NCERT Class 9 Science Chapter 3: Atoms and Molecules - A Comprehensive Guide

Introduction to Atoms and Molecules

Welcome to this comprehensive guide on Chapter 3 of the NCERT Class 9 Science textbook, "Atoms and Molecules." This chapter forms the very bedrock of chemistry. Everything we see, touch, and feel is made up of matter, but what is matter itself made of? Ancient Indian and Greek philosophers pondered this question for centuries. The Indian philosopher Maharishi Kanad postulated that if we go on dividing matter (padarth), we shall get smaller and smaller particles, eventually reaching a stage where further division is impossible. He named these indivisible particles 'Parmanu'. Another Indian philosopher, Pakudha Katyayama, elaborated on this doctrine and said that these particles normally exist in a combined form which gives us various forms of matter.

Around the same era, ancient Greek philosophers Democritus and Leucippus also suggested that if we go on dividing matter, a stage will come when particles obtained cannot be divided further. Democritus called these indivisible particles 'atoms' (meaning indivisible). However, these were philosophical ideas without much experimental evidence. It wasn't until the late 18th century that scientists began to use experimental methods to understand the nature of matter and the transformations it undergoes. This chapter will take you on a journey from these early ideas to the scientifically established laws and theories that govern the world of atoms and molecules. Understanding these concepts is crucial as they provide the foundation for almost all other topics in chemistry.

Laws of Chemical Combination

By the end of the 18th century, scientists recognized the difference between elements and compounds and became interested in how and why elements combine and what happens when they do. The meticulous experimental work of Antoine L. Lavoisier and Joseph L. Proust laid the foundation of chemical sciences by establishing two fundamental laws of chemical combination.

Law of Conservation of Mass

This is one of the most fundamental principles in all of science. It deals with the mass of reactants and products in a chemical reaction.

Statement: The law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction. In other words, the total mass of the reactants is equal to the total mass of the products in a chemical reaction.

Explanation and Experiment (Lavoisier's Contribution): Antoine Lavoisier conducted careful experiments where he carried out chemical reactions in closed containers. He meticulously measured the mass of the reactants before the reaction and the mass of the products after the reaction. He found that in every case, the total mass remained unchanged.

An Illustrative Activity: Imagine you take a solution of sodium sulphate (Na₂SO₄) in a conical flask and a solution of barium chloride (BaCl₂) in a small ignition tube. You hang the ignition tube carefully in the flask so that the solutions do not mix. You put a cork on the flask and weigh it. Let's say the initial mass is 'x' grams. Now, you tilt and swirl the flask so that the two solutions mix. A chemical reaction occurs:

BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)

You will observe the formation of a white, insoluble substance called barium sulphate (a precipitate). After the reaction is complete, you weigh the flask again. You will find that the final mass is still 'x' grams. This experiment beautifully demonstrates that the total mass before and after the reaction remains constant, thus verifying the law of conservation of mass.

Law of Constant Proportions

This law, also known as the Law of Definite Proportions, was put forth by the French chemist Joseph Proust. It deals with the composition of a chemical compound.

Statement: The law of constant proportions states that in a chemical substance, the elements are always present in definite proportions by mass, irrespective of the source from which it is obtained or the method by which it is prepared.

Explanation and Example: Let's take the familiar example of water (H₂O). Water can be obtained from various sources – a river, a well, rain, or it can be synthesized in a laboratory by burning hydrogen in oxygen. Regardless of its origin, if you decompose 18g of pure water, you will always get 2g of hydrogen and 16g of oxygen. The ratio of the mass of hydrogen to the mass of oxygen in water is always:

Mass of Hydrogen : Mass of Oxygen = 2 : 16 = 1 : 8

This ratio of 1:8 by mass is constant for water. If you have 90g of water, it will contain (1/9) * 90g = 10g of hydrogen and (8/9) * 90g = 80g of oxygen. The ratio remains 10:80, which simplifies to 1:8. This consistency is the essence of the Law of Constant Proportions. Similarly, in ammonia (NH₃), the ratio of the mass of nitrogen to the mass of hydrogen is always 14:3.

What is an Atom? Dalton's Atomic Theory

The laws of chemical combination provided a framework, but they didn't explain the underlying nature of matter that led to these laws. The next logical step was to develop a theory that could explain these observations. British chemist and schoolteacher John Dalton provided the first comprehensive scientific theory of matter.

John Dalton's Contribution

John Dalton, born in a poor weaver's family in England in 1766, was a brilliant mind. He picked up the ancient Greek idea of 'atoms' and built a robust scientific theory around it. His theory was a turning point in the study of matter. He proposed that all matter, whether an element, a compound, or a mixture, is composed of very small particles called atoms. His theory provided a powerful explanation for the laws of chemical combination.

Postulates of Dalton's Atomic Theory

Dalton's atomic theory can be summarized in the following postulates:

• All matter is made of very tiny particles called atoms, which participate in chemical reactions. This explains the particulate nature of matter.
• Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction. This postulate is a direct explanation for the Law of Conservation of Mass. If atoms are simply rearranged during a reaction and not created or destroyed, the total mass must remain constant.
• Atoms of a given element are identical in mass and chemical properties. For example, every atom of gold (Au) is identical to every other atom of gold in terms of its mass and how it behaves chemically.
• Atoms of different elements have different masses and chemical properties. An atom of gold is fundamentally different from an atom of iron (Fe) or an atom of oxygen (O).
• Atoms combine in the ratio of small whole numbers to form compounds. This postulate explains the Law of Constant Proportions. For example, in water (H₂O), two atoms of hydrogen always combine with one atom of oxygen. Because atoms have fixed masses, their combination in a fixed numerical ratio results in a fixed ratio by mass.
• The relative number and kinds of atoms are constant in a given compound. This reiterates the idea of a fixed composition for any given compound.

While we now know that some of these postulates are not entirely correct (we know atoms can be divided into subatomic particles, and isotopes of an element have different masses), Dalton's theory was a monumental achievement for its time and remains the foundation upon which modern chemistry is built.

What is an Atom? Size and Symbols

An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are incredibly small, far smaller than anything we can imagine or see with the naked eye. Atomic radius is measured in nanometres (nm).

1 nanometre (nm) = 10⁻⁹ metres (m)

To give you a perspective, a single strand of hair is about 100,000 nm thick. An atom of hydrogen has a radius of about 0.037 nm. Millions of atoms stacked on top of each other would barely make a sheet as thick as a piece of paper.

Symbols of Atoms of Different Elements: Dalton was the first scientist to use symbols to represent elements in a very specific sense. However, his symbols were difficult to draw and use. Later, J.J. Berzelius suggested that the symbols of elements be made from one or two letters of the name of the element. The modern system, standardized by the International Union of Pure and Applied Chemistry (IUPAC), follows these rules:

• The symbols for many elements are the first one or two letters of the element's name in English. The first letter is always written as a capital letter (uppercase) and the second letter as a small letter (lowercase). For example, H for Hydrogen, Al for Aluminium, Co for Cobalt (not CO, which is carbon monoxide).
• Symbols for some elements are formed from the first letter of the name and a letter appearing later in the name. For example, Cl for Chlorine, Zn for Zinc.
• Some symbols have been taken from the names of elements in Latin, German, or Greek. For example, the symbol of iron is Fe from its Latin name ferrum, sodium is Na from natrium, and potassium is K from kalium.

What is a Molecule?

A molecule is the smallest particle of an element or a compound that is capable of independent existence and shows all the properties of that substance. It is a group of two or more atoms that are chemically bonded together, i.e., held together by attractive forces. These atoms can be of the same element or of different elements.

Molecules of Elements

The molecules of an element are constituted by the same type of atoms. For many elements, like the noble gases (Helium, Neon, Argon), the molecules are made up of only one atom of that element. But this is not the case for most non-metals. For example, a molecule of oxygen consists of two atoms of oxygen (O₂), and hence it is known as a diatomic molecule. The number of atoms constituting a molecule is known as its atomicity.

Molecules of Compounds

Atoms of different elements join together in definite proportions to form molecules of compounds. For example, a molecule of water (H₂O) contains two atoms of hydrogen and one atom of oxygen. A molecule of ammonia (NH₃) contains one atom of nitrogen and three atoms of hydrogen. The ratio by number of atoms for water is H:O = 2:1, and for ammonia is N:H = 1:3.

What is an Ion?

Atoms are electrically neutral. However, they can gain or lose electrons to become electrically charged particles. These charged particles are known as ions. Compounds composed of metals and non-metals contain charged species (ions).

Cations and Anions

An ion can be positively charged or negatively charged.

• Cation: A positively charged ion is called a cation. It is formed when a neutral atom loses one or more electrons. Metals typically form cations. For example, a sodium atom (Na) loses one electron to become a sodium ion (Na⁺). A magnesium atom (Mg) loses two electrons to become a magnesium ion (Mg²⁺).
• Anion: A negatively charged ion is called an anion. It is formed when a neutral atom gains one or more electrons. Non-metals typically form anions. For example, a chlorine atom (Cl) gains one electron to become a chloride ion (Cl⁻). An oxygen atom (O) gains two electrons to become an oxide ion (O²⁻).

Simple and Polyatomic Ions

Ions can consist of a single charged atom or a group of atoms that have a net charge on them.

• Simple Ions: These are formed from single atoms, like Na⁺, Mg²⁺, Al³⁺, Cl⁻, S²⁻.
• Polyatomic Ions: These are groups of atoms carrying a charge, which behave as a single unit. For example, the ammonium ion (NH₄⁺), the hydroxide ion (OH⁻), the nitrate ion (NO₃⁻), the sulphate ion (SO₄²⁻), and the carbonate ion (CO₃²⁻).

Writing Chemical Formulae

The chemical formula of a compound is a symbolic representation of its composition. It tells us the kinds of atoms present and the number of each kind of atom in one molecule of the compound. To write a chemical formula, we need to know the symbols of the elements and their combining capacity, which is known as valency.

Valency and Ionic Charge

Valency is the number of electrons an atom can lose, gain, or share to form a chemical bond. For ions, the charge on the ion is considered its valency. For example, the valency of Na⁺ is 1, Mg²⁺ is 2, Cl⁻ is 1, and O²⁻ is 2.

Rules for Writing Formulae

1. The valencies or charges on the ion must balance. The overall charge on a compound must be zero.
2. When a compound consists of a metal and a non-metal, the symbol of the metal is written first. For example, in calcium oxide (CaO) and sodium chloride (NaCl), calcium and sodium are written first as they are metals.
3. In compounds formed with polyatomic ions, the ion is enclosed in a bracket before writing the number to indicate the ratio. The bracket is not needed if the number of the polyatomic ion is one.

Examples of Writing Formulae (Criss-Cross Method)

The simplest method is the 'criss-cross' method. We write the symbols of the constituent elements/ions and then write their valencies below them. Then, we cross over the valencies.

Example 1: Formula of Aluminium Oxide

• Step 1: Write symbols: Al O
• Step 2: Write charges/valencies below: Al³⁺ O²⁻ (or simply 3 and 2)
• Step 3: Criss-cross the numerical values of the valencies. The valency of O (2) becomes the subscript for Al, and the valency of Al (3) becomes the subscript for O.
• Resulting Formula: Al₂O₃

Example 2: Formula of Calcium Hydroxide

• Step 1: Write symbols: Ca OH (Hydroxide is a polyatomic ion)
• Step 2: Write charges/valencies below: Ca²⁺ OH⁻ (or simply 2 and 1)
• Step 3: Criss-cross the valencies. The valency of OH (1) goes to Ca, and the valency of Ca (2) goes to OH. Since OH is a polyatomic ion and we need two of them, we must put it in brackets.
• Resulting Formula: Ca(OH)₂

Example 3: Formula of Sodium Carbonate

• Step 1: Write symbols: Na CO₃ (Carbonate is a polyatomic ion)
• Step 2: Write charges/valencies below: Na⁺ CO₃²⁻ (or simply 1 and 2)
• Step 3: Criss-cross the valencies. The valency of CO₃ (2) goes to Na, and the valency of Na (1) goes to CO₃.
• Resulting Formula: Na₂CO₃ (No bracket is needed for CO₃ as its subscript is 1).

Molecular Mass and Mole Concept

This section introduces some of the most important quantitative aspects of chemistry, allowing us to relate mass to the number of particles.

Molecular Mass

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

Example: Calculate the molecular mass of water (H₂O).

• Atomic mass of Hydrogen (H) = 1 u
• Atomic mass of Oxygen (O) = 16 u
• Molecular mass of H₂O = (2 × Atomic mass of H) + (1 × Atomic mass of O)
• = (2 × 1 u) + (1 × 16 u) = 2 u + 16 u = 18 u

Example: Calculate the molecular mass of nitric acid (HNO₃).

• Atomic mass of H = 1 u
• Atomic mass of Nitrogen (N) = 14 u
• Atomic mass of O = 16 u
• Molecular mass of HNO₃ = (1 × At. mass of H) + (1 × At. mass of N) + (3 × At. mass of O)
• = (1 × 1) + (1 × 14) + (3 × 16) = 1 + 14 + 48 = 63 u

Formula Unit Mass

The term 'formula unit mass' is used for substances whose constituent particles are ions (like ionic compounds). The formula unit mass is calculated in the same way as the molecular mass, by summing the atomic masses of all atoms in a formula unit of a compound.

Example: Calculate the formula unit mass of sodium chloride (NaCl).

• Atomic mass of Sodium (Na) = 23 u
• Atomic mass of Chlorine (Cl) = 35.5 u
• Formula unit mass of NaCl = (1 × 23 u) + (1 × 35.5 u) = 58.5 u

The Mole Concept: A Chemist's Counting Unit

Atoms and molecules are \textremely small, so dealing with them individually is impossible. Chemists need a convenient unit to count them in large numbers. Just as we use 'a dozen' for 12 items or 'a gross' for 144 items, chemists use the 'mole' to count atoms, molecules, or ions.

One mole of any species (atoms, molecules, ions, or particles) is that quantity in number having a mass equal to its atomic or molecular mass in grams.

The number of particles present in one mole of any substance is fixed, with a value of 6.022 × 10²³. This is called the Avogadro Constant or Avogadro's Number (represented by Nₐ), named in honour of the Italian scientist Amedeo Avogadro.

• 1 mole of atoms = 6.022 × 10²³ atoms
• 1 mole of molecules = 6.022 × 10²³ molecules
• 1 dozen bananas = 12 bananas
• 1 mole of bananas = 6.022 × 10²³ bananas

The mass of 1 mole of a substance is called its molar mass. The molar mass of a substance is numerically equal to its atomic/molecular/formula unit mass, but the unit is changed from 'u' to 'g' (grams).

• Atomic mass of Oxygen (O) = 16 u
• Molar mass of Oxygen atoms (O) = 16 g/mol
• Molecular mass of Oxygen molecule (O₂) = 32 u
• Molar mass of Oxygen molecules (O₂) = 32 g/mol
• Molecular mass of Water (H₂O) = 18 u
• Molar mass of Water (H₂O) = 18 g/mol

Relationship between Mole, Mass, and Number of Particles

The mole concept provides a bridge between the mass of a substance and the number of particles it contains. The following relationships are crucial for solving numerical problems:

1. Number of moles (n) = Given mass (m) / Molar mass (M)
2. Number of moles (n) = Given number of particles (N) / Avogadro's number (Nₐ)
3. Given mass (m) = Number of moles (n) × Molar mass (M)
4. Given number of particles (N) = Number of moles (n) × Avogadro's number (Nₐ)

By combining these, we can also relate mass directly to the number of particles:

(Given mass / Molar mass) = (Number of particles / Avogadro's number)

Important Questions and Answers

Question 1: 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 acetate. Show that these observations are in agreement with the law of conservation of mass.

Answer:

To verify the law of conservation of mass, we need to show that the total mass of reactants is equal to the total mass of products.

Reactants are: Sodium carbonate and Acetic acid

Total mass of reactants = Mass of sodium carbonate + Mass of acetic acid

= 5.3 g + 6.0 g = 11.3 g

Products are: Carbon dioxide, Water, and Sodium acetate

Total mass of products = Mass of carbon dioxide + Mass of water + Mass of sodium acetate

= 2.2 g + 0.9 g + 8.2 g = 11.3 g

Since the total mass of reactants (11.3 g) is equal to the total mass of products (11.3 g), these observations are in agreement with the law of conservation of mass.

Question 2: Calculate the number of moles for the following: (i) 52 g of He (finding mole from mass) (ii) 12.044 × 10²³ number of He atoms (finding mole from number of particles). (Atomic mass of He = 4 u)

Answer:

(i) To find moles from mass:

We use the formula: Number of moles (n) = Given mass (m) / Molar mass (M)

Given mass of Helium (m) = 52 g

Atomic mass of Helium = 4 u

Therefore, Molar mass of Helium (M) = 4 g/mol

n = 52 g / 4 g/mol = 13 moles

So, 52 g of Helium is equal to 13 moles.

(ii) To find moles from the number of particles:

We use the formula: Number of moles (n) = Given number of particles (N) / Avogadro's number (Nₐ)

Given number of Helium atoms (N) = 12.044 × 10²³

Avogadro's number (Nₐ) = 6.022 × 10²³

n = (12.044 × 10²³) / (6.022 × 10²³)

n = 2 moles

So, 12.044 × 10²³ atoms of Helium is equal to 2 moles.

Question 3: Calculate the mass of the following: (i) 0.5 mole of N₂ gas (ii) 3.011 × 10²³ number of N atoms

Answer:

(i) Mass from moles of N₂ gas (molecule):

We use the formula: Mass (m) = Number of moles (n) × Molar mass (M)

Number of moles (n) = 0.5 mole

Molar mass of N₂ molecule (M) = 2 × Atomic mass of N = 2 × 14 g/mol = 28 g/mol

m = 0.5 mol × 28 g/mol = 14 g

So, the mass of 0.5 mole of N₂ gas is 14 g.

(ii) Mass from the number of N atoms:

First, we need to find the number of moles (n).

n = Given number of particles (N) / Avogadro's number (Nₐ)

n = (3.011 × 10²³) / (6.022 × 10²³) = 0.5 moles

Now, we find the mass from these moles.

Mass (m) = Number of moles (n) × Molar mass of N atoms (M)

Molar mass of N atoms = 14 g/mol

m = 0.5 mol × 14 g/mol = 7 g

So, the mass of 3.011 × 10²³ N atoms is 7 g.

Question 4: Write the chemical formula of (i) Magnesium chloride (ii) Aluminium sulphate (iii) Sodium sulphide

Answer:

(i) Magnesium chloride

Symbols: Mg Cl

Charges: 2+ 1-

By criss-crossing the valencies, we get the formula: MgCl₂

(ii) Aluminium sulphate

Symbols: Al SO₄

Charges: 3+ 2-

By criss-crossing the valencies (and using brackets for the polyatomic ion), we get the formula: Al₂(SO₄)₃

(iii) Sodium sulphide

Symbols: Na S

Charges: 1+ 2-

By criss-crossing the valencies, we get the formula: Na₂S

Chapter Summary

Here are the key takeaways from the chapter "Atoms and Molecules":

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
• Law of Constant Proportions: In a pure chemical compound, elements are always present in a definite proportion by mass.
• Atom: The smallest particle of an element that retains the properties of that element and can participate in a chemical reaction.
• Dalton's Atomic Theory: Provided the first scientific basis for the concept of atoms as fundamental, indivisible particles of matter.
• Molecule: The smallest particle of an element or compound that can exist independently and shows all the properties of that substance. It is a cluster of atoms held together by chemical bonds.
• Ion: An electrically charged atom or group of atoms, formed by the loss (cation, positive) or gain (anion, negative) of electrons.
• Chemical Formula: A shorthand representation of a compound using symbols and subscripts to indicate the number of atoms of each element present.
• Valency: The combining capacity of an element.
• Molecular Mass: The sum of the atomic masses of all atoms in a molecule, expressed in atomic mass units (u).
• Mole: A unit representing a specific number of particles (6.022 × 10²³), known as Avogadro's number.
• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol). It is numerically equal to the molecular mass.