Atoms and Molecules

Detailed NCERT Solutions & Analysis

In-Text Questions (Page 32)

Q1. In a reaction, 5.3 g of sodium carbonate reacted with 6 g of acetic 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.

Answer:

Mass of Reactants:
Mass of Sodium Carbonate + Mass of Acetic Acid
= 5.3 g + 6 g = 11.3 g

Mass of Products:
Mass of Sodium Ethanoate + Mass of CO2 + Mass of Water
= 8.2 g + 2.2 g + 0.9 g = 11.3 g

Since, Mass of Reactants = Mass of Products, the observations are in agreement with the Law of Conservation of Mass.

Q2. 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:

Ratio of Hydrogen : Oxygen = 1 : 8
This means 1 g of Hydrogen requires 8 g of Oxygen.
Therefore, 3 g of Hydrogen will require = 3 × 8 = 24 g of Oxygen.

In-Text Questions (Page 33)

Q1. Which postulate of Dalton's atomic theory is the result of the law of conservation of mass?

Postulate: "Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction."

Q2. Which postulate of Dalton's atomic theory can explain the law of definite proportions?

Postulate: "The relative number and kinds of atoms are constant in a given compound."

In-Text Questions (Page 35)

Q1. Define the atomic mass unit.

Atomic mass unit (amu or u) is defined as a mass unit equal to exactly one-twelfth (1/12th) the mass of one atom of carbon-12.

In-Text Questions (Page 39)

Q1. Write down the formulae of:

(i) Sodium Oxide: Na₂O
(ii) Aluminium Chloride: AlCl₃
(iii) Sodium Sulphide: Na₂S
(iv) Magnesium Hydroxide: Mg(OH)₂

In-Text Questions (Page 40)

Q1. Calculate the molecular mass of H2, O2, Cl2, CO2, CH4, C2H6, C2H4, NH3, CH3OH.

Given: Atomic masses: H=1, O=16, Cl=35.5, C=12, N=14.

H₂ = 2 × 1 = 2 u
O₂ = 2 × 16 = 32 u
Cl₂ = 2 × 35.5 = 71 u
CO₂ = 12 + (2 × 16) = 12 + 32 = 44 u
CH₄ = 12 + (4 × 1) = 16 u
C₂H₆ = (2 × 12) + (6 × 1) = 24 + 6 = 30 u
C₂H₄ = (2 × 12) + (4 × 1) = 24 + 4 = 28 u
NH₃ = 14 + (3 × 1) = 17 u
CH₃OH = 12 + (3 × 1) + 16 + 1 = 32 u

In-Text Questions (Page 42)

Q1. If one mole of carbon atoms weighs 12 gram, what is the mass (in grams) of 1 atom of carbon?

1 mole of Carbon = 6.022 × 10²³ atoms = 12 g
Mass of 1 atom = 12 / (6.022 × 10²³)
= 1.99 × 10^-23 g

Q2. Which has more number of atoms, 100 grams of sodium or 100 grams of iron?

Sodium (Na): Atomic mass = 23 u
Moles of Na = 100 / 23 = 4.35 moles

Iron (Fe): Atomic mass = 56 u
Moles of Fe = 100 / 56 = 1.78 moles

Since moles of Na > moles of Fe, 100 g of Sodium has more atoms.

Main Textbook Exercises

Q1. 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.

Given: Total mass = 0.24 g, Boron = 0.096 g, Oxygen = 0.144 g

% of Boron = (0.096 / 0.24) × 100 = 40%

% of Oxygen = (0.144 / 0.24) × 100 = 60%

Q6. Calculate the molar mass of the following substances.

(a) Ethyne, C₂H₂ = (2 × 12) + (2 × 1) = 24 + 2 = 26 g
(b) Sulphur molecule, S₈ = 8 × 32 = 256 g
(c) Phosphorus molecule, P₄ = 4 × 31 = 124 g
(d) Hydrochloric acid, HCl = 1 + 35.5 = 36.5 g
(e) Nitric acid, HNO₃ = 1 + 14 + (3 × 16) = 1 + 14 + 48 = 63 g

Atoms and Molecules

Detailed Chapter Analysis & The Mole Concept

1. Laws of Chemical Combination

The ancient maharishi, Kanad, postulated that if we go on dividing matter (padarth), we shall get smaller and smaller particles. A stage will come when we shall come across the smallest particles beyond which further division will not be possible. He named these particles 'Parmanu'.

Later, Antoine L. Lavoisier laid the foundation of chemical sciences by establishing two important laws of chemical combination.

Law of Conservation of Mass

Mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of the reactants is equal to the total mass of the products.

Law of Constant Proportions

In a chemical substance, the elements are always present in definite proportions by mass. Example: In water (H₂O), the ratio of the mass of hydrogen to the mass of oxygen is always 1:8, whatever the source of water.

2. Dalton's Atomic Theory

According to Dalton's atomic theory, all matter, whether an element, a compound or a mixture is composed of small particles called atoms.

The Postulates

All matter is made of very tiny particles called atoms.
Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
Atoms of a given element are identical in mass and chemical properties.
Atoms of different elements have different masses and chemical properties.
Atoms combine in the ratio of small whole numbers to form compounds.
The relative number and kinds of atoms are constant in a given compound.

3. Ions and Formulae

Compounds composed of metals and nonmetals contain charged species known as ions. A negatively charged ion is called an anion and the positively charged ion is called a cation.

Rules for Writing Formulae:

The valencies or charges on the ion must balance.
When a compound consists of a metal and a non-metal, the name or symbol of the metal is written first. For example: calcium oxide (CaO), sodium chloride (NaCl).
In compounds formed with polyatomic ions, the ion is enclosed in a bracket before writing the number to indicate the ratio. Ex: Mg(OH)₂.

Valency

The combining power (or capacity) of an element is known as its valency. It can be used to find out how the atoms of an element will combine with the atom(s) of another element to form a chemical compound.

4. The Mole Concept

Since atoms and molecules are extremely small, we use a unit called Mole to count them. One mole is the amount of substance that contains the same number of particles (atoms/ions/molecules) as there are atoms in exactly 12 g of carbon-12.

Avogadro Constant (N_A)

6.022 × 10²³

This is the number of particles in 1 mole of any substance.

Molar Mass

The mass of 1 mole of a substance in grams. It is numerically equal to atomic/molecular mass in u.

Relationship Formulae

1. Number of Moles (from Mass) = Given Mass / Molar Mass

2. Number of Moles (from Particles) = Given Number of Particles / Avogadro Number

3. Mass = Moles × Molar Mass

Key Facts & Definitions

50+ Important Points to Remember

1. Parmanu

Indian philosopher Kanad's term for the smallest indivisible particle of matter.

2. Atom

Greek term for indivisible. The building block of all matter.

3. Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

4. Law of Constant Proportions

In a chemical substance, elements are always present in definite proportions by mass.

5. Antoine L. Lavoisier

Father of modern chemistry; established the laws of chemical combination.

6. Dalton's Atomic Theory

Proposed that all matter is composed of atoms.

7. Atom Radius Unit

Measured in nanometers (nm). 1 nm = 10^-9 m.

8. Hydrogen Atom Radius

Smallest of all atoms, radius is 10^-10 m.

9. IUPAC

International Union of Pure and Applied Chemistry - approves names of elements.

10. Chemical Symbol

First one or two letters of the element's name. First letter uppercase, second lowercase.

11. Atomic Mass Unit (u)

Mass unit exactly equal to 1/12th the mass of one atom of carbon-12.

12. Relative Atomic Mass

Average mass of an atom as compared to 1/12th mass of one C-12 atom.

13. Molecule

Smallest particle of a compound capable of independent existence.

14. Atomicity

The number of atoms constituting a molecule.

15. Diatomic

Molecules having 2 atoms. Ex: O₂, H₂, N₂.

16. Tetra-atomic

Molecules having 4 atoms. Ex: Phosphorus (P₄).

17. Polyatomic

Molecules having more than 4 atoms. Ex: Sulphur (S₈).

18. Ion

A charged particle (positive or negative).

19. Cation

Positively charged ion (loses electrons). Ex: Na⁺.

20. Anion

Negatively charged ion (gains electrons). Ex: Cl^-.

21. Polyatomic Ion

A group of atoms carrying a charge. Ex: Sulphate (SO₄^2-).

22. Valency

The combining capacity of an element.

23. Binary Compounds

The simplest compounds separated by two different elements.

24. Chemical Formula

The symbolic representation of the composition of a compound.

25. Molecular Mass

Sum of atomic masses of all atoms in a molecule.

26. Formula Unit Mass

Sum of atomic masses for ionic compounds (e.g., NaCl).

27. Mole

One mole is the amount of substance having 6.022 × 10²³ particles.

28. Avogadro Number (N_A)

6.022 × 10²³.

29. Molar Mass

Mass of 1 mole of a substance in gram. Numerical value is same as atomic mass, unit is 'g'.

30. Gram Atomic Mass

Atomic mass expressed in grams (mass of 1 mole atoms).

31. 1 Mole of H atoms

Contains 6.022 × 10²³ atoms of H and weighs 1 g.

32. 1 Mole of Water molecules

Weighs 18 g and contains 6.022 × 10²³ molecules of H₂O.

33. Mole - Latin origin

Means 'heap' or 'pile'. Coined by Wilhelm Ostwald.

34. Mass of Electron

Negligible compared to Proton and Neutron.

35. H₂O ratio by mass

1 : 8 (Hydrogen : Oxygen).

36. NH₃ ratio by mass

14 : 3 (Nitrogen : Hydrogen).

37. CO₂ ratio by mass

3 : 8 (Carbon : Oxygen).

38. Ferrum

Latin name for Iron (Symbol Fe).

39. Natrium

Latin name for Sodium (Symbol Na).

40. Kalium

Latin name for Potassium (Symbol K).

41. Valency of Aluminium

+3.

42. Valency of Carbon

4. It is tetravalent.

43. Sulphate Ion

SO₄^2-.

44. Carbonate Ion

CO₃^2-.

45. Ammonium Ion

NH₄⁺.

46. Hydroxide Ion

OH^-.

47. Phosphate Ion

PO₄^3-.

48. Quick Lime

Calcium Oxide (CaO).

49. Baking Soda

Sodium Hydrogen Carbonate (NaHCO₃).

50. Limestone

Calcium Carbonate (CaCO₃).

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