ATOMS AND MOLECULES

Class IX Science
Notes for Atoms and Molecules

Syllabus

Particle nature, Basic units, Atoms and molecules, Law of constant proportions, Atomic and molecular masses.

Facts that Matter

• Law of Chemical Combination

Given by Lavoisier and Joseph L. Proust as follows:

(i) Law of conservation of mass: Mass can neither be created nor destroyed in a chemical reaction.

e.g.,A+B-C+D

Reactants R Products

Mass of reactants = Mass of products

(ii) Law of constant proportion: In a chemical substance the elements are always present in definite proportions by mass.

E.g., in water, the ratio of the mass of hydrogen to the mass of oxygen is always 1 : 8 respectively.

These laws lacked explanation. Hence, John Dalton gave his theory about the matter. He said that the smallest particle of matter is called ‘atom’.

• Dalton’s Atomic Theory

1. Every matter is made up of very small or tiny particles called atoms.

2. Atoms are not divisible and cannot be created or destroyed in a chemical reaction.

3. All atoms of a given element it are same in size, “mass and chemical properties.

4. Atoms of different elements are different in size, mass and chemical properties.

5. Atoms combine in the ratio of small whole number to form compounds.

6. The relative number and kinds of atoms are constant in a given compound. 9

• Atom

Atoms are the smallest particles of an element which can take reaction.

Size of an atom: Atomic radius is measured in nanometres.

Atomic radii of hydrogen atom = 1 × 10–10 m.

Symbols of atoms:

(a) Symbols for some elements as proposed by Dalton:

(b) Symbols of some common elements:

Name of the element	Latin name	Symbol
Hydrogen	ï¿½	H
Helium	ï¿½	He
Carobon	ï¿½	C
Copper	Cuprum	Cu
Cobalt		Co
Chlorine		Cl
Cadmium		Cd
Boron		B
Barium		Ba
Bromine		Br
Bismuth		Bi
Sodium	Natrium	Na
Potassium	Kalium	K
Iron	Ferrum	Fe
Gold	Aurum	Au
Silver	Argentum	Ag
Mercury	Hydragyrum	Hg

• Molecule

It is the smallest particle of an eleme4It- or a dolnpound which can wxist independently.

• Molecules of an element constitutes same type of atoms. 1.w

• Molecules may be monoatomic, di-atomic or polyatomic. IT .

• Molecules of compounds join together in defmite proportionsrand constitutes different type of atoms.

• Atomicity

The number of atoms constituting a Molecule is known as its atomicity.

Name of the element	Atomicity	Molecules formula
Helium	Monoatomic	He
Neon	Monoatomic	Ne
Argon	Monoatomic	Ar
Sodium	Monoatomic	Na
Iron	Monoatomic	Fe
Aluminium	Monoatomic	Al
Hydrogen	Di-atomic	H2
Oxygen	Di-atomic	O2
Chlorine	Di-atomic	Cl2
Nitrogen	Di-atomic	N2
Phosphorus	Polyatomic (Tetra)	P4
Sulphur	Polyatomic (Octa)	S8

• Ions

The charged particles (atoms) are called ions, they charge or negative charge on it:

Negatively charged ionis called anion (C1ï¿½).

Positively charge ion is called cation (Na+).

• Valency

The combining capacity of an element is known as its valency: Valency is used to fmd out how atom of an element will combine with the atom of another element to form a chemical compound.

(Every atom want, to become stable, to do so it may loose, gain or share electrongs.

(i) If an atom consists of 1, 2 or 3 electrons in its valgncesI ell then its valency is 1, 2 or 3 respectively,

(ii) If an atom consists of 5, 6 or 7 electrons in the outermost shell, then it will gain 3, 2 or 1 electron respectively and its valency will be 3, 2 or 1 respectively.

(ii) If an atom has 4 electrons in the outermost shell than it will she this electron and hence its valency will be 4.

(iv) If an atom has 8 electrons in the outermost shell then its valency is 0.

• Chemical Formulae

Rules: (i) The valencies or charges on the ion must balance.

(ii) A metal and non-metal compound should show the name or symbols of the metal first.

e.g., Na+ Cl– → NaCl

(iii) If a compound consist of polyatomic ions. The ion before writing the number to indicate the ratio.

e.g., [SO4]2– → polyatomic radical

H1+ SO42– → H2SO4.

Chemical formula of some simple compounds

(a) Calcium hydroxid

(b) Aluminium oxide

• Molecular Mass

It is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass unit (u).

e.g., 2H+ + O2 H2O [H = 1, 0 = 16]

1 × 2 + 16 = 18 u

• Formula Unit Mass

It is the sum of the atomic masses of all atoms in a formula unit of a compound. The constituent particles are ions.

e.g., Na+ + Cl– → NaCl

1 × 23 + 1 × 35.5 = 58.5 u

• Mole Concept

Definition of mole: It is defined as one mole of any species (atom, molecules, ions or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.

1 mole = 6.022 × 1023 in number

Molar mass = mass of 1 mole → is is always expressed in r gram, and is also known as gram atomic mass.

1u of hydrogen has → 1 atom of hydrogen

lg of hydrogen has → 1 mole of hydrogen

= 6.022 × 1023 atoms of hydrogens

Class IX Science
NCERT Solution for Atoms and Molecules

IN-TEXT QUESTIONS SOLVED

NCERT TEXTBOOK PAGE 32

1. In a reaction 5.3 g of sodium carbonate reacted with 6 g of ethanoic acid. The products were 2.2 g of carbon dioxide, 0.9 g water and 8.2 g of sodium ethanoate. Show that these observations are in agreement with the law of conservation of mass carbonate.

Ans.

Sodium carbonate	+	Ethanoic acid	→	Sodium ethanoate	+	Carbon dioxide	+	Water
5.3 g	+	6 g	→	8.2 g	+	2.2 g	+	0.9 g

LHS RHS

11.3 g = 11.3 g

(Mass of reactant) (Mass of product)

This shows, that during a chemical reaction mass of reactact = mass of product.

2. 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?

Ans. Ratio of H : O by mass in water is:

Hydrogen : Oxygen → H2O

1 : 8 = 3 : x

x = 8 × 3

∴ 24 g of oxygen gas would be required to react completely with 3 g of hydrogen gas.

3. Which postulate of Dalton's atomic theory the result of the law of conservation of mass?

Ans. The postulate of Dalton's atomic theory that is the result of the law of conservation of mass is-the relative number and kinds of atoms are constant in a given compound. Atoms cannot be created nor destroyed in a chemical reaction.

4. Which postulate of Dalton's atomic theory can explain the law of definite proportions?

Ans. The relative number and kinds of atoms are constant in a given compound.

NCERT TEXTBOOK PAGE 35

1. Define the atomic mass unit.

Ans. One atomic mass unit is equal to exactly one-twelfth (1/12th) the mass of one atom of carbon-12. The relative atomic masses of all elements have been found with respect to an atom of carbon-12.

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

Ans. Atom is too small to be seen with naked eyes. It is measured in nanometres.

1 m = 109 nm

NCERT TEXTBOOK PAGE 35

1. Write down the formulae of

(i) Sodium oxide

(ii) Aluminium chloride

(iii) Sodium sulphide

(iv) Magnesium hydroxide

Ans. The formulae. are

(i) Formula bf Sodium Oxide

Symbol → Na O

Charge → +1 –2

Formula → Na2O

(iii) Formula of Sodium Oxide

Symbol → Na S

Charge → + 1 –2

Formula → Na2S

(ii) Formula of aluminium chloride

Symbol → Al Cl

Charge → +3 –1

Formula → A1C13

(iv) Formula of magnesium hydroxide

Symbol → Mg OH

Charge 1

Formula → Mg(OH)2

2. Write down the names of compounds represented btthe following formulae:

(i) Al2(SO4)3 (ii) CaCl2 (iii) K2SO4 (iv) KNO3

(v) CaCO3

Ans. (i) A12(SO4)3 → Aluminium sulphate

(ii) CaC12 → Calcium chloride

(iii) K∴SO4 → Potassium sulphate

(iv) KNO3 → Potassium nitrate

(v) CaCO3 → Calcium carbonate

3. What is meant by the term chemical formula?

Ans. The chemical formula of the compound is a symbolic representation of its composition, e.g., chemical formula of sodium chloride is NaCl.

4. How many atoms are present in a

(i) H2S molecule and

(ii) PO43– ion?

Ans. (i) H2S → 3 atoms are present

(ii) PO43– → 5 atoms are present

NCERT TEXTBOOK PAGE 40

1. Calculate the molecular masses of H2, O2, C12, CO2, CH4, C2H6, C2H4, NH3, CH3OH.

Ans. The molecular masses are:

H2 ⇒ 1 × 2 → 2u

O2⇒ 16 × 2 → 32u

Cl2 ⇒ 35.5 × 2 → 71 u

CO2 ⇒ 1 × 12 + 2 × 16 = 12 +32 = 44 u

CH4 ⇒ 1 × 12 + 4 × 1 = 16 u

C2H6 ⇒ï¿½n2 × 12 + 6 × 1 = 30 u

C2H4 ⇒ (2 × 12) + (4 × 1) = 28 u

NH3 ⇒ (1 × 14) + (3 × 1) = 17 u

CH3OH ⇒ 12 + (3 × 1) + 16 + 1 = 32 u

2. Calculate the formula unit masses of ZnO, Na2O, K2Co3 given atomic masses of Zn=65 u, Na = 23 u, K = 39 u, C = 12u, and O = 16 u.

Ans. The formula unit mass of

(i) ZnO = 65 u + 16 u =.81 u

(ii) Na2O = (23 u × 2) + .16 u : = 46u+ 161.u.

= 62 u

(iii) K2CO3 = (39 u × 2) + 12 u + 16 u × 3

= 75 u + 12u = 48u = 138u

NCERT TEXTBOOK PAGE 42

1. If one mole of carbon atoms weigh 12 grams, what is the mast (in grams) of 1 atom of carbon?

Ans. 1 mole of carbon atoms 6.022 × 1023 atoms = 12 g

Mass of 1 atom = ?

∴ Mass of 1 atom of carbon

=1.99 × 10–23 g

2. 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)?

Ans. 23 g of Na = 6.022 × 1023 atoms (1 mole).

100 g of Na = ?

= 26.182 ×1023 = 2.6182×1024 atoms

56 g of Fe = 6.022 × 1023 atoms

100 g o Fe = ?

100 g of Na contain → 2.618 × 1024 atoms

100 g of Fe contain → 1.075 × 1024 atoms

∴ 100 g of Na contains more atoms.

QUESTIONS FROM NCERT TEXTBOOK

1. 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.

Ans. Boron and oxygen compound → Boron + Oxygen

0.24 g → 0:096g + 0.144g

Percentage composition of the compound

For boron: 0.24g → 0.096 g

100 g → ?

For oxygen:

0.24 g → 0.144 g of oxygen

100 g → ?

2. 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 formed when. 3.00 g of carbon is burnt in 50.00 g of oxygen? Which law of chemicai combination will govern your answer?

Ans. The reaction of burning of carbon in oxygen may be written as:

It shows that 12 g of carbon burns in 32 g oxygen to form 44 g of carbon dioxide. Therefore 3 g of carbon reacts with 8 g of oxygen to form 11 g of carbon dioxide. It is given that 3.0 g of carbon is burnt with 8 g of oxygen to produce 11.0 g of CO2. Consequently 11.0 g of carbon dioxide will be formed when 3.0 g of C is burnt in 50 g of oxygen consuming 8 g of oxygen, leaving behind 50 ï¿½V 8 = 42 g of 02. The answer governs the law of constant proportion.

3. What are polyatomic ions? Give eaamples:

Ans. The ions which contain more than one atoms (same kind or may be of different kind) and behave as a single unit are called polyatomic ions e.g., OH–, SO2–4, CO32–.

4. Write the chemical formulae of the following:

(a) Magnesium chloride (b) Caldum oxide

(e) Calcium carbonate.

Ans. (a) Magnesium chloride

Symbol → Mg Cl

Change → +2 –1

Formula → MgCl2

(b) Calcium oxide

Symbol → Ca O

Charge → +2 –2

Formula → CaO

Symbol → Cu NO3

Change → +2 –1

Formula → Cu(NO3)2

(d) Aluminium chloride

Symbol → Al Cl

Change → +3 –1

Formula → A1Cl3

(d) Calcium carbonate

Symbol → Ca CO3

Change → +2 –2

Formula → CaCO3

5. Give the names of the elements present in the following compounds:

(a) Quick time (b) Hydrogen bromide

Ans. (a) Quick lime → Calcium oxide

Elements → Calcium and oxygen

(b) Hydrogen bromide

Elements → Hydrogen and bromine

Elements → Calcium, hydrogen, carbon and oxygen

(d) Potassium sulphate

Elements → Potassium, sulphur and oxygen

6. Calculate the molar mass of the following substances.

(a) Ethyne, C2H2

(b) Sulphur molecule, S8

(d) Hydrochloric acid, HCl

(e) Nitric acid, HNO3

Ans. The molar mass of the following: [Unit is ‘g’]

(a) Ethyne, C2H2 = 2 × 12 + 2 × 1 = 24 + 2 = 26 g

(b) Sulphur molecule, S8 = 8 × 32 = 256 g

(d) Hydrochloric acid, HCl = 1 × 1 + 1 × 35.5

= 1 + 35.5 = 36.5 g

(e) Nitric acid, HNO3 = 1 × 1 + 1 × 14 + 3 × 16

= 1 + 14 + 48 = 63 g

7. What is the mass of

(a) 1 mole of nitrogen atoms?

(b) 4 moles of aluminium atoms

Ans. (a) Mass of 1 mole of nitrogen atoms = 14g

(b) 4 moles of aluminium atoms

Mass of 1 mole of aluminium atoms = 27 g

Mass of 4 moles of aluminium atoms = 27 × 4 = 108 g

Mass of 1 mole of Na2SO3 = 2 × 23 + 32 + 3 ×16

=46 + 32 + 48 = 126 g

Mass of 10 moles of Na2SO3 = 126 × 10

= 1260 g

8. Convert into mole.

(a) 12 g of oxygen gas

(b) 20 g of water

Ans. (a) Given mass of oxygen gas = 12 g

Molar mass of oxygen gas (O2) = 32 g

Mole of oxygen gas 12

mole

(b) Given mass of water = 20 g

Molar mass of water (H2O) = (2 × 1) + 16 = 18 g

Mole of water

mole

Molar mass of carbon dioxide (CO2) = (1 × 12) + (2 × 16)

= 12 + 32 = .44 g

Mole of carbon dioxide

mole

9. What is the mass of

(a) 0.2 mole of oxygen atoms?

(b) 0.5 mole of water molecules?

Ans. (a) Mole of Oxygen atoms = 0.2 mole

Molar mass of oxygen atoms = 16 g

Mass of oxygen atoms = 16 × 0.2 = 3.2 g

(b) Mole of water molecule = 0.5 mole

Molar mass of water molecules = 2 × 1 + 16 = 18 g

Mass of H2O = 18 × 0.5 = 9 g

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

Ans. Molar mass of S8 sulphur = 256 g = 6.022 × 1023 molecule

Given mass of sulphur = 16 g

= 0.376 × 1023

= 3.76 × 1022 molecules

11. Calculate the number of aluminium ions present in 0.051 g of aluminium oxide.

Ans. Molar mass of aluminium oxide Al2O3

= (2 × 27) + (3 × 16)

= 54 + 48 = 102g.

102 g of Al2O3 contains = 2 × 6.022 × 1023 aluminium ions

= 0.006022 × 1023

= 6.022 × 1020 Al3+ ions

MULTIPLE CHOICE QUESTIONS
ATOMS AND MOLECULES

1. The atomicity of K2Cr2O7 is

I. 9

II. 11

III. 10

IV. 12

2. The formula for quicklime is

I. CaCl2

II. CaCo3

III. Ca(OH)2

IV. CaO

3. The symbol of cadmium is

I. Ca

II. Cu

III. Cm

IV. Cd

4. All noble gas molecules are

I. Monoatomic

II. Diatomic

III. Triatomic

IV. Both I and II

5. The valency of nitrogen in NH3 is

I. 1

II. 3

III. 4

IV. 5

6. The formula of ethanol is C2H5 – OH. What will be its molecular mass?

I. 46 u

II. 34 u

III. 34 g

IV. 46 g

7. Number of moles present in 28g of nitrogen atoms are

I. 1 mole

II. 2.3 moles

III. 0.5 mole

IV. 2 moles

8. The molecular mass of x is 106. x can be

I. CaCo3

II. So3

III. Na2Co3

IV. NaCl

9. Which among the following is not a postulate of Daltonï¿½s atomic theory?

I. Atoms can not be created or destroyed

II. Atoms of different elements have different sizes, masses and chemical properties

III. Atoms of same elements can combine in only one ratio to produce more than any one compound

IV. Atoms are very tiny particles which can not be further divided

10. Which of the following is a wrong Combination?

I. 6.022 * 1023 molecules of oxygen = 32g of oxygen

II. 6.022 * 1023 ions of sodium = 23g of sodium

III. 6.022 * 1023atoms of C = 24g of carbon

IV. 6.022 * 1023 atoms of H = 1g of hydrogen atoms

SAMPLE PAPER
ATOMS AND MOLECULES

1. Explain the molecular mass of C2H5OH.

2. Explain the law of constant proportion.

3. Explain the difference of O2 and 2O.

4. Find the number of moles in 7g of Na.

5. What is the atomicity of Ca(OH)2?

6. Write the formula for Aluminium Chloride.

7. State the difference between sodium atom and sodium ion.

8. What is formula unit mass? How is it different from molecular mass?

9. Define valency and give valency of copper and iron.

10. Calculate the mass of one molecule of chlorine.

HIGH ORDER THINKING SKILLS
ATOMS AND MOLECULES

1. Define polyatomic ion. Give one example.

2. What is formula unit of mass? How is it different from molecular mass?

3. What is Law of conservation of mass and Law of constant proportions?

4. What is an ion? Explain the types of ion with examples.

5. Find the molecular mass of H2O.

6. Define the term valency. What is the valency for magnesium and copper?

7. What is the difference between cation and anion?

8. What is atomicity? What is the atomicity of phosphorus and nitrogen?

9. Find the number of atoms in 0.5 mole of C atom.

10. Find the mass of 1.5 mole of CO2 molecule.

11. Calculate the formula unit mass of NaCl and CaCl2.

12. What is the difference between molecules 2O and O2?

TOMS AND MOLECULES TEST

Maximum marks- 25

Maximum time-35 minutes

1. What are the rules for writing the symbol of an element? (3)

2. Explain relative atomic mass and relative molecular mass. (2)

3. What is the relationship between mole, Avogadro number and mass? (3)

4. How do atoms exist? (2)

5. Define a mole. What is the importance of a mole? (3)

6. Calculate the mass of one atom of oxygen (1.5)

7. Calculate the mass of one ion of oxygen (1.5)

8. What do you understand by atomic mass and Gramatomic mass of an element? (2)

9. The formula of water is H2O. What do you understand by this formula? (3)

10. State the Law of conservation of mass and the Law of constant proportion with examples. (4)

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