Green Science 9 Final PDF (2076)

18. Argon (Ar) 18p Atomic number = 18 22n Atomic mass = 40 8.26 Shells KL MN Fig. Argon 8× Fig. Fig. No. of e– 2 8 19. Potassium (K) 19p Atomic number = 19 20n Atomic mass = 39 8.27 Shells KL MN Potassium 81 No. of e– 2 8 20. Calcium (Ca) 20p Atomic number = 20 20n Atomic mass = 40 Shells KL MN 82 No. of e– 2 8 8.28 Calcium Electronic Configuration of Elements on the Basis shells and sub- shells Atomic Name and symbol Electronic Electronic Valency number of elements configuration on configuration on the the basis of shells basis of sub - shells 1 1. Hydrogen (H) 1××× 1s1 0 2. Helium (He) 1 3. Lithium (Li) 2××× 1s² 2 4. Beryllium (Be) 1s² 2s1 3 5. Boron (B) 21×× 4 6. Carbon (C) 1s² 2s² 22×× 1s² 2s² 2p1 23×× 1s² 2s² 2p² 24×× GREEN Science (Chemistry) Book-9 151

7. Nitrogen (N) 2 5 × × 1s² 2s² 2p³ 3 8. Oxygen (O) 2 6 × × 1s² 2s² 2p4 2 9. Fluorine (F) 2 7 × × 1s² 1s² 2p5 1 10. Neon (Ne) 2 8 × × 1s² 2s² 2p6 0 11. Sodium (Na) 2 8 1 × 1s² 2s² 2p6 3s1 1 12. Magnesium (Mg) 2 8 2 × 1s² 2s² 2p6 3s² 2 13. Aluminium (Al) 2 8 3 × 1s² 2s² 2p6 3s² 2p1 3 14. Silicon (Si) 2 8 4 × 1s² 2s² 2p6 3s² 3p² 4 15. Phosphorus (P) 2 8 5 × 1s² 2s² 2p6 3s² 3p³ 3, 5 16. Sulphur (S) 2 8 6 × 1s² 2s² 2p6 2s² 2p4 2, 6 17. Chlorine (Cl) 2 8 7 × 1s² 2s² 2p6 3s² 3p5 1 18. Argon (Ar) 2 8 8 × 1s² 2s² 2p6 3s² 3p6­ 0 19. Potassium (K) 2 8 8 1 1s² 2s² 2p6 3s² 3p6 4s1 1 20. Calcium (Ca) 2 8 8 2 1s² 2s² 2p6 3s² 3p6 4s² 2 Sub-Shells or Orbitals The main shells of atoms are K, L, M, N, O, P and Q. Each main shell consists of one or more sub-shells or orbitals which are given below: Main shells Sub-shells or orbitals K(n = 1) s L (n = 2) s, p M (n = 3) s, p, d N (n = 4) s, p, d and f Each sub-shell can accommodate a certain number of electron in it. The maximum number of electrons that can be accommodated by sub-shells or orbitals is given below: Sub-shells/orbitals Maximum number of electrons that can be accommodated in the sub-shell s2 p6 d 10 d 14 Aufbau Principle Aufbau principle states that, \"the elements in an atom are so distributed that they occupy orbits in order of their increasing energy\". Form this principle it becomes clear that the orbits having low energy level are filled faster than the orbits having high energy level as shown in the given order. 152 GREEN Science (Chemistry) Book-9

1s < 2s < 2p < 3s < 3p < 4s < 4s < 3d < 4p < 5s < 4l < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p Do You Know Aufbau principle was propounded by wolfgang Pauli and Neils Bohr. According to this principle, electrons always try to enter into the shell having low energy level. Fig. 8.29 Symbol The abbreviation of the full name of an element is called symbol. It is represented by one or more English letters. If the symbol of an element has only one letter, it is written in the capital letter. If the symbol has two letters, the first letter is written in the capital and the second in small letter. Some elements which are symbolized by only one letter are as follows: Atomic Name of element Symbol Atomic Name of element Symbol number number 1 Hydrogen H 15 Phosphorus P 5 Boron B 16 Sulphur S 7 Nitrogen N 23 Vanadium V 8 Oxygen O 53 Iodine I 9 Fluorine F 92 Uranium U Some elements which are symbolized by two letters are as follows: Atomic Name of element Symbol Atomic Name of element Symbol number number Cl 2 Helium He 17 Chlorine Ca Cr 3 Lithium Li 20 Calcium Mn Co 4 Beryllium Be 24 Chromium Zn 10 Neon Ne 25 Manganese 12 Magnesium Mg 27 Cobalt 13 Aluminium Al 30 Zinc 14 Silicon Si GREEN Science (Chemistry) Book-9 153

Some elements whose names are written in English but symbolized from Latin or German names are as follows: English name Latin/German name Symbol 1. Sodium Natrium (Latin) Na 2. Potassium Kalium (Latin) K 3. Iron Ferrum (Latin) Fe 4. Copper Cuprum (Latin) Cu 5. Silver Argentum (Latin) Ag 6. Gold Aurum (Latin) Au 7. Tungsten Wolfram (German) W Valency Valency is the combining capacity of an atom or a radical with another atom or radical to form a molecule. Valency of an atom is determined by the number of electrons present in the valence shell. Valency is the number of electrons lost, gained or shared. For example, the valency of sodium (Na) is 1 because it loses one electron from its outermost shell. The valency of magnesium (Mg) is 2 because it loses 2 electrons from its outermost shell during chemical reaction. Similarly, the valency of chlorine is 1 as it gains 1 electron from other atoms and the valency of oxygen (O) is 2 as it gains 2 electrons from other atoms during chemical reaction. If an atom has only a shell K having two electrons, the atom does not take part in chemical reaction. For example, the outermost shell of Helium (He) contains 2 electrons. So the valency of Helium (He) atom is zero. The state of an atom having two electrons in the shell K is called a duplet state. The tendency of an element to maintain two electrons in its K-shell is called duplet rule. The elements like Ne, Ar, Kr, Xn and Rn have eight electrons in their valence shell. This state is called an octate state. The state of an atom having eight electrons in the valence shell is called the octate state. The tendency of elements by which they maintain eight electrons in their valence shell is called octet rule. The outermost shell of inert gases is already fulfilled. So, these elements do not take part in chemical reaction. Hence, their valency is zero. The elements having a valency of zero are called inert gases. If elements have 1, 2 or 3 electrons in their outermost shell except shell K, their valency is 1, 2 or 3 respectively. If elements have 4 electrons in their valence shell, their valency is also 4. If elements have 5, 6 or 7 electrons in their valence shell, their valency is 3, 2 and 1 respectively. If an atom has only one shell K and the shell K has two electrons, the valency of the atom is zero. If an atom has eight electrons in the outermost shell, its valency is also zero. Some elements show variable valency. During chemical reaction,most elements lose or gain electrons from the valence shell. But in case of transition elements, electrons from the shell 154 GREEN Science (Chemistry) Book-9

just inner to valence shell also take part in chemical reaction. Due to this, these elements show variable valency. For example, S.N. Name of the elements Symbol Valency 1. Iron Fe 2 and 3 2. Copper Cu 1 and 2 3. Mercury Hg 1 and 2 4. Silver Ag 1 and 2 5. Gold Au 1 and 3 6. Lead Pb 2 and 4 Radicals Radicals are charged atoms or group of atoms having a common charge which acts as a single unit during a chemical reaction. Radicals have either a positive charge or a negative charge. So, they are chemically reactive and do not occur in free state. On the basis of electric charge, there are two types of radicals, viz. electro-positive radicals or basic radicals and electro-negative radicals or acidic radicals. 1. Electro-positive radicals or Basic radicals The atoms or groups of atoms having a positive charge are called electro-positive radicals. They are also called basic radicals. Some common electro-positive radicals with their valencies are given below: a. Electro-positive radicals having valency 1 (Monovalent radicals) i. Hydrogen (H+) ii. Lithium (Li+) iii. Sodium (Na+) iv. Potassium (K+) v. Cuprous (Cu+) vi. Mercurous (Hg+) vii. Aurous (Au+) viii. Ammonium (NH4+) b. Electro-positive radicals having valency 2 (Bivalent radicals) i. Beryllium (Be++) ii. Magnesium (Mg++) iii. Calcium (Ca++) iv. Cupric (Cu++) v. Ferrous (Fe++) vi. Zinc (Zn++) vii. Nickel (Ni++) viii. Manganous (Mn++) c. Electro-positive radicals having valency 3 (Trivalent radicals) i. Boron (B+++) ii. Aluminium (Al+++) iii. Ferric (Fe+++) iv. Auric (Au+++) GREEN Science (Chemistry) Book-9 155

d. Electro-positive radicals having valency 4 (Tetravalent radicals) i. Silicon (Si++++) ii. Plumbic (Pb++++) iii. Stannic (Sn++++) 2. Electro-negative radicals or Acidic radicals The atoms or group of atoms having negative charge/s are called electro-negative radicals. They are also called acidic radicals. Some common electro-negative radicals with their valency are given below: a. Electro-negative radicals having valency 1 i. Fluoride (F–) ii. Chloride (Cl–) iii. Bromide (Br–) iv. Iodide (I–) v. Nitrate (NO3–) vi. Nitrite (NO2–) vii. Hydroxide (OH–) viii. Bisulphate (HSO4–) ix. Bicarbonate (HCO3–) b. Electro-negative radicals having valency 2 i. Oxide (O– –) ii. Sulphate (SO4– –) iii. Carbonate (CO3– –) iv. Sulphide (S– –) v. Sulphite (SO3– –) vi. Silicate (SiO3– –) c. Electro-negative radicals having valency 3 i. Nitride (N– – –) ii. Phosphate (PO4– – –) iii. Phosphide (P– – –) iv. Phosphite (PO3– – –) Molecular Formula A molecular formula is defined as the symbolic representation of the molecule of an element or a compound in a molecular form. It represents the number and type of atoms present in a molecule. The molecular formula of Magnesium Chloride is MgCl2. It means that one molecule of Magnesium Chloride is formed by combination of one atom of magnesium and two atoms of chlorine. Similarly, the molecular formula of calcium carbonate is CaCO3. It means that one molecule of Calcium Carbonate is formed by one atom of calcium, one atom of carbon and three atoms of oxygen. We follow criss-coss method while writing molecular formulae. Steps for writing molecular formulae 1. First of all, the symbols of electro-positive and electro-negative radicals are written side by side. 2. The valency of each radical is written just below the symbol. 156 GREEN Science (Chemistry) Book-9

3. The valency of these radicals is exchanged. HCF is taken when required. 4. Radicals should be combined with exchanged valency to get the correct molecular formula. Examples: 2. Sodium chloride 1. Hydrochloric acid Na Cl 11 H Cl 11 H1 Cl1 Na1 Cl1 = HCl = NaCl 3. Magnesium oxide 4. Carbon dioxide Mg O CO 22 42 Mg2 1 O2 1 C2 1 O4 2 = MgO = CO2 5. Ammonia 6. Magnesium chloride NH Mg Cl 31 21 N1 H3 Mg1 Cl2 = NH3 = MgCl2 Information Obtained from Molecular Formula 1. Molecular formula represents one molecule of a substance. 2. The valency of each element can be calculated from molecular formula. 3. Molecular formula indicates the total number of atoms of the same/different element/s in each molecule. 4. Molecular formula indicates percentage composition of each element of a molecule. Inert gases or Noble gases Inert gases are the elements which have eight electrons in their valence shell (except helium) and do not take part in the chemical reactions. They are also called noble gases or inactive gases. Helium (He), argon, (Ar), neon (ne), etc. are the examples of inert gases. They are kept in the zero group or group 18 of the modern periodic table. They are chemically inert and occur in atomic form in gaseous state. Inert gases have duplet state or octet state in their atoms. GREEN Science (Chemistry) Book-9 157

2p+ 10p+ 18p+ 2n0 10n0 22n0 Fig. 8.30 Helium Neon Argon Atomic Name and Electronic configuration number symbol of inert Atomic state gases K L MNO P 2 Helium (He) 2 Duplet 28 Octet 10 Neon (Ne) 28 Octet 28 Octet 18 Argon (Ar) 28 8 Octet 28 18 8 Octet 36 Krypton (Kr) 18 18 8 18 32 18 8 54 Xenon (Xe) 86 Radon (Rn) Ions An atom is electrically neutral as it contains equal number of protons and electrons having opposite charges. When an atom loses electrons it becomes positively charged and when an atom gains electrons, it becomes negatively charged. These charged particles are called ions. So, ions are the charged atoms formed by gain or loss of electrons. Some examples of the ions having positive charge are Na+, K+, Mg++, Ca++ Cu++, Zn++, Al+++, etc. These ions are formed by losing one or more electrons from their valence shell. Some examples of the ions having negative charge are F–, Cl–, Br–, O– –, N– – –, etc. Chemical bond The atoms of the elements lose or gain electrons from their valence shell while taking part in a chemical reaction. Similarly, some atoms share electrons during chemical reaction. During this process, positively charged and negatively charged particles are formed and there is a force of attraction between those opposite charges, which is called chemical bond. Due to presence of chemical bond, atoms combine together in a molecule. So, chemical bond can be defined as the force between atoms in a molecule which binds them together. 158 GREEN Science (Chemistry) Book-9

Types of chemical bond There are three types of chemical bond. They are as follows: 1. Electrovalent bond or Ionic bond 2. Covalent bond 3. Co-ordinate covalent bond In this unit, we will study electrovalent bond and covalent bond only. 1. Electrovalent bond Inert gases neither gain nor lose electrons as they have octet state. Other elements either lose or gain electrons during chemical reaction to form molecules. Metals have one, two or three electrons in their valence shell. So they lose either one, two or three electrons from their valence shell. Non-metals have 5, 6 or 7 electrons in their valence shell. So, they gain 3, 2 or 1 electron from metallic atoms during chemical reaction. The chemical bond which is formed by transfer of one or more electrons from metal to non-metal is called electrovalent bond. During the formation of electrovalent bond, ions are formed. So, this bond is also called ionic bond. The compounds formed by ionic bond are called ionic compounds. These compounds are generally found in solid state. Sodium chloride (NaCl), Magnesium chloride (MgCl2), Potassium chloride (KCl), Calcium chloride (CaCl2), etc. are some examples of electrovalent or ionic compounds. To know the formation of electrovalent compounds, let us discuss the formation of sodium chloride. p+ = 11 –e– p+ = 11 n° = 12 n° = 12 Fig. Sodium (Na) atom Sodium (Na+) ion 11 proton = +11 charge 11 proton = +11 charge 11 electron = – 11 charge 11 electron = – 10 charge 8.31 Total charge = 0 Total charge = +1 Formation of Sodium ion (Na+) Sodium atom has only one electron in its valence shell. So, sodium loses one electron from its valence shell to obtain octet state like Neon. Its nucleus has 11 protons but the number of electrons become 10. So, sodium ion has one positive charge, i.e. Na+. GREEN Science (Chemistry) Book-9 159

p+ = +17 +e– p+ = +17 n° = –17 n° = –17 Chlorine atom (Cl) Chloride ion (Cl–) 17 proton = +17 charge 17 proton = +17 charge 17 electron = – 17 charge 17 electron = – 18 charge 8.32 Total charge = 0 Total charge = –1 Fig. The outermost shell of chlorine atom contains seven electrons. So, it gains one electronFig. from other atoms to obtain octet state like argon and forms chloride radical. The charge of the chloride radical (Cl–) becomes –1. When sodium atom and chlorine atom are kept in a close contact, sodium atom loses one electron from its valence shell which is taken by chlorine atom. As a result, sodium ion (Na+) and chloride ion (Cl–) are formed. There is a force of attraction between these opposite charges which results in the formation of electrovalent bond between sodium and chlorine in the molecule sodium chloride. p+ = 11 p+ = +17 n° = 12 n° = –17 8.33 Formation of sodium chloride (NaCl) Electrovalent compounds are found in the form of crystalline soilds. They have high melting and boiling point. They are brittle and soluble in water. They conduct electricity in molten or solution state. ii. Covalent bond Non-metallic atoms share electrons pairs during chemical reaction to obtain either duplet or octet state. During this process, covalent bond is formed between non-metallic atoms. The chemical bond formed by sharing of electron pairs is called covalent bond. The compounds containing covalent bond are called covalent compounds. Water, ammonia, methane, hydrochloric acid, etc. are the examples f covalent compounds. To know the formation of covalent bond and covalent compounds, let us study the formation of water molecule. 160 GREEN Science (Chemistry) Book-9

8p+ 1p+ + 8n° + 1p+ Hydrogen atom Oxygen atom Hydrogen atom 1p+ 8p+ 1p+ 8n° Fig. 8.34 Water molecule Hydrogen atom has only one electron in its valence shell. So it requires one more electron to obtain duplet state. Similarly, oxygen atom has six electrons in its valence shell. So it requires two more electrons to obtain octet state. In this condition, one atom of oxygen shares two electrons with two hydrogen atoms. As a result, water molecule is formed. Generally, covalent bond is formed between two or more non-metallic atoms. Covalent compounds exist in all three states, i.e. solid, liquid and gas. They are generally soft. They have low melting point and boiling point. They are insoluble in water and cannot conduct electricity. Differences between electrovalent compound and covalent compound Electrovalent compound 1. Covalent compound 1. It contains electrovalent bond. 2. 2. It is formed by transfer of electrons. It contains covalent bond. 3. It has high melting point and boiling 3. It is formed by sharing of electron point. pairs. 4. It is generally found in solid state. 4. It has low melting point and boiling point. It may be found in solid, liquid or gaseous state. Key Concepts 1. An element is a pure substance which cannot be divided into two or more simpler substances by ordinary chemical means. 2. All the atoms of the same element are identical but atoms of different elements are different. 3. Scientists have discovered 118 elements so far. Among them 92 are found in nature and remaining 26 elements have been discovered by chemists in the laboratory. GREEN Science (Chemistry) Book-9 161

4. The chemical substance formed by a combination of two or more elements in a fixed proportion by weight is called a compound. 5. An atom is the smallest particle of an element which can take part in a chemical reaction. 6. Among 118 elements, most atoms cannot exist freely in nature but atoms of helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn) exist freely in nature. 7. A molecule is smallest particle of an element or a compound which can exist freely in nature. 8. The mass of one proton is equal to that of one hydrogen atom. The mass of one hydrogen atom is considered as one amu (atomic mass unit). 9. Electrons are negatively-charged sub-atomic particles that revolve around the nucleus in elliptical orbits or shells. 10. An atom is electrically neutral due to the presence of equal number of protons and electrons having opposite charges. 11. An atomic number of an atom is the total number of protons present in the nucleus of that atom. 12. The sum of number of protons and neutrons present in the nucleus of an atom is called atomic mass or atomic weight. 13. Electronic configuration is the systematic distribution of electrons in various orbits or energy levels around the nucleus of an atom. 14. Aufbau principle states that, \"the elements in an atom are so distributed that they occupy orbits in order of their increasing energy\". 15. The abbreviation of the full name of an element is called symbol. It is represented by one or more English letters. 16. Valency is the combining capacity of an atom or a radical with another atom or radical to form a molecule. 17. The state of an atom having two electrons in the shell K is called a duplet state. The tendency of an element to maintain two electrons in its K-shell is called duplet rule. 18. The state of an atom having eight electrons in the valence shell is called the octate state. The tendency of elements by which they maintain eight electrons in their valence shell is called octet rule. 19. Radicals are charged atoms or group of atoms having a common charge which acts as a single unit during a chemical reaction. 20. On the basis of electric charge, there are two types of radicals, viz. electro-positive radicals or basic radicals and electro-negative radicals or acidic radicals. 21. The atoms or groups of atoms having a positive charge are called electro-positive radicals. They are also called basic radicals. 22. The atoms or group of atoms having negative charge/s are called electro-negative radicals. They are also called acidic radicals. 22. A molecular formula is defined as the symbolic representation of the molecule of an element or a compound in a molecular form. It represents the number and type of atoms present in a molecule. 162 GREEN Science (Chemistry) Book-9

23. Inert gases are the elements which have eight electrons in their valence shell (except helium) and do not take part in the chemical reactions. 24. When an atom loses electrons it becomes positively charged and when an atom gains electrons, it becomes negatively charged. 25. chemical bond can be defined as the force atoms in a molecule which binds them together. 26. The chemical bond which is formed by transfer of one or more electrons from metal to non-metal is called electrovalent bond. 27. The compounds formed by ionic bond are called ionic compounds. These compounds are generally found in solid state. Sodium chloride (NaCl), Magnesium chloride (MgCl2), Potassium chloride (KCl), Calcium chloride (CaCl2), etc. are some examples of electrovalent or ionic compounds. 28. Electrovalent compounds are found in the form of crystalline solids. They have high melting and boiling point. 29. The chemical bond formed by sharing of electron pairs is called covalent bond. The compounds containing covalent bond are called covalent compounds. Water, ammonia, methane, hydrochloric acid, etc. are the examples f covalent compounds. 30. Generally, covalent bond is formed between two or more non-metallic atoms. Covalent compounds exist in all three states, i.e. solid, liquid and gas. Sequential General Exercise 1 1. Choose the best answer from the given alternatives. a. The smallest particle of an element is called ................... matter atom molecule compound b. Which of the following is a compound? Water Hydrogen Oxygen Gold c. Which of the given elements is an inert gas? H O Ne Na d. Which of the following is the correct molecular formula of Calcium carbonate? CaCO CaCO2 CaCO3 Ca2CO3 e. Which of the following is an electrovalent compound? Water Ammonia Sodium chloride Methane GREEN Science (Chemistry) Book-9 163

2. Answer the following questions. a. What are elements? Give any three examples. b. Define compounds with any five examples. c. What is electronic configuration? Write down the electric configuration of O, Na, Mg, P, Cl and Ca on the basis of shells and sub-shells. d. What is a molecule? Give any two examples. e. What is valency? What are valence electrons? f. What are radicals? Write down their types. g. State duplet rule and octet rule. h. What are ions? Give any five examples. i. What is a chemical bond? Name its types. j. What is an electrovalent bond? k. Write any three features of electrovalent compounds with any three examples. l. What is a covalent bond? How is it formed? m. What are covalent compounds? Give any three examples. n. Write any three properties of covalent compounds. 3. Differentiate between: a. Element and Compound b. Atom and Molecule c. Duplet state and Octet state d. Electrovalent compound and Covalent compound 4. Give reason: a. Atoms are electrically neutral. b. The valency of chlorine is 1. c. The valency of neon is zero. d. Magnesium chloride is called an ionic compound. e. Water is called a covalent compound. 5. Write down the valency of the given radicals: a. Sulphate b. Chloride Chlorate c. Nitrate d. Nitrate Silicate e. Carbonate f. g. Phosphate h. 164 GREEN Science (Chemistry) Book-9

6. Write down the molecular formula of given compounds. a. Calcium sulphate b. Sodium carbonate c. Potassium chloride d. Zinc sulphate e. Potassium nitrate f. Calcium sulphate 7. Describe the formation of sodium chloride with a neat figure. 8. Describe the formation of water molecule with a neat figure. 9. Describe the formation of magnesium chloride molecule with a neat figure. 10. Describe the formation of carbon dioxide molecule with a neat figure. Grid-based Exercise 2 Group ‘A’ (Knowledge Type Questions) (1 Mark Each) 1. What is an element ? 2. What is a compound ? 3. What is electronic configuration ? 4. What is a diatomic molecule? 5. What are sub-atomic particles ? 6. What type of bond is called covalent bond? 7. What is a chemical bond? 8. Define electrovalent bond. 9. What are valence electrons? 10. What is 2n2 rule? 11. What is octet rule? 12. What is duplet state? Give one example. 13. What is molecular formula? Give one example. For Group ’B’ (Understanding Type Questions) (2 Marks Each) 14. Write any two differences between element and compound. 15. Sodium chloride is called an electrovalent compound whereas carbon dioxide is called covalent compound, why? 16. The valency of sodium is + 1 and that of chlorine is -1, why ? 17. Write any two differences between electrovalent bond and covalent bond. 18. Neon atom can exist freely in nature but sodium atom cannot, why ? GREEN Science (Chemistry) Book-9 165

19. What is the valency of calcium and hydroxide in the molecular formula Ca(OH)2 ? Which are the acidic and basic radicals in that molecular formula? 20. Why do radicals contain charge? Write. For Group ‘C’ (Application Type Questions) (3 Marks Each) 21. Calculate the maximum number of electrons present in shell L and N by using 2n2 rule. 22. What is a chemical bond? How is it formed? Write with an example. 23. How is covalent bond formed? Write in brief with are example. 24. How is electrovalent compound formed? Explain with an example. 25. How is sodium chloride formed by combination of sodium and chlorine? Describe with figure. For Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each) 26. Answer the following questions on the basis of given table: i. What is the valency of the elements A Elements Electronic configuration and D? A 2, 8, 8, 2 ii. Name the compound and type of B 2,8,2 bond formed by the combination of C 2,8,7 elements B and C. D 2,6 27. Draw the molecular structure of the following molecules: i. Sodium chloride ii. Carbon dioxide 28. In a part of the periodic table given below, some elements are denoted by symbols A, B, C,D and E. Answer the following questions on this basis. Group I A(1) I A (2) III A IV A V A VI A VII A 0 (18) Period (13) (14) (15) (16) (17) 2 Beryllium D 3B A Sulphur C Argon 4 Potassium E i. Name the elements 'B' and 'D‘ and state their valency also. ii. Name the compound formed by the combination of elements 'A' and 'C'. Also write down the type of bond present between them. 29. Name any four elements having variable valency. Also, write their valencies. 166 GREEN Science (Chemistry) Book-9

UNIT Chemical Reaction 9 Weighting Distribution Theory : 2 Practical: 0 Before You Begin When we burn a piece of paper, it changes into ash and gas. When we burn firewood, kerosene, coal, etc., we can see the formation of smoke. In these cases, one from of a matter changes into another matter having different chemical properties. Such types of change in matter is called chemical change. It occurs as a result of chemical reaction. The process of chemical change is expressed in chemical reaction. The combination, decomposition or replacement that occurs in the molecules of matter during a chemical change is called chemical reaction. We have studied general introduction to chemical reaction in class 8. In this unit, we will study of chemical reaction, word equation, balanced chemical equation methods of writing chemical equation and catalyst with its types in brief. Learning Objectives Syllabus After completing the study of this unit, students will be able to: i. introduce chemical reaction with examples • Introduction to chemical reaction ii. write chemical change in the form of word equation chemical equation. • Word equation iii. explain the method of writing chemical equation. • Reactant and products iv. introduce catalyst with its types and examples. • Chemical equation v. introduce exothermic and endothermic reaction. • Balanced chemical equation • Information obtained from a balanced chemical equation • Catalyst and its types • Exothermic and endothermic reaction Glossary: A dictionary of scientific/technical terms decomposition : Dissociation, breaking down displacement : replacement neutralization : an act or process of neutralizing, to cause a chemical to be neither an acid, nor a base catalyst : the chemical substance which can increase or decrease the rate of a chemical reaction being itself chemically unchanged exothermic : the chemical reaction that evolves heat endothermic : the chemical reaction that requires heat for the chemical change GREEN Science (Chemistry) Book-9 167

Chemical Reaction Do You Know The chemical bond present in the molecules Losing, gaining or sharing of electrons of reactants breaks due to heat, light, by an atom to gain stable electronic electricity, etc. during a chemical change. configuration is the major cause of Then combination, dissociation or exchange chemical reaction. take place between the atoms and molecules of reactants to from new substances called products. This process is called chemical reaction. The combination, decomposition or replacement or neutralization that occurs in molecules of matter during a chemical change is called a chemical reaction. Carbon + Oxygen burn Carbon dioxide C + O2 CO2 [Combination] When Copper carbonate is heated, it decomposes into Copper oxide and Carbon dioxide. Copper carbonate heat Copper oxide + Carbon dioxide CuCO3 CuO + CO2 [Decomposition] When magnesium reacts with hydrochloric acid, it forms magnesium chloride and hydrogen gas. Magnesium + hydrochloric acid Magnesium chloride + Hydrogen Mg + HCl MgCl2 + H2 [Displacement] The above examples show that a chemical change is expressed in both word equation and formula/chemical equation. Word Equation A word equation is the chemical reaction expressed by writing the full names of reactants and products. Examples: Sodium chloride Sodium + Chlorine Hydrogen + Oxygen Water Magnesium + Chlorine Magnesium chloride Calcium carbonate Calcium oxide + Carbon dioxide Zinc + Sulphuric acid Zinc sulphate + Hydrogen Potassium chlorate Potassium chloride + Oxygen Magnesium + Sulphuric acid Magnesium sulphate + Hydrogen 168 GREEN Science (Chemistry) Book-9

Formula Equation A formula equation or chemical equation is the chemical reaction expressed by writing symbols and molecular formulae of reactants and products. It gives more information than a word equation. Examples: 2NaCl Do You Know 2Na + Cl2 2H2 + O2 2H2O Formula equation gives more information Mg + Cl2 MgCl2 than a word equation. CaCO3 CaO + CO2 Zn + H2SO4 2KClO3 ZnSO4 + H2 Mg + H2SO4 2KCl + 3O2 MgSO4 + H2 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. For example CaSO4 + H2 Do You Know Ca + H2SO4 (Products) (Reactants) Reactants are written on the left hand side and products are written on the right hand 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 Mg(NO3)2 + H2 NaOH + H2SO4 Mg + HNO3 Balanced Chemical Equation The chemical equation written by balancing the total number of atoms of each element in reactants and products is called a balanced chemical equation. GREEN Science (Chemistry) Book-9 169

Examples 2H2O 2H2 + O2 ZnCl2 + H2 Zn + 2HCl Na2SO4 + 2H2O Mg(NO3)2 + H2 2NaOH + H2SO4 Mg + 2HNO3 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. Some more examples of balanced chemical equation 1. Carbon + Oxygen burn Carbon dioxide C + O2 CO2 2. Potassium + Chlorine Potassium chloride 2K+ Cl2 2KCl 3. Sodium + Chlorine Sodium chloride 2Na + Cl2 2NaCl 4. Aluminium + Nitrogen Aluminium nitride 2 Al + N2 2AlN 5. Nitrogen + Hydrogen Ammonia N2 + 3H2 2NH3 170 GREEN Science (Chemistry) Book-9

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. Water electricity Hydrogen + Oxygen 2H2O 2H2 + O2 11. Calcium carbonate Calcium oxide + Carbon dioxide CaCO3 CaO + CO2 12. Ammonium hydroxide heat Ammonia + Water NH4OH NH3 + H2O 13. Copper carbonate heat Copper oxide + Carbon dioxide CuCO3 heat CuO + CO2 14. Potassium chlorate heat Potassium chloride + Oxygen 2KClO3 heat 2KCl + 3O2 15. Hydrogen peroxide MnO2 Water + Oxygen 2H2O2 2H2O + O2 16. Silver nitrate heat Silver + Nitrogen dioxide + Oxygen 2AgNO3 heat 2Ag + 2NO2 + O2 17. Copper nitrate heat Copper + Nitrogen dioxide + Oxygen Cu (NO3)2 heat Cu + 2NO2 + O2 18. Potassium + Hydrochloric acid Potassium chloride + Hydrogen 2K + 2HCl 2KCl + H2 19. Magnesium + Hydrochloric acid Magnesium chloride + Hydrogen Mg + 2HCl MgCl2 + H2 GREEN Science (Chemistry) Book-9 171

20. Calcium + Sulphuric acid Calcium sulphate + Hydrogen Ca + H2SO4 CaSO4 + H2 21. Iron + Copper sulphate Iron sulphate + Copper Fe + CuSO4 FeSO4 + Cu 22. Magnesium + Zinc chloride Magnesium chloride + Zinc Mg + ZnCl2 MgCl2 + Zn 23. Sodium chloride + Silver nitrate Sodium nitrate + Silver chloride NaCl + AgNO3 NaNO3 + AgCl 24. Magnesium chloride + Silver nitrate Magnesium nitrate + Silver chloride MgCl2 + 2Ag NO3 Mg(NO3)2 + 2AgCl 25. Sodium hydroxide + Ferrous chloride Sodium chloride + Ferrous hydroxide 2NaOH + FeCl2 2NaCl + Fe(OH)2 26. Calcium chloride + Silver nitrate Calcium nitrate + Silver chloride CaCl2 + 2Ag NO3 Ca(NO3)2 + 2AgCl 27. Sodium sulphate + Lead nitrate Lead sulphate + Sodium nitrate Na2SO4 + Pb (NO3)2 PbSO4 + 2NaNO3 28. Hydrochloric acid + Potassium hydroxide Potassium chloride + Water HCl + KOH KCl + H2O 29. Nitric acid + Potassium hydroxide Potassium nitrate + Water HNO3 + KOH KNO3 + H2O 30. Sulphuric acid + Sodium hydroxide Sodium sulphate + Water H2SO4 + 2NaOH Na2SO4 + 2H2O 31. Sulphuric acid + Calcium oxide Calcium sulphate + Water H2SO4 + CaO CaSO4 + H2O 32. Sulphuric acid + Ferrous oxide Ferrous sulphate + Water H2SO4 + FeO FeSO4 + H2O 172 GREEN Science (Chemistry) Book-9

Limitation of a Balanced Chemical Equation A balanced chemical equation cannot provide information about: i. the physical state of reactants and products ii. conditions required for the chemical reaction iii. the rate of chemical reaction iv. the duration of chemical reaction v. concentration of reactants Modification of Chemical Equation Following modifications are done in a balanced chemical equation to make it more informative. i. The physical states of reactants and products are denoted by 's' for solid, 'l' for liquid 'g' for gas and 'aq' for aqueous solution. ii. The conditions required like heat, light, electricity, etc. are written above or below the arrow. iii. The concentration of reactants is denoted by 'dil.' for dilute and 'conc.' for concentrated solution. iv. A double ways arrow ( ) is used to represent reversible reaction and single way arrow (→) is used to represent irreversible reaction. v. Positive (+D) sign is used for exothermic reaction and negative sign (–D) is used for endothermic reaction. 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 a positive catalyst. Examples: 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). 2KClO3 2KCl + 2O2 The catalyst which decreases the rate of a chemical reaction is called a negative catalyst. GREEN Science (Chemistry) Book-9 173

Example : Glycerol [C3H5(OH)3] acts as a negative catalyst during decomposition of Hydrogen peroxide (H2O2). 2H2O2 Glycerol 2H2O + O2 Characteristics of a 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 decreases 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. 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: C + O2 CO2 + Heat C + 2H2 CH4 + Heat Zn + 2HCl ZnCl2 + H2 + Heat CaO + H2O Ca(OH)2 + Heat CH4 + 2O2 CO2 + 2H2O + Heat 174 GREEN Science (Chemistry) Book-9

Reversible Reaction The chemical reaction which occurs in forward as well as backward directions is called a reversible reaction. In this reaction, the products are changed back into reactants again, e.g. H2 + O2 H2O N2 + 3H2 2NH3 Irreversible Reaction The chemical reaction which occurs only in one direction is called irreversible reaction. In this reaction, we cannot obtain reactants from the products again, e.g. CaCO3 CaCO3 Na + Cl2 2NaCl Key Concepts 1. Losing, gaining or sharing of electrons by an atom to gain stable electronic configuration is the major cause of chemical reaction. 2. The combination, decomposition, replacement or neutralization that occurs in molecules of matter during a chemical change is called a chemical reaction. 3. A word equation is the chemical reaction expressed by writing the full names of reactants and products. 4. A formula equation or chemical equation is the chemical reaction expressed by writing symbols and molecular formulae of reactants and products. 5. Reactants are chemical substances which take part in a chemical reaction. They are written on the left hand side of the arrow. 6. Products are the chemical substances which are produced after a chemical reaction. They are written on the right hand side of the arrow. 7. 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. 8. The chemical equation written by balancing the total number of atoms of each element in reactants and products is called a balanced chemical equation. 9. The chemical substance which is used to increase or decrease the rate of a chemical reaction is called a catalyst. 10. The catalyst which increases the rate of a chemical reaction is called positive catalyst. 11. The catalyst which decreases the rate of a chemical reaction is called a negative catalyst. GREEN Science (Chemistry) Book-9 175

12. A catalyst does not initiate a chemical reaction but increases or decease the rate of chemical reaction. 13. The chemical reaction that absorbs heat during the chemical change is called endothermic reaction. 14. 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. A chemical change is a ....................... change. reversible irreversible physical molecular b. The substances which take part in a chemical reaction are called ....................... products reactants atoms molecules c. The substances which are formed after a ....................... change are called products. chemical physical atomic molecular d. Which of the following is a catalyst? MnO2 H2SO4 CaCO3 MgCO3 e. Which of the following is an endothermic reaction? CaCO3 CaO + CO2 N2 + O2 2NO Zn + 2HCl ZnCl2 + H2 C + O2 CO2 176 GREEN Science (Chemistry) Book-9

2. Answer the following questions. a. Define chemical reaction with any three examples. b. What are reactants and products? Write with an example. c. What is a word equation? Give any three examples. d. What is a chemical equation? Write any three examples. e. What is a catalyst? Write its types. f. Define positive and negative catalyst with any one example of each. g. 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. Positive catalyst and Negative catalyst d. Endothermic reaction and Exothermic reaction 4. 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 5. Write down the given word equations into balanced chemical equations. a. Nitrogen + Hydrogen Ammonia b. Hydrogen + Oxygen Water c. Aluminum + Nitrogen Aluminium nitride 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 GREEN Science (Chemistry) Book-9 177

h. Copper + Oxygen Copper oxide i. Potassium chlorate Potassium chloride + Oxygen j. Nitric acid + Calcium hydroxide Calcium nitrate + Water 6. Write down the method of writing balanced chemical equation with examples. 7. Write any four pieces of information that can be obtained from a balanced chemical equation. Grid-based Exercise 2 Group ‘A’ (Knowledge Type Questions) (1 Mark Each) 1. What is a physical change ? 2. What is a chemical reaction? 3. What is a chemical equation ? 4. What is a word equation? Give one example. 5. What is a chemical (or formula) equation ? 6. What is a balanced chemical equation? 7. Write any two limitations of a balanced chemical equation. 8. What is a catalyst? What are its types? Write 9. What is a chemical change ? 10. What is a catalyst? 11. Write any two characteristics of catalyst. 12. Define 'reactants'. 13. What type of catalyst is called a positive catalyst? For Group ’B’ (Understanding Type Questions) (2 Marks Each) 14. The formation of ice from water is called a physical change, why? 15. The formation of ash by burning a paper is called a chemical change, why? 16. Why is formula equation more meaningful than a word equation? 17. Manganese dioxide is called a positive catalyst and glycerol is called a negative catalyst, why ? 18. Write any two limitations of the given chemical equation. AgBr Ag + Br2 19. Write any two differences between physical change and chemical change. 20. Write any two pieces of information which can be obtained from given chemical equation. A 178 GREEN Science (Chemistry) Book-9

For Group ‘C’ (Application Type Questions) (3 Marks Each) 21. What do you mean by 'reactants'? Write with examples. Write down the balanced chemical equation of the reaction between sulphuric acid and ammonium hydroxide. 22. Write down the information which can be obtained from a balanced chemical equation. Write a balanced equation of the chemical reaction between Iron (III) chloride and Ammonium hydroxide. 23. What type of catalyst is called a positive catalyst? Give an example. Write any two applications of catalysts. 24. 'What are the limitations of a balanced chemical equation? Write down the balanced chemical equation for the following reaction. Calcium chloride + Silver nitrate Calcium nitrate + Silver chloride 25. We can get more information from a formula equation than from a word equation. Justify this statement. Write down the balanced chemical equation of the given word equation. Aluminium + Oxygen Aluminium oxide For Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each) 26. Correct and rewrite the given chemical equations: i. NaCO3 + CaCI2 NaCI2 + CaCO3 ii. HNO2 + NaHCO3 NaNO3 + H2O + CaCO2 iii. NaOH + NH3CI NaCI + NH2 + H2O iv CaCI2 + AgNO3 CaCO3 + AgCl 27. Study the given word equation and answer the following equations: Calcium carbonate Heat Calcium oxide + Carbon dioxide i. Express the above word equation in the form of formula equation. ii. Write any two information that can be obtained from the given equation. iii. Name the factor that brings out chemical change in the above reaction. 28. Change the given word equations into balanced chemical equation. Differentiate between reactants and products. i. Sulphuric acid + Zinc Zinc sulphate + Hydrogen ii. Iron + Oxygen Ferric oxide 29. Complete and balance the given chemical equations: i. Na2CO3 + ..... NaCI + ..... ii. HCI + ... CaCl2 + H2O 30. What happens in the given conditions ? Write with balanced chemical equation : i. When magnesium reacts with nitric acid ii. When silver nitrate reacts with sodium chloride GREEN Science (Chemistry) Book-9 179

UNIT Solubility 10 Weighting Distribution Theory : 5 Practical: 3 Before You Begin Take a beaker with some water. Add a spoonful of sugar and stir it with a glass rod. What happens? The sugar dissolve in water. It is called sugar solution. When we add salt in water, the salt dissolves. It is called salt solution. Sugar solution and salt solution both are homogeneous mixtures. So solution can be defined as a homogeneous mixture of two or more substances. A solution is made of solution and solvent. In a solution, the dissolved component, i.e. solute cannot be seen with our naked eyes. Solution is very important in our daily life. In this unit, we will study solution, types of solution, solubility, solubility curve and crystallization in brief. Learning Objectives Syllabus After completing the study of this unit, students will be able to: • Mixture and its types i. introduce solution and state its types with examples. • Solution, solute and solvent • Types of solution ii. define unsaturated, saturated and super saturated • Solubility solution and prepare such solutions. • Solubility curve • Crystalline solids and iii. define solubility of a substance and explain the relation between solubility and temperature. amorphous solids. • Crystallization iv. describe crystallization in brief. Glossary: A dictionary of scientific/technical terms solution : the homogeneous mixture made of solute and solvent solute : the substance which dissolves in a solvent solvent : the substance in which a solute dissolves crystalline : made up of repeating crystals crystal : a solid substance having a fixed geometrical shape, smooth surface and sharp edges 180 GREEN Science (Chemistry) Book-9

Solution A solution is the homogeneous mixture of two Fig. Sugar + water or more substances. All the components of a Sugar solution solution cannot be seen through our naked eyes. Solution of sugar and water, salt and water, alcohol and water, coke and water, milk and water, etc. are some examples of solution. A solution consists of two components. They 10.1 are solute and solvent. Solute + solvent = Solution A solute is a substance which gets dissolved in a solvent. For example, sugar, salt, copper sulphate, magnesium chloride, etc. A solvent is a substance which dissolves a solute into it. For example, water, alcohol, ether, etc. In the solution of salt and water, salt is a solute and water is a solvent. Green plants absorb minerals from soil in the form of solution. Similarly, animals get nutrients in the form of solution. Therefore, solution plays a great role in the growth and development of living beings. Experiment: 1 To prepare the mixture of various substances with water Materials required Sugar, salt, washing soda, soil, sand, cooking oil, chalk, stone, wood, beaker, water, glass rod Procedure • Take a beaker. Add some water into it. • Now, add one teaspoonful of sugar into water and stir it with a glass rod. Observe whether sugar dissolves in water or not. Fill in the given table after your observation. • By the same way prepare a mixture of above given substances one by one. • Observe whether these substances dissolve in water or not. • Fill in the given table after your observation. S.N. Mixtures Homogeneous mixture Heterogeneous mixture 1 water + sugar √√ 2 water + sand 3 4 GREEN Science (Chemistry) Book-9 181

5 6 7 8 9 10 There are different types of solution on the basis of states of the components of the solution. Some common types of solutions are tabulated below: Types of solution Examples Solution Solvent 1. Solid + liquid Salt water salt water 2. Liquid+ liquid Ink and water ink water 3. Liquid + gas Oxygen dissolved in water oxygen water 1. Gas + gas Atmosphere other gases nitrogen From above table, it becomes clear that the solution of solid and gas is not found in nature. Dilute solution and concentrated solutionFig. Dilute solution is the solution which Concentrated solution contains relatively less amount of solute. Similarly, concentrated solution is the solution which contains relatively more amount of solute. Dilute solution is lighter than the concentrated solution. It means that the density of concentrated solution is more than that of dilute solution. 10.2 Dilute solution Differences between Dilute solution and Concentrated solution Dilute solution Concentrated solution 1. It has relatively less amount of a solute 1. It has relatively more amount of solute in a certain amount of solvent. in a certain amount of solvent. 2. It has less density. 2. It has more density. Activity 1 Take two beakers and keep 100 ml of water in each. Mark the beakers A and B. Add 5 gram of copper sulphate in the beaker A and stir it with a glass rod. Add 50 gram of copper sulphate in the beaker B and stir it with a glass rod. 182 GREEN Science (Chemistry) Book-9

The concentration of copper sulphate is more in the beaker B than in beaker A. The solution in beaker A is dilute and that in beaker B is concentrated. Activity 2 Take two glasses and put about 100 ml pure drinking water in each. Mark one glass A and another glass B. Dissolve 10 gram of table salt in the glass A and 50 gram of salt in glass B. Taste the salt solution from the beaker A. Rinse your mouth. Then taste the same amount of salt solution from the glass B. Which solution is more salty A or B? Which solution is dilute and which one is concentrated? Why? Unsaturated, Saturated and Supersaturated solution Take a beaker and put 50 ml water in it. Put a b c 5 gram of salt in it and stir with a glass rod. Now, add 5 gram more salt and stir till the salt dissolves completely. This type of solution Fig. is called unsaturated solution. The solution 10.3 Unsaturated Saturated Super saturated which can dissolve some more solute at a certain temperature is called unsaturated solution. Keep on adding some more salt, the salt gets dissolved and finally some more amount of salt does not dissolve at that temperature. This type of solution is called saturated solution. The solution in which no more solute can be dissolved at a fixed temperature is called saturated solution. But more amount of solute can be dissolved by raising the temperature of the solution. We can dissolve further amount of salt by raising its temperature. But the solution releases excess solute in the form of crystals when the solution is cooled. This type of solution is called supersaturated solution. The saturated solution prepared at a higher temperature which throws excess solute as a solid when cooled is called supersaturated solution. The density of unsaturated solution is minimum and that of supersaturated solution is maximum. We can identify the type of solution by using a crystal of the solute from which the solution is made. When a crystal of the solute dissolve in the solution at a fixed temperature, the type of solution is unsaturated. If the crystal does not dissolve in the solution at the given temperature, the type of solution is saturated. If the size of the crystal increases gradually, the type of the solution is supersaturated. GREEN Science (Chemistry) Book-9 183

Differences between Unsaturated solution and Saturated solution Unsaturated solution Saturated solution 1. More amount of solute can dissolve at 1. No more amount of solute can dissolve a certain temperature. at a certain temperature. 2. It has less density. 2. It has more density. 3. It has less saturation. 3. It has more saturation. Differences between Saturated solution and Supersaturated solution Saturated solution Supersaturated solution 1. No more amount of solute can dissolve 1. It is the saturated solution prepared at at a certain temperature in this type of higher temperature. solution. 2. It has less density. 2. It has more density. 3. It does not change concentration when 3. Its concentration decreases and crystals it comes in contact of a solute. are formed when it comes in contact of solute. Utilities of solution in Our Daily Life Solution is very important in our daily life. The major utilities of solution are given below: i. We use many medicines in the form of solution. ii. We consume many drinks in the form of solution. iii. The food that we eat is digested and absorbed in the form of solution. iv. Green plants absorb water and minerals in the form of solution. v. Aquatic plants and animals take oxygen in the form of solution. vi. We use ink, paint, etc. in the form of solution. Solubility In order to get the concept of solubility, let's perform the following activity. • Take a beaker and add 100 gram water (100 ml water) into it. Measure the temperature of water. Let the temperature of water be 20°C. • Now keep the beaker on a top pan balance. Note down the weight of the beaker along with water. • Take a spoon and dissolve some sugar in the water. Keep on dissolving sugar until it stops to dissolve. The amount sugar dissolved in 100 gram of water in this condition is taken as the solubility of sugar. It will be approximately 179. It means that 179 gram of sugar dissolves in 100 gram of water of 20°C and forms a saturated solution. Solubility of a solute can be defined as the amount of solute (in gram) dissolved in 100 gram of a solvent at a certain temperature to form a saturated solution. 184 GREEN Science (Chemistry) Book-9

In other words, the solubility of a solute at a fixed temperature can be defined as the amount of solute required to form a saturated solution in 100 gram of water at that temperature. Formula: Weight of solute (in gram) Solubility = Weight of solvent (in gram) × 100 Solution has no unit as it is a percentage ratio of two masses (weights). At 20°C, the solubility of sodium chloride is 35.7. It means that 35.7 gram of sodium chloride is required by 100 gram of water at 20°C to form a saturated solution. Different solutes have different solubility at a fixed temperature. Similarly, same solute has different solubility at different temperature. Therefore, temperature is mentioned while writing solubility of a substance. The solubility of some common solutes at 20°C temperature is given below. S.N. Solute Solubility Temperature 1 Sugar 179 20°C 2 Copper sulphate 20 3 Common salt 35.7 4 Sodium nitrate 88 5 Magnesium chloride 32 Worked out Numerical: 1 At 25°C, 65 gram of a solute dissolves in 75 gram of water. Calculate the solubility of the salt. Solution Given, At 25°C, Weight of solute = 65 gram Weight of solvent = 75 gram Solubility = ? We know, wt. of solute Solubility = wt. of solvent× 100 65 = 86.66 = 75 × 100 \\ At 25°C, the solubility of the solute = 86.66 GREEN Science (Chemistry) Book-9 185

Worked out Numerical: 2 When 50 gram of a saturated solution prepared at 60°C is dried, 20 gram of a salt is left behind. Calculate the solubility of the salt at the given temperature. Solution, Given, At 60°C, Weight of solution = 50 gram Weight of solute = 20 gram Then weight of solvent = (50 – 20) gram = 30 gram Solubility = ? We know, wt. of solute Solubility = wt. of solvent × 100 20 Solubility = 30 × 100 = 66.66 \\ The solubility of the solute = 66.66 at 60°C. Relation between solubility and temperature The solubility of most solutes increases on increasing the temperature of the solution. The reverse of this is also true. When a solvent is heated, the intermolecular space of the solvent increases. As a result, more solute dissolves in a solvent at high temperature. At different temperature same mass of a solvent can dissolve different amount of a solute. Solutes 20°C Solubility at 100°C 1. Copper sulphate 20 76.99 2. Common salt 35.7 39.1 3. Sugar 179 487 4. Ammonia gas 89.9 7.4 5. Oxygen gas 0.0045 0.0033 From above table, it becomes clear that the solubility of solid substances increases on increasing the temperature. But the solubility of gases decreases on increasing the temperature of the solution. In summer season, the amount oxygen dissolve in water decreases. So fishes come on the surface for breathing. 186 GREEN Science (Chemistry) Book-9

Liquids expand more than solids and gases expand more than liquids. So the solubility of gases decreases on increasing the temperature of solvent due to expansion of gases. Solubility curveFig. Solubility curve is defined as the curve obtained by plotting the solubility of a solute at different temperatures along Y-axis and corresponding temperature along X-axis. Let us observe the given graph having solubility curves of different solutes at different temperatures. Above graph shows the solubility of different solutes at different 10.4 temperature. These curves also show that the solubility of different substances at a fixed temperature is also different. Application of solubility curves 1. The solubility of a solute at a fixed temperature can be known. 2. The solubility of different solutes at a fixed temperature can be known. 3. The solubility of different substances at a certain temperature can be compared. 4. The relation between increase in temperature and change in solubility can be observed. 5. The amount of solute that crystallizes out while cooling a solution to a certain temperature can be calculated. 6. We can predict the solute that crystallizes first among many solutes. Activity 4 Take a graph paper and plot a solubility curve on the basis of given table. Also, find out the relation between solubility and increase in temperature. Temperature (°C) 0 10 20 30 40 50 60 70 80 90 100 Solubility 14 17 20 24 29 34 40 48 57 68 77 GREEN Science (Chemistry) Book-9 187

Crystallization When a hot supersaturated solution is cooled, it throws excess solute as solid. As a result, crystals are formed. Crystals are the solid substances having a fixed geometrical shape, smooth surfaces and sharp edges. Crystals have a fixed melting point. The process of getting crystals by cooling down a hot supersaturated solution is called crystallization. This process is applied to purify solid substances. Fig. 10.5 Crystals of sugar Crystals of copper sulphate Crystallization of a certain solid substance depends on the solubility of the substance. The substances whose solubility is less, crystallize faster than the solute having more solubility, for example, at 20°C, the solubility of copper sulphate is 20 and that of sodium chloride is 35.7. So copper sulphate crystallizes earlier than sodium chloride when their saturated solution is cooled. We should cool down the solution for a long time to obtain the crystals of sodium chloride. Activity 5 • Take a beaker and keep 50ml of water into it. • Add some copper sulphate and stir it with a glass rod. • Dissolve some more copper sulphate in the water. • Now heat the solution and add more copper sulphate to form a super-saturated solution. Heat the solution till you see fine crystals of copper sulphate on the wall of beaker. • Now, take a small crystal of copper sulphate and tie it with a thread. • Immerse the crystal into the solution and leave the solution undisturbed for 2-3 days. • Observe the beaker after 3 days. You will see large crystals of copper sulphate. This process is called crystallization. 188 GREEN Science (Chemistry) Book-9

China dish Copper China Copper Wire gauze sulphate dish sulphate solution crystals Boiling water Fig. Burner Tripod Tripod stand stand 10.6 Key Concepts 1. A solution is the homogeneous mixture of two or more substances. 2. A solution consists of two components. They are solute and solvent. 3. A solute is a substance which gets dissolved in a solvent. For example, sugar, salt, copper sulphate, magnesium chloride, etc. 4. A solvent is a substance which dissolves a solute into it. For example, water, alcohol, ether, etc. 5. Dilute solution is the solution which contains relatively less amount of solute. Similarly, concentrated solution is the solution which contains relatively more amount of solute. 6. The solution which can dissolve some more solute at a certain temperature is called unsaturated solution. 7. The solution in which no more solute can be dissolved at a fixed temperature is called saturated solution. 8. The saturated solution prepared at a higher temperature which throws excess solute as a solid when cooled is called supersaturated solution. 9. We can identify the type of solution by using a crystal of the solute from which the solution is made. 10. Solubility of a solute can be defined as the amount of solute (in gram) dissolved in 100 gram of a solvent at a certain temperature to form a saturated solution. 11. At 20°C, the solubility of sodium chloride is 35.7. It means that 35.7 gram of sodium chloride is required by 100 gram of water at 20°C to form a saturated solution. 12. Different solutes have different solubility at a fixed temperature. Similarly, same solute has different solubility at different temperature. Therefore, temperature is mentioned while writing solubility of a substance. GREEN Science (Chemistry) Book-9 189

13. The solubility of most solutes increases on increasing the temperature of the solution. The reverse of this is also true. 14. Liquids expand more than solids and gases expand more than liquids. So the solubility of gases decreases on increasing the temperature of solvent due to expansion of gases. 15. Solubility curve is defined as the curve obtained by plotting the solubility of a solute at different temperatures along Y-axis and corresponding temperature along X-axis. 16. Crystals are the solid substances having a fixed geometrical shape, smooth surfaces and sharp edges. 17. The process of getting crystals by cooling down a hot supersaturated solution is called crystallization. This process is applied to purify solid substances. Sequential General Exercise 1 Water 1. Choose the best answer from the given alternatives. a. Which of the given substances is a solvent? Sugar Salt Copper sulphate b. Which of the following is a solute? Water Alcohol Ether Common salt c. The solution which cannot dissolve any more solute in a fixed temperature is called................... Homogeneous mixture Saturated solution Unsaturated solution Supersaturated solution d. At 25°C, 25g of a salt dissolves into 50 g of water. What is the solubility of the salt? 25 50 75 75 190 GREEN Science (Chemistry) Book-9

e. Which of the following contains maximum solute? Concentrated solution Dilute solution Saturated solution Supersaturated solution 2. Answer the following questions. a. What is a solution? Give any two examples. b. Define solute and solvent with any two examples of each. c. What is unsaturated solution? d. What is saturated solution? e. What is supersaturated solution ? How is it prepared? f. Define dilute and concentrated solution. g. What is solubility? h. Write down the formula to calculate solubility. i. What is solubility curve? j. What are crystals? k. What is crystallization? 3. Write down the relationship between change in temperature and solubility. 4. Write any three applications of solubility curve. 5. Differentiate between. a. Solute and Solvent b. Dilute and Concentrated solution c. Saturated solution and Unsaturated solution d. Unsaturated solution and Supersaturated solution 6. Give reason. a. The solubility increases on increasing the temperature. b. Solubility has no unit. c. Temperature is mentioned while writing solubility of a solute. 7. Write down the method of obtaining crystals of copper sulphate. 8. Study the give table and answer the following questions. Temperature 20°C 40°C 60°C 80°C 100°C Solubility of PbNO3 58 74 94 115 140 a. What is the relation between increase in temperature and solubility of lead nitrate? b. Plot the solubility curve on the basis of above data. c. What is the solubility of lead nitrate at 30°C and 50°C? d. How much gram of lead nitrate crystallizes out when the solution prepared at 80°C is cooled to 40°C? GREEN Science (Chemistry) Book-9 191

9. Numerical Problems a. When 6 gram of ammonium nitrate is dissolved in 25 gram of water and heated upto 30°C, a saturated solution is formed. Calculate the solubility of ammonium nitrate. [Ans : 24] b. A saturated solution is formed by keeping 48 gram of a salt in 200 gram of water at 25°C. Calculate the solubility of the salt. [Ans: 24] c. The solubility of potassium nitrate is 40 at 30°C. Calculate the mass of water required to form a saturated solution by dissolving 15 gram of potassium nitrate at that temperature. [Ans: 37.5g] d. The solubility of a salt at 60°C is 55. Calculate the amount of salt dissolved in 40 gram saturated solution of the salt at that temperature. [Ans: 14.2g] e. How much amount of crystal is formed when 15 gram of a saturated solution prepared at 30°C is cooled to 10°C. The solubility of that salt at 30°C is 95 and at 10° C is 30. [Ans : 3.85g] Grid-based Exercise 2 Group ‘A’ (Knowledge Type Questions) (1 Mark Each) 1. What is a mixture ? 2. What is heterogeneous mixture? 3. Define homogeneous mixture with one example. 4. Define solvent. 5. Define unsaturated solution. 6. What is saturated solution? 7. What is supersaturated solution? 8. What is solubility? 9. What is a solubility curve? 10. Define dilute solution and concentrated solution. 11. What is a crystal? 12. Write down the formula to calculate the solubility. 13. What happens when a saturated solution prepared at 40°C is cooled down to 20°C? For Group ’B’ (Understanding Type Questions) (2 Marks Each) 14. Temperature is also mentioned while writing the solubility of a substance, why? If concentration of a solution reduces and crystals appear when solute is added to the solution at 40°C, name the type of solution. 15. Why is turbid water called a heterogeneous mixture ? Give reason. How is a saturated solution made from an unsaturated solution without adding the solute ? 16. Solute separates as a solid when a saturated solution is cooled. Give reason. What does it mean by the statement that the solubility of copper sulphate at 20°C is 21? 17. Write any two differences between unsaturated and saturated solution. 192 GREEN Science (Chemistry) Book-9

18. Why does a solute dissolve faster in a solvent having high temperature? Give reason. What happens when supersaturated solution of sodium chloride is cooled? 19. Write any two differences between solvent and solute. 20. When a saturated solution at a high temperature is cooled, why does the solute separate out as a solid? For Group ‘C’ (Application Type Questions) (3 Marks Each) 21. Write any three utilities of solubility curve? 22. Describe in brief the utility of solution in our daily life. What happens when the saturated solution of copper sulphate is heated? 23. How is the solubility of a solute calculated ? Explain. 24. Describe in brief the method of finding the solubility of common salt. 25. What happens when a crystal of copper sulphate is added separately in unsaturated, saturated and supersaturated solution of copper sulphate? For Group ‘D’ (Higher Abilities Type Questions) (4 Marks Each) 26. At 20°C, the solubility of sodium chloride is 35. How much gram sodium chloride is required at that temperature to make a saturated solution in 20 gram of water? More amount of solute dissolves in hot water than in cold water, why? 27. How much solute separates as a solid when 30 gram of a saturated solution prepared at 50°C is cooled upto 30°C? The solubility of the given solute is 35 at 50°C and 21 at 30°C. Less amount of gas dissolves in hot water, why? 28. Describe the method of making crystals of copper sulphate with a labelled figure. 29. The solubility of a salt is 36 at 35°C. How much salt is required to make a saturated solution in 144 gram of water at that temperature ? Differentiate between solubility and solubility curve. 30. Study the given solubility curve and answer the Solubility 65 following questions: 60 55 i. What is the solubility of sodium chloride at 50 25°C? 45 40 ii. What is the relationship between increase 35 in temperature and solubility? 30 25 iii. What happens when a saturated solution of 20 sodium chloride at 50°C is cooled to 20°C ? 15 10 iv. What is the solubility of Ammonium 5 chloride at 24°C and 65°C from the solubility curve? 0 10 20 30 40 50 60 70 80 90 100 Temperature°C GREEN Science (Chemistry) Book-9 193

UNIT Some Gases 11 Weighting Distribution Theory : 5 Practical: 3 Before You Begin Air is the mixture of different gases. The gases present in air are nitrogen (78.08%), oxygen (20.95%), argon (0.93%), carbon dioxide (0.03%) hydrogen (0.00005%), ozone (0.000004%) and other gases (Ne, He, CH4, H2O). Various natural as well as artificial substances are made of gases. The formation of plants, animals, fossil fuels, food and drinks is impossible in the absence of gases. Our food as well as fossil fuels require gases directly or indirectly. About 92.2% of human body is made of four elements, viz. nitrogen, oxygen, hydrogen and carbon. Among four elements, three are gases. These gases combine with other elements and form a variety of substance. These gases play a great role in the existence of living beings. In this unit, we will study hydrogen, oxygen and nitrogen gases with their laboratory preparation, properties and uses. Learning Objectives Syllabus After completing the study of this unit, students will be able to: • Laboratory preparation i. introduce some gases like hydrogen, oxygen and of hydrogen, oxygen and nitrogen gases nitrogen. • Physical properties, chemical ii. prepare hydrogen, oxygen and nitrogen gases in properties and uses of laboratory. hydrogen, oxygen and nitrogen gases iii. explain physical properties, chemical properties and uses of hydrogen, oxygen and nitrogen gases. • Industrial preparation of hydrogen, oxygen and nitrogen gases Glossary: A dictionary of scientific/technical terms combustion : the process of burning of a substance nascent : highly reactive and newly formed apparatus : the tools or other equipment that are needed for a particular activity welding : the process of joining metals 194 GREEN Science (Chemistry) Book-9

1. Hydrogen gas Hydrogen is the lightest and most abundant element of the solar system. This gas is found in the sun and stars. But in earth hydrogen gas is found only about 0.00005%. Since, hydrogen is very reactive gas so it is found interms of compound by reacting with other elements. For example, acid, hydrocarbon, carbohydrate, etc. Henry Cavendish in 1823 named this gas as \"inflammable air\" but later Lavoiser named it as \"hydrogen\" as it forms water while burning. The meaning of hydrogen is \"water producer\". Some facts about hydrogen are tabulated below: Symbol Molecular Atomic Atomic Electronic H formula number weight configuration H2 1 1 1s1 Methods of Preparation of Hydrogen Gas There are mainly three sources of hydrogen. They are acid, base and water. Hydrogen gas can be prepared from these three sources. Here, we will discuss the preparation of hydrogen gas in laboratory and from the electrolysis of water. a. Laboratory Preparation of Hydrogen Do You Know Gas Use of pure zinc decreases the rate of Active metals like zinc, magnesium, etc. reaction. 50 impure zinc is used to prepare zinc in laboratory. The impurity present in replace the hydrogen present in acid and hence it helps to increase the rate of reaction. hydrogen gas is formed. Hydrogen gas can be prepared in the laboratory by the reaction of impure zinc and dilute hydrochloric acid. Zn (s) + 2HCl (aq) Zn Cl2 (aq) + H2 Thistle Dil HCl Delivery tube Gas jar funnel Cork H2 Beehive shelf Fig. Woulf's Water bottle Granulated 11.1 Zinc Laboratory Preparation of Hydrogen Gas Procedure Few pieces of zinc are taken in woulfe's bottle and the apparatus is set as shown in the figure. Then dilute hydrochloric acid is poured gently through the thistle funnel fitted GREEN Science (Chemistry) Book-9 195

in one opening of the bottle so that the zinc pieces immerse in the acid. Then chemical reaction takes place between zinc and hydrochloric acid and hydrogen gas is evolved. Precautions i. The apparatus should be clean. ii. Impure Zinc should be used. iii. The apparatus should be made airtight. iv. Dilute hydrochloric acid should be used because concentrated acid does not produce hydrogen gas. v. The stem end of the thistle funnel must be dipped into acid. vi. The gas jar should be full of water, i.e. there should not be presence of air bubbles. Test of hydrogen gas How we know the gas prepared in laboratory is hydrogen gas? Think yourself. For this purpose, keep the gas jar vertical slowly and take a burning match stick near the mouth of the gas jar. If the gas burns with 'pop' sound, it proves that the gas is hydrogen gas as other gases do not sound 'pop'. b. Preparation of Hydrogen Gas From Electrolysis of Water In industrial scale, hydrogen gas can be prepared from water. Water is found in abundant form in the earth and presence of two molecules of hydrogen in water is the reason of production of hydrogen gas for commercial purpose. Oxygen Hydrogen gas gas Oxygen Water External bubbles with source emf soluble Hydrogen salt bubbles – Battery + Fig. 11.2 Anode Cathode Production of hydrogen gas by electrolysis method Water mixed with a little quantity of sulphuric acid is kept in a electrolytic cell with passage of electricity which causes the decomposition of water into hydrogen ions and oxygen ions. Hydrogen ions move towards the cathode and oxygen ions move towards the anode and collected in the form of gas. Collected gas is passed through pipe and filled in the cylinder. 196 GREEN Science (Chemistry) Book-9

2H2 O 2H2 + O2 We can prepare 100% pure hydrogen gas by this method, where electricity is cheap, hydrogen gas can be manufactured by electrolysis of water. Properties of Hydrogen gas Physical properties i. It is a colourless, odourless and tasteless gas. ii. It is almost insoluble in water. iii. It does not show any reaction on the litmus paper. So it is a neutral gas. iv. It liquefies at – 253°C and solidifies at – 259° C. v. It is lighter than air. Chemical Properties i. Hydrogen gas burns in air or oxygen and forms water. 2H2 + O2 burn 2H2O ii. At very high temperature, hydrogen gas reacts with non-metals and from different compounds. C + 2H2 burn CH4 (Methane) S + H2 burn H2S (Hydrogen sulphide) 3. It reacts with metals like sodium, potassium, calcium, etc. to form hydrides. 2Na + H2 burn 2NaH (Sodium hydride) 2K + H2 burn 2KH (Potassium hydride) Ca + H2 burn CaH2 (Calcium hydride) 4. When hydrogen gas is passed through heated metallic oxides, they are reduced to metal. Fe2 O3 + H2 heat 3Fe + 4H2O CuO + H2 Cu + H2O Pb O + H2 Pb + H2O GREEN Science (Chemistry) Book-9 197

This process is used to obtain metals from metallic oxides. This process is known as reduction process. In this process, hydrogen removes the oxygen present in metal oxide and forms water. So, it is called reduction reaction. Uses of Hydrogen Gas i. It is used to obtain metals from metallic oxides. ii. It is used to produce ammonia gas. iii. It is used in the manufacturing of vegetable ghee. When hydrogen gas is passed into vegetable oil in the presence of heat and catalyst, oil is converted into ghee, which is called vegetable ghee. This process is called hydrogenation. vegetable oil + H2 8 – 10 atm vegetable ghee ∼ 250°C iv. It is used as a fuel in rockets. Sufficient energy is released during the chemical reaction between hydrogen and oxygen which helps in the projection of rocket. v. A mixture of hydrogen gas and oxygen gas burns to produce a very high temperature at about 3000°C. It is called oxy-hydrogen flame. This flame is used in cutting and welding of metals. 2. Oxygen Gas Oxygen gas is one of the essential gases important to all living creatures on the earth. Oxygen kills the harmful bacteria present inside our body. Oxygen makes up about 21% of atmosphere (by volume). This gas is found in water, carbohydrate, limestone, silicate, wood, etc. in the form of compound. About 49% of the compounds of earth's crust possess oxygen. The oxygen gas was first prepared by a British Scientist Joseph Priestley in 1774 AD by heating red oxide of mercury. Later Lavoisier named it oxygen. Lavoisier proved by an experiment that oxygen gas constitutes about 1/5 part of the air. Some facts about oxygen gas are tabulated below: Symbol Molecular Atomic Atomic weight Electronic O formula number 16 Configuration O2 8 1s²2s²2p4 Methods of Preparation of Oxygen Gas 1. Laboratory Preparation of Oxygen Gas a. By Heating Oxygen is prepared by heating compounds rich in oxygen. In the laboratory, oxygen gas is prepared by heating potassium chlorate (KClO3) with Manganese dioxide (MnO2), Here, manganese dioxide acts as a catalyst, which increases the rate of chemical reaction. 2KClO3 MnO2 2KCl + 3O2 198 GREEN Science (Chemistry) Book-9

Procedure The mixture of potassium chlorate (KClO3) and manganese dioxide (MnO2) in the ratio of 4 : 1 is put in a hard glass test - tube and the rest of the apparatus is fitted as shown in the figure. Oxygen evolves when the mixture in a hard glass test-tube is heated at the temperature of 250°C. The oxygen thus produced is collected in the jar by downward displacement of water. Hard glass test tube Cork Oxygen gas Potassium chlorate Delivery tube Gas jar + Manganese dioxide Fig. Fig. Stand Water trough Beehive shelf Bunsen burner 11.3 Preparation of oxygen gas (by heating) Precautions i. The hard glass test tube should be inclined as shown in figure. ii. The apparatus should be made airtight. iii. The gas jar should be full of water or there should not be presence of air bubbles. b. Without Using Heat In the laboratory, oxygen gas can be prepared form the hydrogen peroxide (H2O2) and manganese dioxide (MnO2) without using heat. In this process, manganese dioxide is used as a catalyst. 2H2O2 MnO2 H2O + O2 Procedure Take a little of manganese dioxide and water in a conical flask and set the apparatus as shown in figure. Pour hydrogen peroxide slowly in the flask through the dropping thistle funnel. Oxygen gas is produced immediately when hydrogen peroxide comes in contact with manganese dioxide. Hydrogen peroxide Delivery tube Stopper Cork Oxygen gas Gas jar Conical flask 11.4 Mixture of H2O2 + MnO2 Beehive shelf Water Water trough Preparation of oxygen (without heating) GREEN Science (Chemistry) Book-9 199

Precautions i. The apparatus should be made airtight. ii. The top of thistle funnel should be immersed in hydrogen peroxide. iii. Pure manganese dioxide should be used. Test of Oxygen How we know the gas prepared in laboratory is oxygen? Think yourself. When a burning matchstick is put near the mouth of the gas jar containing oxygen gas, it burns more brightly. This is because oxygen gas is a supporter of burning. Manufacture of Oxygen gas Oxygen is found in abundant from in water and air. So, manufacture of oxygen is possible from these two sources. a. From electrolysis of water Oxygen gas can be manufactured by the electrolysis of water. 2H2 O 2H2 + O2 b. From liquid air The air can be changed into liquid state with a high pressure after cooling. In liquid air there is mostly nitrogen and then oxygen. The boiling point of liquid nitrogen is (– 190°C) and that of liquid oxygen is (– 183°C). So, when liquid air is heated nitrogen is evolved in first section and oxygen still remains in liquid state. This oxygen in liquid state is heated at – 183°C and liquid oxygen is changed into gas. The oxygen and nitrogen gases evolved are collected in separate cylinders. In this way, manufacture of oxygen is possible from liquid air. This method helps to manufacture both oxygen and nitrogen gases. Properties of Oxygen gas a. Physical properties i. Oxygen is colourless, odourless and tasteless gas. ii. Oxygen is slightly heavier than air. iii. It is slightly soluble in water. iv. It is neutral. So, it has no effect on indicator. v. It can be liquefied at – 183°C and solidified at – 219°C. 200 GREEN Science (Chemistry) Book-9