CBSE Class 9 Science Atoms and Molecules Sure Shot Questions Set 01

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ASSIGNMENT QUESTIONS SET - 1
CHAPTER - 3
ATOMS AND MOLECULES

Calculate the mass of one mole of these substances.
1. AlCl₃
2. TeF₄
3. PbS
4. Cu₂O
5. AgI
6. N₂O
7. MoCl₅
8. Hg₂Br₂
9. Ta₂O₅
10. HgF₂
11. KCl
12. KF
13. ZnO
14. Ba(SCN)₂
15. K₂S
16. NH₄Cl
17. KH₂PO₄
18. C₂H₅NBr
19. Ba(ClO₃)₂
20. Fe(OH)₃
21. (NH₄)₂S
22. CoCl₂
23. KMnO₄
24. CaSO₄
25. H₂CO₃
26. CO₂
27. LiH
28. CO
29. SnI₄
30. KOH
31. K₂O
32. H₂SO₄
33. Hg₃N₂
34. SiF₄
35. NH₄OH
36. N₂O₅
37. SnCrO₄
38. Al₂O₃
39. CuCO₃
40. Ba(BrO₃)₂
41. Hg₂Cl₂
42. Cr₂(SO₃)₃
43. Al(MnO₄)₃
44. CoSO₄
45. Ca(NO₃)₃
46. NaH₂PO₄
47. (NH₄)₃PO₄
48. KAl(SO₄)₂
49. Hg₂SO₄
50. Al₂(SO₄)₃
51. FePO₄
52. Ca(C₂H₃O₂)₂
53. AlBr₃
54. P₂O₅
55. NH₄NO₃
56. Ba(OH)₂
57. PbSO₄
58. Ba₃(PO₄)₂
59. NaC₂H₃O₂
60. Ba(OH)₂
61. NaHCO₃
62. Al(OH)₃
63. NH₄MnO₄
64. Fe₂O₃
65. CaCO₃
66. HCl
67. K₂SO₄
68. NaCl
69. LiI
70. Hg₂O
71. HF
72. FeCl₃
73. NaHSO₄
74. Ag₂O
75. Pb(ClO₂)₂
76. CoF₃
77. Al(C₂H₃O₂)₃
78. Na₂Al₂(SO₄)₄
79. (HOOCCH₂)₂NCH₂CH₂N(CH₂COOH)₂
80. (NH₄)₂CH(CH₂)₅COOH

Answer:
1. AlCl₃ = 133.34 g/mol
2. TeF₄ = 203.59 g/mol
3. PbS = 239.3 g/mol
4. Cu₂O = 143.09 g/mol
5. AgI = 234.77 g/mol
6. N₂O = 44.01 g/mol
7. MoCl₅ = 273.20 g/mol
8. Hg₂Br₂ = 560.98 g/mol
9. Ta₂O₅ = 441.89 g/mol
10. HgF₂ = 238.59 g/mol
11. KCl = 74.55 g/mol
12. KF = 58.10 g/mol
13. ZnO = 81.38 g/mol
14. Ba(SCN)₂ = 253.48 g/mol (corrected: the printed key gives 255.26, which is wrong; Ba = 137.33, S = 32.06, C = 12.01, N = 14.01 gives 253.48 g/mol)
15. K₂S = 110.26 g/mol
16. NH₄Cl = 53.49 g/mol
17. KH₂PO₄ = 136.08 g/mol
18. C₂H₅NBr = 122.97 g/mol
19. Ba(ClO₃)₂ = 304.23 g/mol
20. Fe(OH)₃ = 106.87 g/mol
21. (NH₄)₂S = 68.14 g/mol
22. CoCl₂ = 129.84 g/mol
23. KMnO₄ = 158.03 g/mol
24. CaSO₄ = 136.14 g/mol
25. H₂CO₃ = 62.02 g/mol
26. CO₂ = 44.01 g/mol
27. LiH = 7.95 g/mol
28. CO = 28.01 g/mol
29. SnI₄ = 626.31 g/mol
30. KOH = 56.106 g/mol
31. K₂O = 94.20 g/mol
32. H₂SO₄ = 98.07 g/mol
33. Hg₃N₂ = 629.78 g/mol
34. SiF₄ = 104.08 g/mol
35. NH₄OH = 35.046 g/mol
36. N₂O₅ = 108.01 g/mol
37. SnCrO₄ = 234.68 g/mol
38. Al₂O₃ = 101.96 g/mol
39. CuCO₃ = 123.555 g/mol
40. Ba(BrO₃)₂ = 393.1314 g/mol
41. Hg₂Cl₂ = 472.09 g/mol
42. Cr₂(SO₃)₃ = 344.1666 g/mol
43. Al(MnO₄)₃ = 383.788 g/mol
44. CoSO₄ = 154.99 g/mol
45. Ca(NO₃)₃ = 226.09 g/mol
46. NaH₂PO₄ = 119.977 g/mol
47. (NH₄)₃PO₄ = 149.087 g/mol
48. KAl(SO₄)₂ = 258.195 g/mol
49. Hg₂SO₄ = 497.24 g/mol
50. Al₂(SO₄)₃ = 342.136 g/mol
51. FePO₄ = 150.82 g/mol
52. Ca(C₂H₃O₂)₂ = 158.169 g/mol
53. AlBr₃ = 266.69 g/mol
54. P₂O₅ = 141.944 g/mol
55. NH₄NO₃ = 80.04 g/mol
56. Ba(OH)₂ = 171.34 g/mol
57. PbSO₄ = 303.26 g/mol
58. Ba₃(PO₄)₂ = 601.93 g/mol
59. NaC₂H₃O₂ = 82.03 g/mol
60. Ba(OH)₂ = 171.34 g/mol
61. NaHCO₃ = 84.007 g/mol
62. Al(OH)₃ = 78.00 g/mol
63. NH₄MnO₄ = 136.97 g/mol
64. Fe₂O₃ = 159.69 g/mol
65. CaCO₃ = 100.09 g/mol
66. HCl = 36.461 g/mol
67. K₂SO₄ = 174.25 g/mol
68. NaCl = 58.443 g/mol
69. LiI = 133.846 g/mol
70. Hg₂O = 417.179 g/mol
71. HF = 20.006 g/mol
72. FeCl₃ = 162.206 g/mol
73. NaHSO₄ = 120.055 g/mol
74. Ag₂O = 231.74 g/mol
75. Pb(ClO₂)₂ = 342.10 g/mol
76. CoF₃ = 115.928 g/mol
77. Al(C₂H₃O₂)₃ = 204.12 g/mol
78. Na₂Al₂(SO₄)₄ = 484.173 g/mol
79. (HOOCCH₂)₂NCH₂CH₂N(CH₂COOH)₂ = 292.246 g/mol
80. (NH₄)₂CH(CH₂)₅COOH = 164.248 g/mol

Teacher's Note:
a) Molar mass is the sum of the atomic masses of all atoms in the formula, with brackets multiplied out.
b) Multiply each atomic mass by the number of atoms, including the subscript outside a bracket, such as (NH₄)₂.
c) Check that the formula is chemically valid before you calculate, because two printed entries had a wrong key or an invalid formula.

1. Who established the two important laws of chemical combinations?

Answer:
Antoine Lavoisier established the law of conservation of mass, and Joseph Proust established the law of constant proportions (definite proportions).

Teacher's Note:
a) Lavoisier worked on the conservation of mass in reactions.
b) Proust gave the law of constant composition for compounds.

2. What are the laws of chemical combinations?

Answer:
1. Law of conservation of mass: mass is neither created nor destroyed in a chemical reaction.
2. Law of constant proportions: a chemical compound always contains the same elements combined in the same fixed mass ratio.

Teacher's Note:
a) Both laws apply to every chemical reaction and compound.
b) Dalton's law of multiple proportions is a third related law, but these two are the basic laws.

3. What is the law of conservation of mass?

Answer:
The law of conservation of mass states that the total mass of the reactants is equal to the total mass of the products in a chemical reaction. Mass can neither be created nor destroyed.

Teacher's Note:
a) Count the mass of all reactants and all products, and they must be equal.
b) Mass is conserved, but the substances change their form and properties.

4. Give an example to show Law of conservation of mass applies to physical change also.

Answer:
1. Take 18 g of ice.
2. When the ice melts, it forms 18 g of water.
3. The mass remains 18 g, so the total mass is conserved in a physical change too.

Teacher's Note:
a) Ice and water are the same substance in different states.
b) The mass remains the same even though the state changes.

5. Explain with example that law of conservation of mass is valid for chemical reactions.

Answer:
1. Take 2 g of hydrogen and 16 g of oxygen.
2. They react to form water: 2H₂ + O₂ → 2H₂O.
3. Mass of reactants = 2 g + 16 g = 18 g.
4. Mass of water formed = 18 g.
5. Hence, the mass is conserved in the chemical reaction.

Teacher's Note:
a) Balance the equation first, then compare the masses.
b) Mass of reactants equals mass of products, which is the proof of conservation.

6. The 2.8 g of nitrogen gas was allowed to react with 0.6 g of hydrogen gas to produce 3.4 g of ammonia. Show that these observations are in agreement with the law of Conservation of mass.

Answer:
1. Mass of nitrogen = 2.8 g.
2. Mass of hydrogen = 0.6 g.
3. Total mass of reactants = 2.8 g + 0.6 g = 3.4 g.
4. Mass of ammonia formed = 3.4 g.
5. Since the mass of reactants equals the mass of product, the observation agrees with the law of conservation of mass.

Teacher's Note:
a) Add the masses of all reactants first.
b) The product mass must equal the sum of the reactant masses.

7. If 12 g of carbon is burnt in the presence of 32 g of oxygen, how much carbon dioxide will be formed?

Answer:
1. Carbon + oxygen → carbon dioxide.
2. Mass of carbon = 12 g and mass of oxygen = 32 g.
3. Total mass of reactants = 12 g + 32 g = 44 g.
4. By the law of conservation of mass, mass of carbon dioxide formed = 44 g.

Teacher's Note:
a) The reactants are used up completely in this case.
b) Do not add the masses unless all the substances react.

8. Who proposed Law of Definite Proportions (or Law of Constant Composition)?

Answer:
Joseph Proust proposed the law of definite proportions, also called the law of constant composition.

Teacher's Note:
a) Proust was a French chemist.
b) The law is also called the law of constant proportions.

9. State Law of constant proportions. Explain with an example.

Answer:
1. Law of constant proportions: a chemical compound always contains the same elements combined in the same fixed mass ratio, whatever its source.
2. Example: in water, hydrogen and oxygen are always present in the mass ratio 1 : 8, whether the water is from a river or is made in a laboratory.

Teacher's Note:
a) The ratio is fixed for a given compound.
b) Water from any source gives the same ratio.

10. Show that water illustrates the law of constant proportions.

Answer:
1. Take 9 g of water. It contains 1 g of hydrogen and 8 g of oxygen. Ratio H : O = 1 : 8.
2. Take 18 g of water from another source. It contains 2 g of hydrogen and 16 g of oxygen. Ratio H : O = 2 : 16 = 1 : 8.
3. The ratio is the same in both samples, so water illustrates the law of constant proportions.

Teacher's Note:
a) Always compare the ratio, not the total mass.
b) Any amount of water gives the same 1 : 8 ratio.

11. Hydrogen and oxygen combine in the ratio of 1:8 by mass to form water. What mass of oxygen gas would be required to react completely with 3 g of hydrogen gas?

Answer:
1. Ratio of hydrogen to oxygen by mass = 1 : 8.
2. Let the mass of oxygen be x g. Then 3 : x = 1 : 8.
3. x = 3 × 8 = 24 g.
4. Hence, 24 g of oxygen is required.

Teacher's Note:
a) Set up the ratio with hydrogen first.
b) The answer is 8 times the hydrogen mass, which is 24 g.

12. A 0.24 g sample of compound of oxygen and boron was found by analysis to contain 0.096 g of boron and 0.144 g of oxygen. Calculate the percentage composition of the compound by weight.

Answer:
1. Mass of the compound = 0.24 g.
2. Percentage of boron = \( \frac{0.096}{0.24} \) × 100 = 40 %.
3. Percentage of oxygen = \( \frac{0.144}{0.24} \) × 100 = 60 %.
4. Hence, the compound contains 40 % boron and 60 % oxygen by weight.

Teacher's Note:
a) Check that the two percentages add up to 100 %.
b) Use the mass of the whole sample as the denominator.

13. When 3.0 g of carbon is burnt in 8.00 g oxygen, 11.00 g of carbon dioxide is produced. What mass of carbon dioxide will be formed when 3.00 g of carbon is burnt in 50.00 g of oxygen? Which law of chemical combination will govern your answer?

Answer:
1. In the first reaction, carbon and oxygen combine in the mass ratio 3.0 : 8.00 to form 11.00 g of carbon dioxide.
2. In the second reaction, 3.00 g of carbon combines with only 8.00 g of oxygen, because the ratio of carbon to oxygen must stay 3 : 8.
3. Oxygen left over = 50.00 g - 8.00 g = 42.00 g, which does not react.
4. Mass of carbon dioxide formed = 3.00 g + 8.00 g = 11.00 g.
5. The answer is governed by the law of constant proportions, and it also follows the law of conservation of mass.

Teacher's Note:
a) The extra oxygen is left unreacted, so do not add it to the product.
b) Carbon and oxygen always combine in a fixed ratio.

14. Magnesium and oxygen combine in the ratio of 3 : 2 by mass to form magnesium oxide. How much oxygen is required to react completely with 12 g of magnesium?

Answer:
1. Ratio of magnesium to oxygen by mass = 3 : 2.
2. Let the mass of oxygen be x g. Then 3 : 2 = 12 : x.
3. x = \( \frac{12 \times 2}{3} \) = 8 g.
4. Hence, 8 g of oxygen is required.

Teacher's Note:
a) Oxygen is always two-thirds of the magnesium mass here.
b) Check: 12 g + 8 g = 20 g of magnesium oxide.

15. What are the postulates of Dalton's atomic theory?

Answer:
1. All matter is made of very tiny particles called atoms, which cannot be divided further in a chemical reaction.
2. Atoms of the same element are identical in mass and properties.
3. Atoms of different elements have different masses and properties.
4. Atoms combine in simple whole number ratios to form compounds.
5. Atoms are neither created nor destroyed in a chemical reaction; they are only rearranged.

Teacher's Note:
a) Dalton's theory explains the laws of chemical combination.
b) The theory was later modified because atoms can be divided into subatomic particles.

16. Which of the following statements is NOT true about an atom?
(a) Atoms are the building blocks from which molecules and ions are formed.
(b) Atoms cannot exist independently.
(c) Atoms are neutral in nature
(d) Atoms combine together to form matter that we can see, feel or touch.

Answer: (b) Atoms cannot exist independently.

1. Statements (a), (c) and (d) are true. Atoms are the building blocks of molecules and ions, they are electrically neutral, and atoms combine to form the matter we see.
2. Noble gases such as helium, neon and argon exist as single independent atoms.
3. Hence, the statement that atoms cannot exist independently is NOT true, and the correct option is (b).

Teacher's Note:
a) Most elements do not exist as free atoms, but some do.
b) Learn the exceptions: helium, neon and argon are monoatomic.

17. What is an atom?

Answer:
An atom is the smallest particle of an element that can take part in a chemical reaction and still shows the properties of that element.

Teacher's Note:
a) Atoms are extremely small, about 0.1 nanometre across.
b) An atom is the basic building block of all matter.

18. Why is it not possible to see an atom with naked eyes?

Answer:
1. Atoms are extremely small, about 10^-10 m in size.
2. The human eye can see only objects larger than about 0.1 mm, so atoms are far too small to be seen.

Teacher's Note:
a) Even the most powerful light microscope cannot show single atoms.
b) Electron microscopes and scanning probes are needed to observe atoms.

19. Who proposed the chemical notation based on first two letters of the name of the element?

Answer:
Jöns Jacob Berzelius proposed the chemical notation in which the symbol of an element is based on the first one or two letters of its name.

Teacher's Note:
a) Berzelius introduced the system that is used today.
b) The first letter is always capital and the second is small.

20. Name the international organization who approves names of elements.

Answer:
The International Union of Pure and Applied Chemistry (IUPAC) approves the names of elements.

Teacher's Note:
a) IUPAC sets the standard names and symbols for elements.
b) Its name is the abbreviation IUPAC.

21. What is the chemical symbol for iron?

Answer:
The chemical symbol for iron is Fe.

Teacher's Note:
a) Fe comes from the Latin name ferrum.
b) Do not write the symbol as FE or Fa.

22. Name five elements have single letter chemical symbol.

Answer:
1. Hydrogen (H)
2. Carbon (C)
3. Nitrogen (N)
4. Oxygen (O)
5. Sulphur (S)

Teacher's Note:
a) Several elements have a one-letter symbol, such as B, F, I, K, P, U, V, W and Y.
b) The symbols are written in capital letters.

23. Name the element having following Latin names
(i) Stibium
(ii) Cuprum
(iii) Argentum
(iv) Natrium
(v) Stannum
(vi) Wolfram
(vii) plumbum
(viii) Kalium

Answer:
1. (i) Stibium is antimony (Sb).
2. (ii) Cuprum is copper (Cu).
3. (iii) Argentum is silver (Ag).
4. (iv) Natrium is sodium (Na).
5. (v) Stannum is tin (Sn).
6. (vi) Wolfram is tungsten (W).
7. (vii) Plumbum is lead (Pb).
8. (viii) Kalium is potassium (K).

Teacher's Note:
a) Many symbols come from Latin names, which is why they do not match English names.
b) Lead, tin and silver are common examples.

24. Write the chemical symbols of the following:
(i) Gold
(ii) Iron
(iii) Chlorine
(iv) Mercury

Answer:
1. (i) Gold: Au
2. (ii) Iron: Fe
3. (iii) Chlorine: Cl
4. (iv) Mercury: Hg

Teacher's Note:
a) Au comes from the Latin word aurum.
b) Chlorine has two letters, so the second letter is small.

25. How will you define chemical symbol?

Answer:
A chemical symbol is a one or two letter short form of the name of an element. The first letter is always capital, and the second letter, if present, is small.

Teacher's Note:
a) Symbols are used for elements in formulae and equations.
b) They are based on English or Latin names.

26. What is the significance of a chemical symbol?

Answer:
1. It represents the name of the element.
2. It represents one atom of the element.
3. It represents the element's fixed mass, such as one mole of atoms when used with its atomic mass.

Teacher's Note:
a) For example, O stands for oxygen, one atom of oxygen, and 16 u.
b) Symbols help us to write formulae and equations easily.

27. Can atoms of an element exist independently? Give examples of elements which exist in atomic form. Give examples of elements that do not exist in atomic form.

Answer:
1. Most elements cannot exist independently as atoms.
2. Elements that exist in atomic form: helium (He), neon (Ne) and argon (Ar).
3. Elements that do not exist in atomic form: hydrogen (H₂), oxygen (O₂) and nitrogen (N₂).

Teacher's Note:
a) Noble gases are stable in atomic form.
b) H₂, O₂ and N₂ exist as molecules of two atoms.

28. Why do atoms of the most of the elements not exist independently?

Answer:
1. Atoms of most elements are not stable on their own because their outer shells are incomplete.
2. To attain a stable electronic configuration, the atoms combine with each other or with other atoms.
3. Hence, they exist as molecules or compounds, not as free atoms.

Teacher's Note:
a) Noble gases have full outer shells, so they are stable as single atoms.
b) Bonding makes the atoms more stable.

29. Which element has the smallest atom in size?

Answer:
Hydrogen has the smallest atom in size.

Teacher's Note:
a) Hydrogen has only one proton and one electron.
b) Atomic size increases down a group in the periodic table.

30. What is the atomic mass unit?

Answer:
The atomic mass unit (u) is one-twelfth of the mass of one carbon-12 atom. It is used to express the masses of atoms.

Teacher's Note:
a) 1 u is about 1.66 × 10^-27 kg.
b) The unit is written as u.

31. Magnesium is two times heavier than C-12 atom, what shall be the mass of Mg atom in terms of atomic mass units? (Given mass of C-12 atom = 12u)

Answer:
1. Mass of C-12 atom = 12 u.
2. Magnesium is two times heavier, so mass of Mg atom = 2 × 12 u = 24 u.
3. Hence, the mass of a Mg atom is 24 u.

Teacher's Note:
a) Multiply the given mass by 2, not add 2.
b) The answer matches the atomic mass of magnesium.

32. What is relative atomic mass of an element? How it is related to atomic mass unit?

Answer:
1. Relative atomic mass of an element is the number of times an atom of the element is heavier than 1/12 of the mass of one carbon-12 atom.
2. It is related to the atomic mass unit because 1 u is 1/12 of the mass of one carbon-12 atom, so the relative atomic mass is expressed in u.
3. For example, the relative atomic mass of magnesium is 24 u.

Teacher's Note:
a) Relative atomic mass has no unit in its ratio form, but it is expressed in u.
b) It is the same number as the atomic mass in u.

33. Define molecule. What are its important properties?

Answer:
A molecule is the smallest particle of an element or a compound that can exist independently and shows all the properties of that substance.
Its important properties are: it is formed by two or more atoms chemically bonded together; it can exist independently; and it shows the same chemical properties as the substance.

Teacher's Note:
a) Molecules of compounds have atoms of different elements.
b) A molecule of an element has atoms of the same element.

34. Based on type of substance, how molecules are classified?

Answer:
1. Molecules of elements, which have atoms of one element only, such as O₂.
2. Molecules of compounds, which have atoms of two or more different elements, such as H₂O and CO₂.

Teacher's Note:
a) Elements and compounds are the two main classes of substances.
b) Look at the number of different elements in the molecule.

35. What is atomicity?

Answer:
Atomicity is the number of atoms present in one molecule of an element or a compound.

Teacher's Note:
a) Atomicity is the number of atoms, not the number of elements.
b) For example, the atomicity of O₂ is 2.

36. Based on atomicity, how molecules are categorized?

Answer:
1. Monoatomic molecules: one atom per molecule, such as helium (He).
2. Diatomic molecules: two atoms per molecule, such as hydrogen (H₂).
3. Triatomic molecules: three atoms per molecule, such as water (H₂O).
4. Tetra-atomic molecules: four atoms per molecule, such as phosphorus (P₄).
5. Polyatomic molecules: more than four atoms per molecule, such as sulphur (S₈).

Teacher's Note:
a) Classify by the number of atoms, not by the elements.
b) S₈ is a polyatomic molecule with eight atoms.

37. Give three examples of monoatomic molecules.

Answer:
1. Helium (He)
2. Neon (Ne)
3. Argon (Ar)

Teacher's Note:
a) Noble gases exist as monoatomic molecules.
b) Each of these has one atom in its molecule.

38. Give four examples of diatomic molecules.

Answer:
1. Hydrogen (H₂)
2. Oxygen (O₂)
3. Nitrogen (N₂)
4. Chlorine (Cl₂)

Teacher's Note:
a) Each molecule has two atoms of the same element.
b) Carbon monoxide (CO) is a diatomic compound.

39. (i) What is the chemical formula of Water molecule? (ii) What is its atomicity? (iii) Calculate the ratio of masses of atoms of elements present in water molecule. (iv) Calculate the ratio by number of atoms of elements present in water molecule.

Answer:
1. (i) Chemical formula of water: H₂O.
2. (ii) Atomicity of water = 3, since one molecule has 2 hydrogen atoms and 1 oxygen atom.
3. (iii) Ratio of masses = H : O = (2 × 1) : (1 × 16) = 2 : 16 = 1 : 8.
4. (iv) Ratio by number of atoms = H : O = 2 : 1.

Teacher's Note:
a) Mass ratio and atom ratio are different, so do not mix them up.
b) The mass ratio of water is always 1 : 8.

40. What is an ion?

Answer:
An ion is an atom or a group of atoms that carries an electric charge. A positively charged ion is called a cation and a negatively charged ion is called an anion.

Teacher's Note:
a) Gaining electrons gives a negative ion.
b) Losing electrons gives a positive ion.

41. What are polyatomic ions? Give examples?

Answer:
Polyatomic ions are charged groups of two or more atoms that behave as a single unit. Examples: ammonium (NH₄⁺), sulphate (SO₄^2-) and phosphate (PO₄^3-).

Teacher's Note:
a) The whole group carries the charge, not just one atom.
b) Polyatomic ions appear in many common compounds.

42. Give examples of triatomic molecules.

Answer:
1. Water (H₂O)
2. Carbon dioxide (CO₂)
3. Ozone (O₃)

Teacher's Note:
a) Each of these has three atoms in one molecule.
b) Sulphur dioxide (SO₂) is another example.

43. What is valency of an element?

Answer:
Valency of an element is the combining capacity of its atom, that is, the number of electrons an atom gains, loses or shares to attain a stable electronic configuration.

Teacher's Note:
a) Valency is the number of bonds an atom can form.
b) For example, the valency of oxygen is 2.

44. What is meant by the term chemical formula?

Answer:
A chemical formula is the symbolic representation of a compound, showing the elements present and the number of atoms of each element in one molecule or one formula unit.

Teacher's Note:
a) For example, H₂O shows two hydrogen atoms and one oxygen atom.
b) The formula is written using symbols of the elements.

45. Write down the formulae of
(i) sodium oxide
(ii) aluminium chloride
(iii) sodium suphide
(iv) magnesium hydroxide

Answer:
1. (i) Sodium oxide: Na₂O.
2. (ii) Aluminium chloride: AlCl₃.
3. (iii) Sodium sulphide: Na₂S.
4. (iv) Magnesium hydroxide: Mg(OH)₂.

Teacher's Note:
a) Balance the charges of the ions to get the formula.
b) Sodium forms Na⁺, and oxygen forms O^2-, so two sodium ions are needed.

46. Write down the names of compounds represented by the following formulae:
(i) Al₂(SO₄)₃
(ii) CaCl₂
(iii) K₂SO₄
(iv) KNO₃
(v) CaCO₃

Answer:
1. (i) Aluminium sulphate.
2. (ii) Calcium chloride.
3. (iii) Potassium sulphate.
4. (iv) Potassium nitrate.
5. (v) Calcium carbonate.

Teacher's Note:
a) Name the metal first, then the non-metal radical.
b) The ending of the radical changes with the number of oxygen atoms.

47. Write the chemical formulae of the following. Also identify the ions present.
(a) Magnesium chloride
(b) Calcium oxide
(c) Copper nitrate
(d) Aluminium chloride
(e) Calcium carbonate.

Answer:
1. (a) Magnesium chloride: MgCl₂; ions Mg²⁺ and Cl^-.
2. (b) Calcium oxide: CaO; ions Ca²⁺ and O^2-.
3. (c) Copper nitrate: Cu(NO₃)₂; ions Cu²⁺ and NO₃^-.
4. (d) Aluminium chloride: AlCl₃; ions Al³⁺ and Cl^-.
5. (e) Calcium carbonate: CaCO₃; ions Ca²⁺ and CO₃^2-.

Teacher's Note:
a) Balance the positive and negative charges in each formula.
b) Use brackets when a polyatomic ion appears more than once.

48. Give the names of the elements present in the following compounds.
(a) Quick lime
(b) Hydrogen bromide
(c) Baking powder
(d) Potassium sulphate.

Answer:
1. (a) Quick lime (CaO): calcium and oxygen.
2. (b) Hydrogen bromide (HBr): hydrogen and bromine.
3. (c) Baking powder (sodium hydrogen carbonate, NaHCO₃): sodium, hydrogen, carbon and oxygen.
4. (d) Potassium sulphate (K₂SO₄): potassium, sulphur and oxygen.

Teacher's Note:
a) Quick lime is calcium oxide.
b) Baking powder is sodium hydrogen carbonate.

49. How many atoms are present in a (i) H₂S molecule and (ii) PO₄^3- ion?

Answer:
1. (i) H₂S molecule has 2 hydrogen atoms and 1 sulphur atom, so 3 atoms in total.
2. (ii) PO₄^3- ion has 1 phosphorus atom and 4 oxygen atoms, so 5 atoms in total.

Teacher's Note:
a) Count each subscript and add them.
b) The charge on an ion does not change the number of atoms.

50. (a) Write a chemical formula of a compound using zinc ion and phosphate ion. (b) Calculate the ratio by mass of atoms present in a molecule of carbon dioxide. (Given C =12, O =16)

Answer:
1. (a) Zinc ion is Zn²⁺ and phosphate ion is PO₄^3-. To balance the charges, 3 zinc ions and 2 phosphate ions are needed, so the formula is Zn₃(PO₄)₂.
2. (b) In CO₂, mass of C = 12 and mass of O = 2 × 16 = 32.
3. Ratio by mass C : O = 12 : 32 = 3 : 8.

Teacher's Note:
a) Balance the charges, not just the number of ions.
b) The mass ratio is 3 : 8, not the atom ratio 1 : 2.

51. What is the molecular mass of a substance?

Answer:
Molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of the substance, expressed in u.
For example, the molecular mass of water (H₂O) is 2 × 1 + 16 = 18 u.

Teacher's Note:
a) Add the atomic mass of every atom, including repeated atoms.
b) Molecular mass is used for covalent compounds.

52. What is Formula Unit Mass? How it is different from molecular mass?

Answer:
1. Formula unit mass is the sum of the atomic masses of all the atoms in one formula unit of an ionic compound, expressed in u.
2. It is different from molecular mass because an ionic compound does not contain separate molecules. It is represented by a formula unit, which is the simplest ratio of ions, while molecular mass is for a discrete molecule.

Teacher's Note:
a) Sodium chloride has a formula unit, not a molecule.
b) Both are calculated by adding atomic masses.

53. Calculate the formula unit masses of ZnO, Na₂O, K₂CO₃, given atomic masses of Zn = 65 u, Na = 23 u, K = 39 u, C = 12 u, and O = 16 u.

Answer:
1. ZnO: 65 + 16 = 81 u.
2. Na₂O: 2 × 23 + 16 = 62 u.
3. K₂CO₃: 2 × 39 + 12 + 3 × 16 = 78 + 12 + 48 = 138 u.

Teacher's Note:
a) Multiply each atomic mass by the number of atoms first.
b) Do not forget the oxygen atoms in the carbonate.

54. What are ionic compounds?

Answer:
Ionic compounds are compounds formed by the transfer of electrons from a metal to a non-metal, so that they consist of positive and negative ions held together by strong electrostatic forces. An example is sodium chloride (NaCl).

Teacher's Note:
a) Ionic compounds are usually solids at room temperature.
b) They conduct electricity when molten or in solution.

55. How do we know the presence of atoms if they do not exist independently for most of the elements?

Answer:
1. Atoms of most elements combine to form molecules and compounds.
2. The properties of these compounds show that they are made of atoms of different elements; for example, water is made of hydrogen and oxygen.
3. Chemical reactions and their mass relations, such as the laws of chemical combination, also show that atoms exist.

Teacher's Note:
a) Chemical reactions can be explained only by atoms.
b) Atoms are inferred from their combinations, not seen directly.

56. An element 'Z' forms the following compound when it reacts with hydrogen, chlorine, oxygen and phosphorus.
ZH₃, ZCl₃, Z₂O₃ and ZP
(a) What is the valency of element Z?
(b) Element 'Z' is metal or non-metal?

Answer:
1. (a) In ZH₃, ZCl₃ and Z₂O₃, the valency of Z is 3.
2. (b) Z is a metal, because it forms an oxide Z₂O₃ and a chloride ZCl₃ of the type formed by metals with valency 3, such as aluminium (AlH₃, AlCl₃, Al₂O₃ and AlP).

Teacher's Note:
a) Valency is the number of atoms of hydrogen or chlorine that one atom of Z combines with.
b) A metal of valency 3 forms such compounds.

57. Name one element each which forms diatomic and tetra atomic molecule.

Answer:
1. Diatomic molecule: hydrogen (H₂).
2. Tetra-atomic molecule: phosphorus (P₄).

Teacher's Note:
a) Oxygen and nitrogen are also diatomic.
b) White phosphorus is the common tetra-atomic form.

58. Name one element which forms diatomic and triatomic molecule.

Answer:
Oxygen forms a diatomic molecule (O₂) and a triatomic molecule (O₃, ozone).

Teacher's Note:
a) O₃ is called ozone.
b) The same element can have different atomicities.

59. What is gram-atomic mass of an element?

Answer:
Gram-atomic mass of an element is the atomic mass of the element expressed in grams. It is the mass of one mole of atoms of that element.

Teacher's Note:
a) For oxygen, the gram-atomic mass is 16 g.
b) It is numerically equal to the atomic mass in u.

60. What is gram-molecular mass of a substance?

Answer:
Gram-molecular mass of a substance is its molecular mass expressed in grams. It is the mass of one mole of molecules of that substance.

Teacher's Note:
a) For water, the gram-molecular mass is 18 g.
b) It is numerically equal to the molecular mass in u.

61. Define mole. What is its significance?

Answer:
A mole is the amount of a substance that contains 6.022 × 10²³ particles, such as atoms, molecules or ions. This number is called Avogadro's number.
Its significance is that it gives a convenient unit for expressing very large numbers of particles, and it relates the number of particles to the mass of the substance.

Teacher's Note:
a) One mole of any substance contains the same number of particles.
b) A mole of carbon is 12 g, and a mole of water is 18 g.

62. What is molar mass?

Answer:
Molar mass is the mass of one mole of a substance. It is expressed in grams per mole (g/mol) and is numerically equal to the atomic or molecular mass in u.

Teacher's Note:
a) The unit of molar mass is g/mol.
b) For example, the molar mass of water is 18 g/mol.

63. Who introduced the term 'mole' in chemistry?

Answer:
Wilhelm Ostwald introduced the term 'mole' in chemistry.

Teacher's Note:
a) The term comes from the Latin word moles, meaning a heap or pile.
b) Ostwald was a German chemist.

64. When 'mole' was chosen internationally standard way to express larger number of chemical units?

Answer:
The mole was adopted as the international standard unit for the amount of substance in 1971, by the International Committee of Weights and Measures.

Teacher's Note:
a) It is one of the seven SI base units.
b) The year to remember is 1971.

65. How many moles are there in 4.6 gms of Sodium(Na)?

Answer:
1. Atomic mass of Na = 23 g/mol.
2. Number of moles = \( \frac{\text{Given mass}}{\text{Molar mass}} \) = \( \frac{4.6}{23} \) = 0.2 mol.
3. Hence, there are 0.2 moles of sodium in 4.6 g.

Teacher's Note:
a) Divide the mass by the molar mass, not by the atomic number.
b) The answer is less than one mole, as expected for a small mass.

66. If one mole of carbon atoms weighs 12 gram, what is the mass (in gram) of 1 atom of carbon?

Answer:
1. One mole of carbon atoms = 6.022 × 10²³ atoms, and its mass is 12 g.
2. Mass of 1 atom = \( \frac{12}{6.022 \times 10^{23}} \) g = 1.99 × 10^-23 g.
3. Hence, the mass of one carbon atom is 1.99 × 10^-23 g.

Teacher's Note:
a) Divide the mass of one mole by Avogadro's number.
b) The answer is a very small number, written in exponent form.

67. Which has more number of atoms, 100 grams of sodium or 100 grams of iron (given, atomic mass of Na = 23 u, Fe = 56 u)?

Answer:
1. Moles of sodium = \( \frac{100}{23} \) = 4.35 mol.
2. Moles of iron = \( \frac{100}{56} \) = 1.79 mol.
3. Since the number of atoms is proportional to the number of moles, 100 g of sodium has more atoms than 100 g of iron.
4. Hence, 100 g of sodium has more atoms.

Teacher's Note:
a) The element with the smaller atomic mass gives more atoms for the same mass.
b) Compare moles, not grams, to compare the number of atoms.

68. What is the mass of
(a) 1 mole of nitrogen atoms?
(b) 4 moles of aluminium atoms (Atomic mass of aluminium = 27)?
(c) 10 moles of sodium sulphite (Na₂SO₃)?

Answer:
1. (a) Mass of 1 mole of nitrogen atoms = 14 g.
2. (b) Mass of 4 moles of aluminium = 4 × 27 = 108 g.
3. (c) Molar mass of Na₂SO₃ = 2 × 23 + 32 + 3 × 16 = 46 + 32 + 48 = 126 g/mol. Mass of 10 moles = 10 × 126 = 1260 g.

Teacher's Note:
a) Use the atomic mass for atoms and the molar mass for compounds.
b) Mass = number of moles × molar mass.

69. Convert into moles:
(a) 12 g of oxygen gas
(b) 20 g of water
(c) 22 g of carbon dioxide

Answer:
1. (a) Molar mass of O₂ = 32 g/mol. Moles = 12 / 32 = 0.375 mol.
2. (b) Molar mass of H₂O = 18 g/mol. Moles = 20 / 18 = 1.11 mol.
3. (c) Molar mass of CO₂ = 44 g/mol. Moles = 22 / 44 = 0.5 mol.

Teacher's Note:
a) Oxygen gas is O₂, so use 32 g/mol, not 16.
b) Divide the given mass by the molar mass.

70. What is the mass of: (a) 0.2 mole of oxygen atoms? (b) 0.5 mole of water molecules?

Answer:
1. (a) Mass of oxygen atoms = 0.2 × 16 = 3.2 g.
2. (b) Mass of water = 0.5 × 18 = 9 g.

Teacher's Note:
a) Oxygen atoms have a molar mass of 16 g/mol.
b) Water molecules have a molar mass of 18 g/mol.

71. Find out number of atoms in 15 moles of He.

Answer:
1. Helium is monoatomic, so one mole of He contains 6.022 × 10²³ atoms.
2. Number of atoms in 15 moles = 15 × 6.022 × 10²³ = 9.03 × 10²⁴ atoms.

Teacher's Note:
a) Helium exists as single atoms, so the atom count equals the molecule count.
b) Multiply the number of moles by Avogadro's number.

72. Calculate the number of molecules of sulphur (S₈) present in 16 g of solid sulphur.

Answer:
1. Molar mass of S₈ = 8 × 32 = 256 g/mol.
2. Moles of S₈ = \( \frac{16}{256} \) = 0.0625 mol.
3. Number of molecules = 0.0625 × 6.022 × 10²³ = 3.76 × 10²² molecules.

Teacher's Note:
a) Use the molar mass of the molecule S₈, not of a single sulphur atom.
b) The answer is a very large number of molecules.

73. Calculate the number of aluminium ions present in 0.051 g of aluminium oxide.
(Hint: The mass of an ion is the same as that of an atom of the same element. Atomic mass of Al = 27 u)

Answer:
1. Molar mass of Al₂O₃ = 2 × 27 + 3 × 16 = 102 g/mol.
2. Moles of Al₂O₃ = \( \frac{0.051}{102} \) = 0.0005 mol.
3. Each formula unit has 2 aluminium ions, so moles of Al³⁺ = 2 × 0.0005 = 0.001 mol.
4. Number of aluminium ions = 0.001 × 6.022 × 10²³ = 6.022 × 10²⁰ ions.

Teacher's Note:
a) Count the aluminium ions per formula unit first.
b) The hint tells you the ion has the same mass as the atom.

74. Calculate the mass percentage of Carbon(C), Hydrogen (H) and Oxygen (O) in one molecule of glucouse (C₆H₁₂O₆). (Atomic mass of C = 12u, H = 1u and O = 16u)

Answer:
1. Molar mass of C₆H₁₂O₆ = 6 × 12 + 12 × 1 + 6 × 16 = 72 + 12 + 96 = 180 u.
2. Mass % of C = \( \frac{72}{180} \) × 100 = 40 %.
3. Mass % of H = \( \frac{12}{180} \) × 100 = 6.67 %.
4. Mass % of O = \( \frac{96}{180} \) × 100 = 53.33 %.
5. Check: 40 + 6.67 + 53.33 = 100 %.

Teacher's Note:
a) Find the total mass of the molecule first.
b) Use the mass of each element as a fraction of the total.

75. Calculate the number of molecules of phosphorus (P₄) present in 31 gram of phosphorus.

Answer:
1. Molar mass of P₄ = 4 × 31 = 124 g/mol.
2. Moles of P₄ = \( \frac{31}{124} \) = 0.25 mol.
3. Number of molecules = 0.25 × 6.022 × 10²³ = 1.51 × 10²³ molecules.

Teacher's Note:
a) Use P₄, not P, because phosphorus exists as P₄ molecules.
b) The mass of one mole of P₄ is 124 g.