NaCl) ii. Magnesium chloride (MgCl2) iii. Potassium chloride (KCl) iv. Sodium nitrate (NaNO3) 176 GREEN Science (Chemistry) Book-10

v. Potassium nitrate (KNO3), etc. Base/Alkali + Acid Neutral salt + Water NaOH + H2SO4 Na2SO4 + H2O KOH + HCl KCl + H2O KOH + H2O + HNO3 KNO3 2. Acidic salt Acidic salt is the salt which is formed by chemical reaction between a strong acid and a weak base (alkali). Examples of acidic salts: ii. Copper chloride (CuCl2) i. Copper sulphate (CuSO4) iv. Lead nitrate [Pb(NO3)2] iii. Lead chloride (PbCl2) Acidic salt + Water v. Lead sulphate (PbSO4), etc. Strong acid + Weak acid H2SO4 + CuO CuSO4 + H2O H2SO4 + 2NH4OH (NH4)2SO4 + H2O 2HNO3 + CuO Cu(NO3)2 + H2O HCl + NH4OH NH4Cl + H2O 3. Basic salt Basic salt is the salt which is formed by the chemical reaction between a weak acid and a strong alkali (base). Examples of basic salts: i. Sodium carbonate (Na2CO3) ii. Potassium carbonate (K2CO3) iii. Magnesium carbonate (MgCO3) iv. Sodium acetate (CH3COONa), etc. Weak acid + Strong alkali Basic salt + Water H2CO3 + 2NaOH Na2CO3 + 2H2O CH3COOH + KOH CH3COOH + NaOH CH3COOK + H2O CH3COONa + H2O 4. Acid salt Acid salt is the salt which is formed by partial replacement of hydrogen (H+) ion by a metal. Examples of acid salts: i. Sodium bicarbonate (NaHCO3) ii. Potassium bicarbonate (KHCO3) iii. Magnesium bicarbonate [Mg (HCO3)2] iv. Calcium bicarbonate [Ca (HCO3)2], etc. NaOH + H2CO3 NaHCO3 + H2O KOH + H2CO3 KHCO3 + H2O Mg(OH)2 + 2H2CO3 Mg(HCO3)2 + H2O NaOH + H2SO4 NaHSO4 + H2O GREEN Science (Chemistry) Book-10 177

5. Base salt Base salt is the salt formed by partial replacement of hydroxyl (OH–) radical of a base by an acidic radical. Examples of base salts: i. Zinc hydroxychloride [Zn(OH) Cl] ii. Lead hydroxychloride [Pb(OH)Cl], etc. HCl + Zn(OH)2 Zn(OH)Cl + H2O HCl + Pb(OH)2 Pb(OH)Cl + H2O Activity 3 Take five test tubes and prepare a solution of NaCl, CuSO4, Na2CO3, NaHCO3 and Zn(OH)Cl separately. Identify normal salt, acid salt, base salt, acidic salt and basic salt using indicators (litmus paper, phenolphthalein and methyl orange). 6. Hydrated salt Hydrated salt is the salt which contains certain molecules of water associated with it. These water molecules are called water of crystallization or water of hydration. Examples of hydrated salts: i. Sodium carbonate decahydrate (Na2CO3. 10 H2O) ii. Calcium sulphate heptahydrate(CaSO4 . 7H2O) iii. Copper II sulphate pentahydrate (CuSO4. 5H2O) or Blue vitriol iv. Ferrous sulphate heptahydrate (FeSO4. 7H2O) v. Zinc sulphate heptahydrate (Zn SO4. 7H2O) or White vitriol vi. Magnesium sulphate heptahydrate (MgSO4.7H2O) or Epsom salt vii. Sodium sulphate decahydrate (Na2SO4.10H2O) Fig. 9.4 Sodium carbonate Zinc sulphate 178 GREEN Science (Chemistry) Book-10

Properties of salts 1. Most salts are bitter in taste. Some are salty (e.g. NaCl) and others are tasteless. 2. Most salts are neutral but some may be acidic or basic. 3. Most salts dissolve in water. All salts of Na, K and NH4 are soluble in water. All nitrate and bicarbonate salts also dissolve in water. All chloride salts are water soluble except chloride salts of Ag and Pb. All sulphate salts are water soluble except sulphates of Pb and Ba. 4. They may be white, colourless or colourful. The salts of Cu, Fe, Mn, Cr, etc. are coloured. 5. They conduct electricity in solution or molten state. 6. They have high melting point and boiling point. Uses of salts Do You Know 1. Common salt (NaCl) is used in our foods. Salts of metals like Na, K, Mg, Ca, Al It is also used as preservative and in and Ba are white or colourless. the manufacture of sodium hydroxide, hydrochloric acid and washing soda. Salts of metals like Cu, Co, Mn, Fe, Ni and Cr are colourful. 2. Copper sulphate (CuSO4) is used to The salts that dissolve in water can be make fungicides and for electroplating. electrolysed. 3. Calcium sulphate (CaSO4.7H2O) is used for making cement, chalk and plastering of fractured bones. 4. Ammonium sulphate [(NH4)2SO4] and Potassium nitrate (KNO3) are used for making chemical fertilizers. 5. Sodium bicarbonate (NaHCO3) is used for making baking powder and to make fire extinguisher. 6. Magnesium sulphate (MgSO4) is used for treating constipation. 7. Sodium carbonate (Na2CO3) is used for making glass, soap and detergent. 8. Ammonium chloride (NH4Cl) is used in dry cells as electrolyte. 9. Zinc sulphate (ZnSO4) is used to make white pigment. 10. Anhydrous Ferrous sulphate (FeSO4) is used to make medicines for anaemia. 11. Copper sulphate (CuSO4) is used for electroplating and preserving woods. 12. Sodium carbonate (Na2CO3) is used to make soap, detergent and glass. 13. Silver nitrate is used as a laboratory reagent. Indicators We cannot identify whether a given chemical substance is an acid, base or salt just by observing it. We use some chemicals to identify them. These chemicals are called indicators. These chemical substances are used to indicate whether a substance is acidic, basic or neutral in nature. For example, litmus paper (red and blue), methyl orange, phenolphthalein, etc. These indicators are called ordinary indicators. GREEN Science (Chemistry) Book-10 179

Fig.9.5 Blue litmus paper Methyl orange Fig. Red litmus paper Indicators are obtained from different parts of plants like roots, flowers, leaves, etc. These parts are collected, crushed and mixed with organic solvents to get indicators. Following table shows the various indicators and their effects in acid, base and salt. S.No. Indicators Colour in acid Colour in basic Colour in neutral solution solution salt solution 1. Litmus paper (red) No change in Changes into No change in 2. Litmus paper (blue) colour blue colour 3. Methyl orange Changes into No change in No change in red colour colour 4. Phenolphthalein Changes into Changes into No change in red yellow colour No change in Changes into No change in colour pink colour Universal Indicator Ordinary indicators can indicate whether a substance is acidic, basic or neutral in nature but cannot measure the strength. Therefore, a special kind of indicator is used to measure the strength of the given substance which is called the universal indicator. So, a universal indicator is a special kind of indicator which is used to measure the strength of acidity or alkalinity of a solution. It changes colour when kept in an acidic, basic or neutral solution which is strength of the solution. A universal indicator is prepared by mixing several ordinary indicators of different colour. 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 9.6 acidic neutral alkaline pH colour chart 180 GREEN Science (Chemistry) Book-10

pH and pH Scale A pH is the measure of hydrogen ion concentration present in a solution. It is measured by using pH paper and pH meter. A pH scale is the standard scale which is used to Fig. measure the strength of acidic or alkaline solution. It consists of numbers 1 to 14 with their corresponding 9.7 colours in the scale. The pH value 1 to 6 represents acidity, pH value 7 represents neutrality and pH value 8 to 14 represents the basicity or alkalinity. The solution having pH value 1 is the strongest acid and that having 14 is the strongest alkali. pH meter Acidity increases Neutral Alkalinity increases 7 pH 1 23 456 8 9 10 11 12 13 14 Red Green Rose Yellow Light green Greenish blue Blue Deep blue The pH value of some common chemicals Acidic Chemicals pH value Neutral Hydrochloric acid (HCl) 1 Basic/Alkaline Sulphuric acid (H2SO4) 1.2 Lemon juice (citric acid) 2.5 Carbonic acid, vinegar 3 Common salt solution Water 7 Sugar solution Human blood 7.3 Baking soda 8.5 Washing soda 11.5 Sodium hydroxide (NaOH) 13 Activity 4 Take red and blue litmus paper and three test tubes. Mark the test tubes 1, 2 and 3. Keep the solution of acid in test tube 1, solution of alkali in test tube 2 and solution of common salt in test tube 3. Now, take red litmus papers and immerse one litmus paper separately in each test tube. Observe the change in colour. Take blue litmus papers and repeat the above activity. Prepare a chart after your observation. GREEN Science (Chemistry) Book-10 181

Activity 5 Take solution of acids, bases and salts in different test tubes. Measure the pH value of each by using a pH meter. Neutralization Reaction The chemical reaction that takes place between an acid and a base to from neutral substances, i.e. salt and water is called neutralization reaction. During chemical reaction, both acid and base lose their properties and form two neutral substance, i.e. salt and water. So, acid-base reaction is called neutralization reaction Examples: Strong acid + Strong base Neutral salt + Water HCl + NaOH NaCl + H2O H2SO4 + 2KOH K2SO4 + 2H2O HNO3 + NaOH NaNO3 + H2O H2SO4 + 2NaOH Na2SO4 + 2H2O Weak acid + Weak base Neutral salt + Water H2CO3 + Cu(OH)2 CuCO3 + 2H2O H2CO3 + Pb(OH)2 PbCO3 + 2H2O H2CO3 + NH4OH (NH4)2CO3 + H2O Application of Neutralization Reaction 1. Neutralization reaction is utilized to treat hyperacidity or gastritis. Magnesium hydroxide [Mg (OH)2] is used as an antacid to neutralize the acidity caused by Hydrochloric acid (HCl) in our stomach. 2. Calcium oxide or lime (CaO) is used by farmers to neutralize the acidity of soil. If the soil is basic, it is treated with compost made of rotting vegetables or leaves. The acidic gas formed from the decomposition of compost neutralizes the alkalis in the basic soil. 3. Soap and baking powder are used to neutralize the effect of acidic poison of sting of red ant and bees. 4. Acetic acid or vinegar is used to neutralize the acidity caused by the sting of yellow bumble bee and wasp. 5. Tooth decay occurs due to acid produced during decomposition of food particles in the mouth. It is neutralized by brushing teeth with alkaline toothpaste. 182 GREEN Science (Chemistry) Book-10

Key Concepts 1. The word 'acid' has been derived from the Latin word 'acidus' which means sour in taste. 2. The chemical substances which give hydrogen (H+) ions when dissolved in water are called acids. For example: Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid (H2CO3), etc. 3. The acids which are obtained from minerals are called inorganic acids. Acids like Hydrochloric acid (HCl), Sulphuric acid (H2SO4), Nitric acid (HNO3), Carbonic acid (H2CO3), etc are inorganic acids. 4. The acids which are obtained from living organisms are called organic acids. Citric acid, Acetic acid, Tartaric acid, Formic acid, Lactic acid, etc. are examples of organic acids. 5. The acids which produce a high concentration of hydrogen (H+) ions when dissolved in water are called strong acids. 6. The acids which produce a low concentration of hydrogen (H+) ions when dissolved in water are called weak acids. 7. Bases are metallic oxides and metallic hydroxides which react with acids and produce salt and water. For example, Sodium oxide (Na2O), Calcium oxide (CaO), Magnesium oxide (MgO), Sodium hydroxide (NaOH). 8. The bases that produce a high concentration of hydroxyl (OH–) ions in water are called strong bases. Examples: Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Magnesium hydroxide [Mg(OH)2], Calcium hydroxide [Ca(OH)2], etc. 9. The bases that produce a low concentration of hydroxyl (OH–) ions in water are called weak bases. Examples: Copper hydroxide ([Cu)OH)2], Ferric hydroxide [Fe(OH)3], etc. 10. Salt can also be defined as a substance formed by partial or complete replacement of hydrogen atom by a metal or ammonium radical. 11. Neutral salt is the salt which is formed by the chemical reaction between a strong acid and a strong alkali or a weak acid and a weak base. 12. Acidic salt is the salt which is formed by chemical reaction between a strong acid and a weak base (alkali). 13. Basic salt is the salt which is formed by the chemical reaction between a weak acid and a strong alkali (base). 14. Acid salt is the salt which is formed by partial replacement of hydrogen (H+) ion by a metal. 15. Base salt is the salt formed by partial replacement of hydroxyl (OH–) radical of a base by an acidic radical. 16. Hydrated salt is the salt which contains certain molecules of water. 17. A special kind of indicator is used to measure the strength of the given substance which is called the universal indicator. 18. The chemical reaction that takes place between an acid and a base to from neutral substances, i.e. salt and water is called neutralization reaction. GREEN Science (Chemistry) Book-10 183

Sequential General Exercise 1 1. Choose the best answer from the given alternatives. a. Which of the given substance is a weak acid? HCl H2SO4 HNO3 H2CO3 b. ....................... is used as a source of vitamin C. Carbonic acid Citric acid Sulphuric acid Acetic acid c. Bases react with carbon dioxide and form ...................... and water. nitrates carbonates sulphates oxides d. Which of the following is a neutral salt? NaHSO4 CuSO4 NaCl NaHCO3 e. Which of the given alkalis is used as an antacid? Mg(OH)2 KOH NaOH Ca(OH)2 2. Answer the following questions. a. Define acids with any four examples. b. What are strong acids? Give any two examples of weak acids. c. Define organic acids. How do they differ from inorganic acids? d. Write any three physical properties of acids. e. Write any two chemical properties of acids. f. Write down the uses of given acids. i. Acetic acid ii. Nitric acid iii. Carbonic acid iv. Sulphuric acid v. Tartaric acid vi. Formic acid g. Define bases with any five examples. h. What are alkalis? Give any three examples of the bases that dissolve in water. i. Write any three physical properties and two chemical properties of bases (alkalis). j. Write any four uses of alkalis. k. Define salt and write any five examples. l. Write any four properties and four uses of salt. m. Define indicator and universal indicator. 184 GREEN Science (Chemistry) Book-10

n. What is pH? Write down the pH value of the strongest acid, neutral salt and the strongest alkali. o. What is a pH-scale? 3. Give reason. a. We should not touch and taste acids in a science laboratory. b. Sodium chloride is called a neutral salt. c. All bases are not alkalis but all alkalis are bases. d. Methyl orange is called an indicator. e. Universal indicator is better than an ordinary indicator. f. We eat aluminium hydroxide to reduce hyperacidity. g. Acids are sour in taste. 4. Differentiate between: a. Inorganic acids and Organic acids b. Base and Alkali c. Acids and Alkalis d. Ordinary indicator and Universal indicator 5. All alkalis are bases but all bases are not alkalis. Justify this statement. 6. Write down the effects of litmus paper, methyl orange and phenolphthalein on acid, base and salt. 7. What are hydrated salts? Give any three examples. 8. Define neutralization reaction. Explain with examples. 9. Neutralization reactions are highly applicable in our daily life. justify this statement giving any three examples. GREEN Science (Chemistry) Book-10 185

Grid-based Exercise 2 Group ‘A’ (Knowledge Type Questions) (1 Mark Each) 1. Define acid with any two examples. 2. Name the acid found in each of the given substances: i. Juice of lemon ii. Vinegar 3. Define organic acids with any two examples. 4. Mention any two physical properties of acids. 5. Define weak acid with one example. 6. Define base and give any two examples. 7. What is an alkali? Give any two examples. 8. Name any two alkalis that react with skin. (Ans: NaOH, KOH) 9. What is a salt? Give any two examples. 10. Define acidic salt with one example. 11. Define hydrated salt with one example. 12. Write any two physical properties of bases. 13. Define strong alkali with one example. 14. Name the bases which are used for given activities: i. To soften hard water ii. To make soft soap [Ans: (a) CaO (b) KOH] 15. Write down a name of alkali which is used to balance the pH of human stomach. 16. Name the bases which are used for given activities: i. To purify sugar ii. To purify petroleum products [Ans: (a) CaO (b) NaOH] Group ‘B’ (Understanding Type Questions) (2 Marks Each) 17. Write any two differences between acid and base. 18. All alkalis are bases but all bases are not alkalis, why? 19. Sodium hydroxide is called a base but sodium chloride is called a salt. Why? 20. It is dangerous to touch or taste acids, why? 21. Write any two differences between alkali and base. 22. Write any two differences between acidic radical and basic radical. 23. Explain why water can be considered as an acid as well as a base? 186 GREEN Science (Chemistry) Book-10

24. Write any two differences between neutral salt and basic salt. 25. Acetic acid is called weak acid and Sulphuric acid is called strong acid, why? 26. Sodium hydroxide is called a strong alkali, why? 27. Write any two differences between strong acid and weak acid. 28. What is meant by neutralization reaction? Give one example. Group ‘C’ (Application Type Questions) (3 Marks Each) 29. Name any three acids which are used in our daily life. Also, write an application of each. 30. Write down any three examples of neutralization reaction applied in our daily life. 31. Write any three uses of acids in our daily life. 32. Give an application of each of the given compounds. i. Calcium sulphate ii. Sodium hydroxide iii. Sulphuric acid 33. Write any three uses of salts in our daily life. 34. How are neutral salt and acid salt formed? Write with one example of each. Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each) 35. Mention any two chemical properties of alkalis with chemical equations. ‘ 36. How is salt prepared? Mention any two methods. 37. A compound gives hydrogen ion and chlorine ion in the solution state: i. Write down the name and molecular formula of the compound. ii. In what colour is methyl orange changed when it is treated with above compound ? iii. Write down the name of salt formed by the chemical reaction of above compound with zinc. Write chemical equation. 38. Write the balanced chemical equation of the reaction between strong base and weak acid, and also mention the type of salt obtained in the reaction. 39. Name the compound which gives hydrogen and sulphate ion in solution. Write down the balanced chemical equation of the chemical reaction occurred when above compound is treated with Sodium hydroxide. GREEN Science (Chemistry) Book-10 187

UNIT Some Gases 10 Weighting Distribution Theory : 5 Practical: 2 Before You Begin Air consists of different types of gases like nitrogen, oxygen, carbon dioxide, argon, neon, etc. Nitrogen occupies about 78.1% of the air by volume, oxygen gas occupies about 20.9% of the air by volume. Similarly, carbon dioxide gas occupies about 0.03% of air by volume. Green plants use carbon dioxide during photosynthesis and all animals and plants release carbon dioxide gas while breathing. When dead bodies of plants and animals decay and decompose, two gases, viz. carbon dioxide and ammonia are released in air. In this unit, we will study about laboratory preparation, properties and uses of carbon dioxide and ammonia gas. Learning Objectives Syllabus After completing the study of this unit, students will be able to: • Introduction to carbon dioxide and ammonia gases i. prepare carbon dioxide gas in the laboratory and explain its properties and uses. • Occurrence of carbon dioxide and ammonia ii. prepare ammonia gas in the laboratory and explain its properties and uses. • Laboratory preparation of carbon dioxide and ammonia • Manufacture of carbon dioxide and ammonia • Properties of carbon dioxide and ammonia • Uses of carbon dioxide and ammonia Glossary: A dictionary of scientific/technical terms photosynthesis : the process of making food by green plants in the presence of sunlight carbogen urea : the mixture of 10 – 15 % oxygen and carbon dioxide gas dry ice : a chemical fertilizer produced by heating carbon dioxide and ammonia under high pressure : the white solid form of carbon dioxide obtained after cooling carbon dioxide to – 78°C 188 GREEN Science (Chemistry) Book-10

A. Carbon dioxide Carbon dioxide is a compound gas having molecular formula CO2. It means that one molecule of carbon dioxide (CO2) is made of one atom of carbon and two atoms of oxygen. The molecular weight of carbon dioxide is 44 amu. This gas is very essential for living beings as green plants use CO2 gas to prepare food during photosynthesis. Discovery of carbon dioxide Carbon dioxide gas was discovered by Van Helmont in 1630 AD by burning wood. In 1755 AD, Joseph Black prepared this gas by burning magnesium carbonate (MgCO3). Similarly, in 1783 AD, Lavoisier proved that carbon dioxide is the compound made of carbon and oxygen. Occurrence of Carbon dioxide In nature, carbon dioxide gas is found in free as well as in combined state. In atmosphere, carbon dioxide occupies 0.03% by volume. All animals and plants release carbon dioxide in air while breathing. Some amount of carbon dioxide is found dissolved in water. So some carbon dioxide is found in water of sea, river, lake, pond, etc. When carbon compounds like wood, coal, petrol, diesel, oil, fat, wax, etc. burn, they release carbon dioxide in air. In compound state, carbon dioxide is found in mineral carbonates such as calcium carbonate (CaCO3), magnesite (MgCO3), etc. Laboratory Preparation of Carbon dioxide Principle In laboratory, carbon dioxide gas is prepared by the chemical reaction between pieces of marble or limestone (CaCO3) with dilute hydrochloric acid (HCl). CaCO3 + 2HCl CaCl2 + H2O + CO2 (dil.) Materials required Do You Know i. Apparatus Woulfe's bottle The pure form of calcium carbonate is the Thistle funnel limestone found in nature. Corks Delivery tube Gas jar ii. Chemicals Pieces of limestone or marble Dilute hydrochloric acid GREEN Science (Chemistry) Book-10 189

Procedure • Some pieces of limestone are kept in a Woulfe's bottle. • The apparatus is set as shown in the figure and dilute hydrochloric acid is poured in the Woulfe's bottle till the acid covers the pieces of limestone. • Chemical reaction takes place between dilute hydrochloric acid and limestone. Brisk effervescence can be seen during the chemical reaction. As a result, carbon dioxide gas is produced. • Carbon dioxide thus produced is collected in the gas jar by upward displacement of air as carbon dioxide gas is heavier than air. Dil. hydrochloric acid Delivery tube Thistle funnel Gas jar Fig. Pieces of CaCO3 CO2 gas 10.1 Laboratory Preparation of Carbon dioxide gas Precautions 1. The apparatus should be made air tight. Do You Know 2. Carbon dioxide gas should be collected Carbon dioxide is heavier than air and in the gas jar by upward displacement soluble in water. So, it can be collected in of air. the gas jar by upward displacement of air. 3. The lower end of the thistle funnel should be immersed in the acid. 4. The lower end of delivery tube inside the round bottom flask should not touch the acid. 5. Carbon dioxide gas dissolves in water. So it should not be collected in the gas jar by passing it through water. Test of carbon dioxide 1. When a burning piece of wood is inserted in the gas jar, it extinguishes. It shows that the gas is carbon dioxide because carbon dioxide is non-supporter of combustion. 2. When a moist blue litmus paper is kept in the gas jar containing carbon dioxide, the litmus paper turns red because carbon dioxide gas is acidic in nature. 190 GREEN Science (Chemistry) Book-10

3. When carbon dioxide is passed through clear solution of lime water, the lime water turns milky due to formation of water insoluble calcium carbonate. CO2 + Ca(OH)2 CaCO3 + H2O Water insoluble Some more methods of preparation of carbon dioxide gas 1. Carbon dioxide is prepared by burning carbon in sufficient oxygen. C(s) + O2 (g) CO2 (g) 2. When hydrocarbons like methane, ethane, propane, butane, etc. burn in oxygen, carbon dioxide is formed. CH4 (g) + 2O2 (g) CO2(g) + 2H2O (l) 2C2H6 (g) + 7O2 (g) 4CO2 (g) + 6H2O(l) 3. Carbon dioxide gas is prepared by heating calcium carbonate in a kiln. CaCO3 (s) CaO (s) + CO2 (g) 4. Carbon dioxide gas can be prepared by the reaction of carbonates or bicarbonates with acids. Na2 CO3 (s) + 2HCl (aq) 2NaCl (aq) + H2O (l) + CO2 (g) Ca (HCO3)2 (s) + 2HCl (aq) CaCl2 (s) + 2H2O (l) + CO2 (g) Mg (HCO3)2 + 2HCl (aq) MgCl2(s) + 2H2O (l) + CO2 (g) Manufacture of carbon dioxide gas In industrial scale, carbon dioxide gas is prepared by heating limestone or calcium carbonate in a kiln. When limestone is heated in a kiln at high temperature, it decomposes into calcium oxide and carbon dioxide gas. CaCO3   CaO + CO2 Carbon dioxide gas thus produced is used for industrial purpose and calcium oxide is used for white washing. When calcium oxide is mixed with water, it forms calcium hydroxide. CaO + H2O Ca(OH)2 Physical properties of carbon dioxide Do You Know 1. Carbon dioxide is colourless and The clear solution of calcium odourless gas. hydroxide is called lime water. 2. It is slightly acidic in taste when Calcium oxide is called quicklime dissolves in water. whereas calcium hydroxide is called slaked lime. 3. It is about 1.5 times heavier than air. 4. It turns moist blue litmus paper into red. GREEN Science (Chemistry) Book-10 191

5. It changes into white solid when cooled down to – 78° C, which is commonly known as dry ice. 6. It does not support combustion and extinguishes burning substances. 7. It is a non-poisonous gas. However, animals die in the atmosphere of carbon dioxide due to suffocation. Chemical properties of carbon dioxide gas 1. Carbon dioxide reacts with alkali solution (e.g. potassium hydroxide) and forms corresponding carbonate and water. CO2 + 2KOH K2 CO3 + H2O CO2 + 2 NaOH Na2 CO3 + H2O 2. When carbon dioxide gas dissolves in water, it forms carbonic acid. It is used in cold drinks to make them sour. CO2 + H2O H2CO3 3. Carbon dioxide neither burns itself nor supports combustion. But burning magnesium ribbon keeps on burning when inserted in the gas jar containing carbon dioxide and forms white solid powder (MgO) and carbon (C). 2Mg + CO2 2MgO + C 4. When carbon dioxide is passed in the clear solution of lime water for a while, the lime water turns milky due to formation of water insoluble calcium carbonate. CO2 + Ca (OH)2 CaCO3 + H2O When carbon dioxide is passed in lime water for a long time, milky colour disappears due to formation of water soluble calcium bicarbonate. CO2 + CaCO3 + H2O Ca (HCO3)2 5. Green plants use carbon dioxide gas to prepare food in leaves in the presence of the sunlight. This process is called photosynthesis. Sunlight 6CO2 + 6H2O chlorophyll C6H12O6 + 6O2 6. Carbon dioxide reacts with red hot coke at about 900°C and forms carbon monoxide. CO2 + C 900°C 2CO Activity 1 Take a beaker and prepare lime water using quick lime (Calcium oxide). Take a straw and blow air into the lime water. What do you observe? What is its reason? Discuss among friends and write down the chemical reaction involved in this process. 192 GREEN Science (Chemistry) Book-10

Uses of carbon dioxide gas Do You Know 1. Carbon dioxide is used in cold drinks Carbon dioxide gas does not extinguishes like coca-cola, beer, soda water, etc. fire itself but helps to reduce flame by displacing oxygen supply as it is a 2. Green plants used carbon dioxide gas heavier gas. for photosynthesis. 3. Liquid carbon dioxide is used in sugar mills for carbonation process. 4. It is used for manufacturing urea Do You Know (NH2CONH2) and sodium carbonate or washing soda (Na2CO3.10H2O). Carbon dioxide along with calcium hydroxide helps to remove gum, colour 5. It is used for making dry ice to preserve and amino acid impurities. meat, fish, fruits and vegetables. 6. The mixture of 95% oxygen and 5% Do You Know carbon dioxide (carbogen) is used to stimulate breathing to treat pneumonic Carbogen is also called Meduna's mixture patients. after its inventor Ladislas Meduna. 7. Carbon dioxide is used in fire extinguishers. In a fire extinguisher, sodium bicarbonate or sodium carbonate and conc. sulphuric acid are kept separately. During fire, the fire extinguisher is inverted to prepare a large amount of carbon dioxide by mixing those chemicals. The CO2 gas thus produced is used to extinguish fire. 2NaHCO3 + H2SO4 (conc.) Na2SO4 + 2H2O + 2CO2 Fig. Na2CO3 + H2SO4(conc.) Na2SO4 + H2O + CO2 B. Ammonia 10.2 Fire extinguisher Ammonia is a compound gas having molecular formula NH3. It means that one molecule of ammonia is formed by combination of one atom of nitrogen and three atoms of hydrogen. The molecular weight of ammonia is 17 amu. This gas has strong pungent odour and highly soluble in water. Discovery of ammonia Ammonia gas was disovered by Lavoisier by heating the mixture of ammonium chloride (NH4Cl) and Calcium hydroxide [Ca(OH)2]. The composition of ammonia gas was discovered by Davy and Berthecol. GREEN Science (Chemistry) Book-10 193

Occurrence of ammonia Ammonia gas is found in free as well as in the form of mixture in nature. Some amount of ammonia is found in air and soil in free state. It is formed when nitrogenous compounds decay in the absence of air (oxygen). In the form of compound, ammonia is found in Ammonium nitrate (NH4NO3), Ammonium sulphate [(NH4)2SO4], etc. Laboratory preparation of ammonia gas Principle In laboratory, ammonia gas is prepared by heating two parts of ammonium chloride (NH4Cl) and one part of calcium hydroxide [Ca(OH)2]. 2NH4Cl (s) + Ca (OH)2 (s) CaCl2 (s) + 2H2O (l) + 2NH3 (g) Materials required i. Apparatus Hard glass test tube Bunsen burner Stand with a clamp (2) Cork Delivery tube Lime tower Gas jar Match box or gas lighter ii. Chemicals required Ammonium chloride (NH4Cl) Calcium hydroxide [Ca (OH)2] Procedure • First of all, a mixture of Ammonium chloride (NH4Cl) and Calcium hydroxide [Ca (OH)2] is made in ratio 2:1 and kept in a hard glass test tube. • Apparatus is set as shown in the figure and the mixture is heated with the help of a Bunsen burner. • When the mixture of Ammonium chloride and Calcium hydroxide is heated, ammonia gas is produced which is passed through lime tower. • Dry ammonia gas is collected in the gas jar by downward displacement of air. 194 GREEN Science (Chemistry) Book-10

Mixture of NH4Cl Gas jar Stand and Ca(OH)2 Ammonia gas Moist red litmus Hard glass test tube Delivery tube paper Bunsen burner Lime tower CaO Fig. 10.3 Laboratory Preparation of Ammonia (NH3) gas Precautions i. The apparatus is made airtight. ii. The gas is collected in the gas jar by downward displacement of air because ammonia is lighter than air and highly soluble in water. iii. The hard glass test tube should be kept in inclined position facing the mouth of the test tube downwards to prevent it from cracking due to evaporation of water produced during the chemical reaction. iv. Ammonia gas should be passed through lime tower to get dry ammonia because calcium oxide absorbs moisture present in the gas. Test of ammonia gas i. When a moist red litmus paper is inserted in the gas jar containing ammonia, the litmus turns into blue because ammonia is basic in nature. NH3 + H2O NH4OH ii. Ammonia gas can be identified from its strong pungent odour. iii. White fumes of Ammonium chloride (NH4Cl) are formed when a glass rod dipped in conc. Hydrochloric acid is kept in the gas jar containing ammonia gas. NH3 + conc. HCl NH4Cl Some other methods of preparation of ammonia gas 1. Ammonia gas can be prepared by heating ammonium salts like Ammonium carbonate, Ammonium sulphate, Ammonium chloride, etc. (NH4)2 CO3 (s) H2O (l) + CO2 (g) + 2 NH3(g) (NH4)2 SO4 (s) H2O(l) + CO2 (g) + 2NH3 (g) GREEN Science (Chemistry) Book-10 195

NH4Cl (s) HCl(g) + NH3 (g) 2. Ammonia gas can also be prepared by reacting ammonium salts with strong bases/ alkalis. (NH4)2 SO4 (s) + 2NaOH (aq) Na2 SO4 (aq) + 2H2O (l) + 2 NH3 (g) NH4 Cl (s) + KOH (aq) KCl (s) + H2O (l) + NH3 (g) NH4NO3 (s) + NaOH (aq) NaNO3 (s) + H2O (l) + NH3 (g) Manufacture of ammonia gas In industrial scale, ammonia gas is prepared by heating one part nitrogen gas and three parts hydrogen gas under high temperature and pressure. This process is called Haber's process. 500°C, Fe/Mo N2 + 3H2 200 – 900 atm. 2NH3 + Heat This process is reversible. So this reaction is very slow under normal conditions. Therefore, following conditions are required to increase the rate of chemical reaction. Temperature about 500° C Pressure 200 – 900 atm. (atmospheric pressure) Catalyst Iron (Fe) Promoter Molybdenum (Mo) Do You Know The method of manufacture of ammonia was More ammonia can be prepared under discovered by German Chemist Haber. So this high pressure. However, it is dangerous to process is called Haber's Process. apply high pressure as it may explode. Physical Properties of Ammonia gas Do You Know 1. Ammonia is a colourless gas. 2. It has a strong pungent odour which A promoter is a substance that enhances may produce tears in eyes. the function of a catalyst. For example, Molybdenum. 3. It is highly soluble in water. 4. It is lighter than air. 5. It turns moist red litmus paper into blue as it is basic in nature. 6. It neither burns itself nor supports burning. 7. It solidfies at –78°C and liquifies at –33.4°C. Chemical Properties of Ammonia gas 1. Ammonia gas is highly soluble in water. It forms Ammonium hydroxide when dissolved in water. NH3 (g) + H2O (l) NH4OH (aq) 196 GREEN Science (Chemistry) Book-10

2. Ammonia reacts with acids and produces salt. 2NH3 (g) + H2SO4 (aq) (NH4)2SO4 (aq) NH3 (g) + HNO3 (aq) NH4NO3 (aq) 3. Ammonia reacts with conc. Hydrochloric acid and forms solid Ammonium chloride. NH3 (g) + conc. HCl (aq) NH4Cl (s) 4. Ammonia reacts with oxygen and forms greenish yellow flame which contains nitrogen and water. 4NH3 (g) + 3O2 (g) 6H2O (l) + 2N2 (g) 5. Ammonia reacts with carbon dioxide at about 1500° C and under certain pressure, (30 atm.), it forms urea and water. NH3 (g) + CO2 (g) 1500°C NH2 CONH2 (s) + H2O (l) Pressure 6. Ammonia solution, i.e. Ammonium hydroxide reacts with acid and forms salt and water. NH4OH + HCl (aq) NH4Cl (aq) + H2O (l) 2NH4OH (aq) + H2SO4 (aq) (NH4)2 SO4 (aq) + 2H2O (l) NH4 OH (aq) + HNO3 (aq) NH4NO3 (aq) + H2O (l) Uses of Ammonia gas 1. Liquid ammonia is used in refrigerator for cooling purpose. 2. It is used for manufacturing nitric acid, plastic, washing soda, alkalis, etc. 3. It is used to develop blue print of maps. 4. It is used as a cleansing agent to remove oil, grease, etc. 5. It is used for making chemical fertilizers like urea, ammonium sulphate, ammonium chloride, ammonium nitrate, etc. 6. It is used for manufacturing ammonium salts like NH4Cl, (NH4)2 SO4, etc. that are used in medicines. 7. It is used for making dyes, rayon, nylon, explosives, etc. 8. It is used in cold stores for cooling purpose. 9. It is used in water and waste water treatment such as pH control. 10. It is used in rubber, leather and paper industries. 11. It is used as a source of nitrogen for yeast and microorganisms in food and beverage industries. GREEN Science (Chemistry) Book-10 197

Key Concepts 1. Carbon dioxide is a compound gas having molecular formula CO2. 2. Carbon dioxide was discovered by Van Helmont in 1630 AD by burning wood. 3. In atmosphere, carbon dioxide occupies 0.03% by volume. All animals and plants release carbon dioxide in air while breathing. 4. In laboratory, carbon dioxide is prepared by the chemical reaction between pieces of limestone (CaCO3) with dilute hydrochloric acid (HCl). 5. Carbon dioxide gas should be collected in the gas jar by upward displacement of air. 6. Carbon dioxide is heavier than air. So, it can be collected in the gas jar by upward displacement of air. 7. When a burning piece of wood is inserted in the gas jar, it extinguishes. It shows that the gas is carbon dioxide because carbon dioxide is non-supporter of combustion. 8. In industrial scale, carbon dioxide gas is prepared by heating limestone or calcium carbonate in a kiln. 9. When carbon dioxide is passed in the clear solution of lime water for a while, the lime water turns milky due to formation of water insoluble calcium carbonate. 10. When carbon dioxide is passed in lime water for a long time, milky colour disappears due to formation of water soluble calcium bicarbonate. 11. Green plants use carbon dioxide gas to prepare food in leaves in the presence of the sunlight. This process is called photosynthesis. 12. Ammonia is a compound gas having molecular formula NH3. 13. Ammonia gas was discovered by Lavoisier by heating the mixture of ammonium chloride (NH4Cl) and Calcium hydroxide [Ca(OH)2]. 14. Ammonia gas is found in free as well as in the form of mixture in nature. Some amount of ammonia is found in air and soil in free sate. 15. Ammonia gas should be passed through lime tower to get dry ammonia because calcium oxide absorbs moisture present in the gas. 16. Ammonia gas can be prepared by heating ammonium salts like Ammonium carbonate, Ammonium sulphate, Ammonium chloride, etc. 17. In industrial scale, ammonia gas is prepared by heating one part nitrogen gas and three parts hydrogen gas under high temperature and pressure. This process is called Haber's process. 18. Ammonia reacts with carbon dioxide at about 1500° C and under certain pressure, (30 atm.), it forms urea and water. 19. Ammonia is used for manufacturing nitric acid, plastic, washing soda, etc. It is also used to develop blue print of maps. 198 GREEN Science (Chemistry) Book-10

Sequential General Exercise 1 1. Choose the best answer from the given alternatives. a. The molecular weight of carbon dioxide is ............... 42 amu 44 amu 17amu 6 amu b. When limestone pieces are heated in a kiln, it forms ................. calcium hydroxide and water calcium oxide and water calcium oxide carbon dioxide carbon oxide and water c. Which of the following gases is used to extinguish fire? CO2 NH3 N2 O2 d. Which of the given gases convert moist red litmus into blue? NH3 CO2 O2 H2 e. What is the ratio of ammonium chloride and calcium hydroxide to prepare ammonia gas in laboratory? 1:2 2:1 2:3 3:1 2. Answer the following questions. a. Write down the molecular formula and molecular weight of carbon dioxide. b. Where is carbon dioxide gas found in nature? c. How is carbon dioxide gas prepared in laboratory? Write with the balanced chemical equation. d. Draw a neat and labelled figure showing the laboratory preparation of carbon dioxide. e. Write any three methods of preparation of carbon dioxide gas. f. How is carbon dioxide gas prepared in industries? Describe in brief. g. Write any four physical properties of carbon dioxide. h. Write any four chemical properties of carbon dioxide with balanced formula equations. GREEN Science (Chemistry) Book-10 199

i. Write down the major uses of carbon dioxide. j. Where is ammonia gas found in nature? Write its molecular formula and molecular weight. k. How is ammonia gas prepared in laboratory? Write with the balanced chemical equation. l. How is ammonia gas prepared in industries? m. Write any three physical properties of ammonia gas. n. Write any four chemical properties of ammonia gas. o. Write down the major uses of ammonia gas. 3. Study the given figure and answer the following questions: i. Which gas is collected in the gas jar? ii. Write down the balanced chemical equation involved in this process. iii. Why is the gas jar kept erect? iv. How can you test this gas? Dil. hydrochloric acid Delivery tube Thistle funnel Gas jar Pieces of CaCO3 4. Study the given figure and answer the following questions. i. Which gas is collected in the gas jar? ii. Write down the balanced chemical equation involved in this process. iii. How can we test this gas? iv. Write down one sure test of this gas. v. What is the method of collection of this gas? 200 GREEN Science (Chemistry) Book-10