Class 9 ATOMS AND MOLECULES

📝 ATOMS AND MOLECULES

Introduction

👉Ancient Indian and Greek philosophers have always wondered about the unknown and unseen form of matter. The idea of divisibility of matter was considered long back in India, around 500 BC.

👉 Maharishi Kanad, said that if we go on dividing the matter, we will get smaller and smaller particles of matter. Ultimately we will get the smallest particle of matter, which cannot be divided any further. Kanad was one of the first persons to propose that matter is made up of very small particles called 'parmanu'.

👉 Another Indian philosopher, Pakudha Katyayama, elaborated 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 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).

All this was based on philosophical considerations and not much experimental work to validate these ideas could be done till the eighteenth century.

👉By the end of the eighteenth century, scientists recognised the difference between elements and compounds and naturally became interested in finding out how and why elements combine and what happens when they combine.

Laws of Chemical Combination

Law of conservation of mass
Law constant proportions

1. Law of conservation of mass

The Law of conservation of mass was given by Antoine Lavoisier in 1774. The law of conservation of mass means that in a chemical reaction, the total mass of the products is equal to the total mass of the reactants.

Example 1: It has been found by experiments that if 100 grams of calcium carbonate are decomposed completely then 56 grams of calcium oxide and 44 grams of carbon dioxide are formed.

$CaCO_3 overset{triangle} rightarrow CaO + CO_2$

Calcium carbonate (100g)

$rightarrow$ Calcium oxide (56 g) + Carbon dioxide (44g)

Since the total mass of the products (56 g + 44 g = 100g) is equal to the total mass of the reactant (100g), there is no change of mass during this chemical reaction. The mass remains the same or conserved.

Example 2

$BaCl_2 + Na_2SO_4 rightarrow BaSO_4 + NaCl$

Barium chloride = 20.8 g

Sodium sulphate = 14.2 g

Barium sulphate = 23.3 g

Sodium chloride = 11.7 g

Mass of the reactants = 20.8 g + 14.2 g = 35.0 g

Mass of the products = 23.3 g + 11.7 g = 35.0 g

Since the total mass of the products (35 g) is equal to the total mass of the reactants (35 g), therefore the given data verify the law of conservation of mass.

👉Law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction.

2. Law constant proportions

The law of constant proportions was given by Joseph Proust in 1779. A chemical compound always consists of the same elements combined together in the same proportion by mass. For example, water is a compound that always consists of the same elements hydrogen and oxygen combined together in the same proportion of 1: 8. Ammonia is a compound. It has been found by analysis that ammonia always consists of the same two elements, nitrogen and hydrogen, combined together in the same ratio of 14:3 by mass.

John Dalton provided the basic theory about the nature of matter. Dalton picked up the idea of divisibility of matter, which was till then just a philosophy. He took the name ‘atoms’ as given by the Greeks and said that the smallest particles of matter are atoms. His theory was based on the laws of chemical combination. Dalton’s atomic theory provided an explanation for the law of conservation of mass and the law of definite proportions.

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 of this theory may be stated as follows:

(i) All matter is made of very tiny particles called atoms, which participate in chemical reactions.

(ii) Atoms are indivisible particles, which 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 chemical properties.

(v) Atoms combine in the ratio of small whole numbers to form compounds.

(vi) The relative number and kinds of atoms are constant in a given compound.

Drawbacks of Dalton's Atomic Theory

(i) One of the major drawbacks is that atoms were thought to be indivisible. But atoms can be further divided into still smaller particles called electrons, protons, and neutrons.

(ii) Dalton's atomic theory says that all atoms of an element have exactly the same mass. It is however we now know that atoms of the same element can have slightly different masses.

(iii) Dalton's atomic theory said that atoms of different elements have different masses. But now known that even atoms of different elements can have the same mass.

What is an Atom?

An atom is the basic unit of matter, made up of three main particles: protons (positively charged), neutrons (neutral), and electrons (negatively charged). The protons and neutrons form the atom's nucleus at the centre, while electrons orbit around the nucleus. Atoms combine to form molecules, which make up all substances.

How big are atoms?

Atoms are very small. More than millions of atoms when stacked would make a layer barely as thick as this sheet of paper.

THE MODERN DAY SYMBOLS OF ATOMS OF DIFFERENT ELEMENTS

👉Dalton was the first scientist to use the symbols for elements in a very specific sense. When he used a symbol for an element he also meant a definite quantity of that element, that is, one atom of that element.

👉In the beginning, the names of elements were derived from the name of the place where they were found for the first time. For example, the name copper was taken from Cyprus. Some names were taken from specific colours. For example, gold was taken from the English word meaning yellow.

👉Now-a-days, IUPAC (International Union of Pure and Applied Chemistry) is an international scientific organisation which approves names of elements, symbols and units. Many of the symbols are the first one or two letters of the element’s name in English. The first letter of a symbol is always written as a capital letter (uppercase) and the second letter as a small letter (lowercase).

For example

(i) hydrogen, H

(ii) aluminium, Al and not AL

(iii) cobalt, Co and not CO.

👉Symbols of some elements are formed from the first letter of the name and a letter, appearing later in the name. Examples are: (i) chlorine, Cl, (ii) zinc, Zn etc.

👉Other 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, potassium is K from kalium. Therefore, each element has a name and a unique chemical symbol.

Atoms

➤ An atom is the smallest particle of an element that can take part in a chemical reaction.

Atoms of most of the elements are very reactive and do not exist in the free state. They exist in combination with the atoms of the same elements or another element.

Atomic Size

➤ The size of an atom is indicated by its radius, which is called 'atomic radius'. Atomic radius is measured in 'nanometres' (very very small unit of measuring length). The symbol of nanometre is nm.

$1nm = frac{1}{10^9} m$

$1nm = 10^-9 m$

The hydrogen atom is the smallest atom of all. The atomic radius of a hydrogen atom is 0.037 nm. If we express the radius of a hydrogen atom in metre, it will be

$0.037 × 10^-9 m$ , which means 0. 000000000037 m! It is really very very small.

HOW DO ATOMS EXIST?

◼ Atoms of most elements are not able to exist independently.

Atoms form molecules and ions. These molecules or ions aggregate in large numbers to form the matter that we can see, feel or touch.

Molecule

➤ A molecule is in general a group of two or more atoms that are chemically bonded together, that is, tightly held together by attractive forces.

or

➤ A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance.

MOLECULES OF ELEMENTS

◼ The molecules of an element are constituted by the same type of atoms.

◼ Molecules of many elements, such as argon (Ar), helium (He) etc. are made up of only one atom of that element.

◼ But this is not the case with most of the non-metals. For example, a molecule of oxygen consists of two atoms of oxygen and hence it is known as a diatomic molecule,

$O_2$ . If 3 atoms of oxygen unite into a molecule, instead of the usual 2, we get ozone,

$O_3$ .

Atomicity

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

Let us look at the atomicity of some non-metals.

MOLECULES OF COMPOUNDS

Atoms of different elements join together in definite proportions to form molecules of compounds. Few examples are given in Table 3.4.

Ions

➤ The charged species are known as ions.

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

Types of ions

Cations
Anions

1. A positively charged ion is called a cation.

2. A negatively charged ion is called an anion.

Take, for example, sodium chloride

$(NaCl)$ . Its constituent particles are positively charged sodium ions

$(Na^+)$ and negatively charged chloride ions $(Cl^–)$ .

Polyatomic ion

➤ A group of atoms carrying a charge is known as a polyatomic ion. Examples:

$OH^-, NO_3^-$ , etc.

Names and symbols of some ions

Chemical Formulae

➤ The chemical formula of a compound is a symbolic representation of its composition.

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 therefore the relative mass of a molecule expressed in atomic mass units (u).

FORMULA UNIT MASS

➤The formula unit mass of a substance is the sum of the atomic masses of all atoms in a formula unit of a compound.

The word formula unit for those substances whose constituent particles are ions.

Mole concept

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

The number of particles (atoms, molecules, or ions) present in 1 mole of any substance is fixed, with a value of

$6.022 × 10^23$ . This is an experimentally obtained value. This number is called the Avogadro Constant or Avogadro Number (represented by

$N_0$ ), named in honour of the Italian scientist, Amedeo Avogadro.

1 mole = $6.022 × 10^23$

Molar Mass

➤ The mass of one mole of a substance in grams is called its molar mass. The molar mass in grams is numerically equal to atomic/molecular/ formula mass in u.

Molar mass of atoms is also known as gram atomic mass.

For example, the atomic mass of hydrogen=1u. So, the gram atomic mass of hydrogen = 1 g.

Examples:

1u hydrogen has only 1 atom of hydrogen, 1 g hydrogen has 1-mole atoms, that is,

$6.022 × 10^23$ atoms of hydrogen.

Similarly, 16 u oxygen has only 1 atom of oxygen, 16 g oxygen has 1-mole atoms, that is,

$6.022 × 10^23$ atoms of oxygen.

18 u water has only 1 molecule of water, 18 g water has 1-mole molecules of water, that is,

$6.022 × 10^23$ molecules of water.

Calculations

$n =$ No. of moles

$m =$ given mass

$M =$ Molar mass

$n = m/M$

$N =$ Given no of particles

$N_A =$ Avogadro number

$n = N/N_A$

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