Chapter 3 — Atoms and Molecules
Welcome to HSLC Guru, dear ASSEB Class 9 students! In this lesson, you will explore the building blocks of all matter — atoms and molecules. This complete English-medium guide includes a clear summary of the laws of chemical combination, Dalton’s atomic theory, the mole concept, and writing of chemical formulas, followed by detailed textbook Question and Answer solutions, additional MCQs, fill in the blanks, true/false statements, a glossary and a formula table prepared strictly according to the Assam State School Education Board (ASSEB) syllabus.
Summary
The chapter Atoms and Molecules begins with two fundamental laws of chemical combination proposed by Antoine Lavoisier and Joseph Proust. The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. The total mass of the reactants is always equal to the total mass of the products. The Law of Constant Proportions (also called Law of Definite Proportions) states that in a pure chemical compound, the elements are always present in a definite proportion by mass. For example, in pure water (H₂O), hydrogen and oxygen are always present in the ratio of 1:8 by mass, no matter what the source of water is.
To explain these laws, John Dalton in 1808 proposed his famous Atomic Theory. According to it: (i) all matter is made of very tiny particles called atoms, (ii) atoms are indivisible particles that cannot be created or destroyed in a chemical reaction, (iii) atoms of a given element are identical in mass and chemical properties, (iv) atoms of different elements have different masses and properties, (v) atoms combine in small whole-number ratios to form compounds, and (vi) the relative number and kinds of atoms in a given compound are constant. An atom is the smallest particle of an element that takes part in a chemical reaction. Each element is represented by a symbol — usually one or two letters of its English or Latin name (for example, H for Hydrogen, Na for Sodium from Natrium, Fe for Iron from Ferrum). The atomic mass of an element is the average mass of its atom compared to 1/12 of the mass of one carbon-12 atom and is expressed in atomic mass units (u).
A molecule is the smallest particle of an element or a compound that can exist independently and shows all the properties of that substance. The number of atoms present in one molecule of an element is called its atomicity. For example, helium (He) is monoatomic, oxygen (O₂) is diatomic, ozone (O₃) is triatomic, phosphorus (P₄) is tetratomic and sulphur (S₈) is polyatomic. The molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance. When atoms or groups of atoms carry an electric charge, they are called ions. Positively charged ions are called cations (e.g. Na⁺, Ca²⁺) and negatively charged ions are called anions (e.g. Cl⁻, O²⁻). A group of atoms carrying a charge is called a polyatomic ion, such as ammonium (NH₄⁺), sulphate (SO₄²⁻), nitrate (NO₃⁻) and carbonate (CO₃²⁻). The combining capacity of an element is called its valency.
A chemical formula is the symbolic representation of the composition of a compound. While writing a formula, the valencies or charges of the constituent ions must be balanced by criss-crossing them. The mole concept was introduced to count very large numbers of atoms and molecules. One mole of any substance contains 6.022 × 10²³ particles — this number is known as Avogadro’s Number (N₀). The mass of one mole of a substance, in grams, is equal to its atomic or molecular mass and is called the molar mass. Thus, 1 mole of carbon = 12 g = 6.022 × 10²³ atoms, and 1 mole of water = 18 g = 6.022 × 10²³ molecules. The mole concept allows scientists to calculate the number of particles, mass, or volume of a substance involved in a chemical reaction with great accuracy.
Textbook Questions and Answers
A. Very Short Answer Type Questions (1 Mark)
Q1. Who proposed the law of conservation of mass?
Answer: Antoine Lavoisier (1774) proposed the law of conservation of mass.
Q2. What is the value of Avogadro’s number?
Answer: The value of Avogadro’s number (N₀) is 6.022 × 10²³ particles per mole.
Q3. Write the chemical symbol of Sodium and Potassium.
Answer: Sodium = Na (from Latin Natrium); Potassium = K (from Latin Kalium).
Q4. Define atomicity. Give an example.
Answer: The number of atoms present in one molecule of an element is called atomicity. For example, the atomicity of oxygen (O₂) is 2.
Q5. What is meant by 1 u (atomic mass unit)?
Answer: One atomic mass unit (1 u) is defined as exactly 1/12 of the mass of one atom of carbon-12.
Q6. Name two cations and two anions.
Answer: Cations: Na⁺ (sodium ion), Ca²⁺ (calcium ion). Anions: Cl⁻ (chloride ion), O²⁻ (oxide ion).
Q7. Write the formula of ammonium sulphate.
Answer: Ammonium sulphate = (NH₄)₂SO₄.
Q8. What is the molar mass of water (H₂O)?
Answer: Molar mass of H₂O = (2 × 1) + 16 = 18 g/mol.
Q9. Define valency.
Answer: Valency is the combining capacity of an element, that is, the number of electrons an atom of an element can lose, gain or share to attain a stable configuration.
Q10. Give one example of a polyatomic ion.
Answer: Sulphate ion (SO₄²⁻) is an example of a polyatomic ion.
B. Short Answer Type Questions (2-3 Marks)
Q1. State the law of conservation of mass with one example.
Answer: The law of conservation of mass states that “mass can neither be created nor destroyed in a chemical reaction.” In other words, in a chemical change, the total mass of the reactants is equal to the total mass of the products.
Example: When 5.3 g of sodium carbonate reacts with 6 g of acetic acid, 2.2 g of carbon dioxide, 0.9 g of water and 8.2 g of sodium acetate are produced. Total reactants (5.3 + 6) = 11.3 g = total products (2.2 + 0.9 + 8.2) = 11.3 g.
Q2. Write the postulates of Dalton’s atomic theory.
Answer: The main postulates of Dalton’s atomic theory are:
- All matter is made up of very small particles called atoms.
- Atoms are indivisible and 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 small whole-number ratios to form compounds.
- The relative number and kinds of atoms in a given compound are constant.
Q3. Differentiate between atom and molecule.
| Atom | Molecule |
|---|---|
| Smallest particle of an element. | Smallest particle of an element or compound that exists independently. |
| Generally cannot exist freely (except noble gases). | Can exist independently in nature. |
| Made of subatomic particles. | Made of two or more atoms joined together. |
| Example: H, O, N. | Example: H₂O, O₂, CO₂. |
Q4. Why is it not possible to see an atom with naked eyes?
Answer: An atom is extremely small in size, with a radius of about 10⁻¹⁰ m (0.1 nm). It is much smaller than the wavelength of visible light, so it cannot be seen even with the most powerful optical microscope. Hence, atoms cannot be seen with naked eyes.
Q5. Define molecular mass and write the molecular mass of CO₂.
Answer: The molecular mass of a substance is the sum of the atomic masses of all the atoms present in one molecule of that substance, expressed in atomic mass units (u).
Molecular mass of CO₂ = 12 + (2 × 16) = 12 + 32 = 44 u.
Q6. Write the rules for writing chemical formulas of simple compounds.
Answer: The main rules are:
- The valencies or charges of the ions must balance.
- When a compound contains a metal and a non-metal, the symbol of the metal is written first.
- In compounds formed with polyatomic ions, the number of ions is shown by enclosing the formula of the ion in brackets and writing the number outside, e.g. Mg(OH)₂.
- The valencies are criss-crossed and written as subscripts.
C. Long Answer / Numerical Questions (5-6 Marks)
Q1. Calculate the molecular mass of (a) H₂O, (b) CO₂, (c) CH₄, (d) NH₃ and (e) C₂H₆.
Answer:
- (a) H₂O = (2 × 1) + 16 = 18 u
- (b) CO₂ = 12 + (2 × 16) = 44 u
- (c) CH₄ = 12 + (4 × 1) = 16 u
- (d) NH₃ = 14 + (3 × 1) = 17 u
- (e) C₂H₆ = (2 × 12) + (6 × 1) = 30 u
Q2. Calculate the number of moles in (a) 22 g of CO₂, (b) 18 g of water and (c) 100 g of CaCO₃.
Answer: Number of moles = Given mass ÷ Molar mass.
- (a) Moles of CO₂ = 22 ÷ 44 = 0.5 mol
- (b) Moles of H₂O = 18 ÷ 18 = 1 mol
- (c) Molar mass of CaCO₃ = 40 + 12 + (3 × 16) = 100 g/mol; Moles = 100 ÷ 100 = 1 mol
Q3. Calculate the number of molecules of sulphur (S₈) present in 16 g of solid sulphur.
Answer:
Molar mass of S₈ = 8 × 32 = 256 g/mol.
Number of moles = 16 ÷ 256 = 0.0625 mol.
Number of molecules = 0.0625 × 6.022 × 10²³ = 3.76 × 10²² molecules.
Q4. What is the mass of (a) 0.5 mole of N₂ gas, (b) 0.5 mole of N atoms, (c) 3.011 × 10²³ atoms of carbon, (d) 6.022 × 10²³ molecules of oxygen?
Answer:
- (a) Mass of 0.5 mol N₂ = 0.5 × 28 = 14 g
- (b) Mass of 0.5 mol N atoms = 0.5 × 14 = 7 g
- (c) Number of moles of C = 3.011 × 10²³ ÷ 6.022 × 10²³ = 0.5 mol; Mass = 0.5 × 12 = 6 g
- (d) 6.022 × 10²³ molecules of O₂ = 1 mole = 32 g
Q5. A 0.24 g sample of a 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 mass.
Answer:
Total mass of compound = 0.24 g.
Percentage of Boron = (0.096 ÷ 0.24) × 100 = 40 %.
Percentage of Oxygen = (0.144 ÷ 0.24) × 100 = 60 %.
Additional Multiple Choice Questions (MCQs)
Q1. The law of conservation of mass was given by:
(a) Proust (b) Lavoisier (c) Dalton (d) Avogadro
Answer: (b) Lavoisier.
Q2. The atomic mass of an element is expressed in:
(a) gram (b) kilogram (c) atomic mass unit (u) (d) milligram
Answer: (c) atomic mass unit (u).
Q3. The chemical symbol of iron is:
(a) Ir (b) In (c) Fe (d) I
Answer: (c) Fe.
Q4. The molecular mass of methane (CH₄) is:
(a) 14 u (b) 16 u (c) 18 u (d) 32 u
Answer: (b) 16 u.
Q5. One mole of any substance contains:
(a) 6.022 × 10²² particles (b) 6.022 × 10²³ particles (c) 6.022 × 10²⁴ particles (d) 6.022 × 10²⁵ particles
Answer: (b) 6.022 × 10²³ particles.
Q6. Which of the following is a polyatomic ion?
(a) Na⁺ (b) Cl⁻ (c) NH₄⁺ (d) K⁺
Answer: (c) NH₄⁺.
Q7. The atomicity of phosphorus (P₄) is:
(a) 1 (b) 2 (c) 3 (d) 4
Answer: (d) 4.
Q8. The mass of one mole of oxygen atoms is:
(a) 8 g (b) 16 g (c) 32 g (d) 18 g
Answer: (b) 16 g.
Q9. The formula of calcium chloride is:
(a) CaCl (b) CaCl₂ (c) Ca₂Cl (d) Ca₂Cl₃
Answer: (b) CaCl₂.
Q10. Which scientist is known as the “Father of modern Atomic Theory”?
(a) Lavoisier (b) Avogadro (c) Dalton (d) Proust
Answer: (c) Dalton.
Fill in the Blanks
Q1. The law of constant proportions was given by ____________.
Answer: Joseph Proust.
Q2. One atomic mass unit is 1/12th of the mass of one atom of ____________.
Answer: carbon-12.
Q3. The number 6.022 × 10²³ is called ____________.
Answer: Avogadro’s number.
Q4. Positively charged ions are called ____________.
Answer: cations.
Q5. The molar mass of CO₂ is ____________ g.
Answer: 44.
True or False
Q1. Atoms can be created and destroyed in a chemical reaction.
Answer: False. Atoms are neither created nor destroyed in a chemical reaction.
Q2. The molecular mass of water is 18 u.
Answer: True.
Q3. Helium is a diatomic gas.
Answer: False. Helium is a monoatomic gas.
Q4. The valency of oxygen is 2.
Answer: True.
Q5. 1 mole of CO₂ contains 6.022 × 10²³ atoms.
Answer: False. 1 mole of CO₂ contains 6.022 × 10²³ molecules (and 3 × 6.022 × 10²³ atoms in total).
Glossary
| Term | Meaning |
|---|---|
| Atom | The smallest particle of an element that takes part in a chemical reaction. |
| Molecule | The smallest particle of a substance that can exist independently. |
| Atomicity | The number of atoms present in one molecule of an element. |
| Atomic Mass | The mass of one atom of an element compared to 1/12th the mass of a C-12 atom. |
| Molecular Mass | The sum of the atomic masses of all atoms in a molecule. |
| Mole | The amount of substance containing 6.022 × 10²³ particles. |
| Avogadro’s Number | The number of particles in one mole = 6.022 × 10²³. |
| Cation | A positively charged ion (e.g. Na⁺). |
| Anion | A negatively charged ion (e.g. Cl⁻). |
| Polyatomic Ion | An ion made up of two or more atoms (e.g. SO₄²⁻). |
| Valency | The combining capacity of an element. |
| Chemical Formula | A symbolic representation of the composition of a compound. |
Important Formulas
| Quantity | Formula |
|---|---|
| Number of moles (n) from mass | n = Given mass (m) ÷ Molar mass (M) |
| Number of moles from particles | n = Number of particles ÷ 6.022 × 10²³ |
| Number of particles | N = n × 6.022 × 10²³ |
| Mass of substance | m = n × M |
| Molecular mass | Sum of atomic masses of all atoms in the molecule |
| Percentage of element | % = (Mass of element ÷ Total mass) × 100 |
| Avogadro’s Number (N₀) | 6.022 × 10²³ per mole |
| Molar Volume of gas at STP | 22.4 L (or 22,400 mL) |
That completes the detailed solutions, additional practice questions, glossary and formula sheet for Chapter 3 — Atoms and Molecules from the ASSEB Class 9 Science textbook. Keep practising the mole concept numericals, as they form the foundation for all higher chemistry topics. Stay connected with HSLC Guru for more chapter-wise English-medium notes and solutions.