Atoms And Molecules

The concept of atoms and molecules developed gradually through the ideas of philosophers and scientists.

Indian philosopher Maharishi Kanad proposed that if matter is divided continuously, a stage would come when further division would not be possible. He called these smallest particles Parmanu.

Another Indian philosopher Maharishi Pakudha Katyayama suggested that these particles generally exist in combined forms which produce different forms of matter.

Greek philosophers Leucippus, Democritus and Epicurus also discussed indivisible particles of matter. They called these particles atoms, meaning “indivisible”.

By the end of the 18th century, scientists clearly distinguished between elements and compounds and established important laws of chemical combination.

Laws Of Chemical Combination

The laws of chemical combination explain how elements combine to form compounds. These laws were established mainly by Antoine Lavoisier and Joseph L. Proust.

Law Of Conservation Of Mass

The Law of Conservation of Mass states that:

• Mass can neither be created nor destroyed during a chemical reaction.

In other words:

Total mass of reactants = Total mass of products

Example:

Carbon + Oxygen → Carbon Monoxide

12 g + 16 g = 28 g

The total mass remains unchanged before and after the reaction.

Law Of Constant Proportion

The Law of Constant Proportion, also called the Law of Definite Proportion, states that:

• In a chemical compound, the elements are always present in a fixed proportion by mass.

Example:

In water, hydrogen and oxygen are always present in the mass ratio:

1 : 8

This ratio remains the same irrespective of the source of water.

Dalton’s Atomic Theory

Since the laws of chemical combination needed an explanation, John Dalton proposed the Atomic Theory.

Postulates Of Dalton’s Atomic Theory

• Matter is made up of extremely small particles called atoms.
• Atoms cannot be further divided during chemical reactions.
• All atoms of a given element are identical in mass, size and chemical properties.
• Atoms of different elements have different masses and properties.
• Atoms combine in simple whole-number ratios to form compounds.
• In a chemical reaction, atoms are only rearranged.
• Atoms are neither created nor destroyed during chemical reactions.

Drawbacks Of Dalton’s Atomic Theory

Modern discoveries revealed some limitations of Dalton’s theory.

Presence Of Subatomic Particles

Atoms are divisible into:

• Electron
• Proton
• Neutron

Therefore, atoms are not indivisible.

Existence Of Isotopes

Atoms of the same element may have different masses.

Example:

• Chlorine-35
• Chlorine-37

Such atoms are called isotopes.

Isotopes

Isotopes are atoms having:

• Same atomic number
• Different mass numbers

Existence Of Isobars

Atoms of different elements may have the same mass number.

Example:

• Argon-40
• Calcium-40

These are called isobars.

Isobars

Isobars are atoms having:

• Same mass number
• Different atomic numbers

Symbols Of Elements

A symbol is the shorthand representation of an element.

Dalton was the first scientist to use symbols systematically for elements.

A symbol also represents:

• One atom of that element

Atom

An atom is the smallest particle of an element that participates in chemical reactions.

Atoms are extremely small.

Modern instruments such as the Scanning Tunneling Microscope (STM) have made it possible to study atoms.

Atomic Mass

Since atoms are very small, their masses are also extremely small.

Therefore, atomic masses are expressed using atomic mass unit (u).

Atomic Mass Unit

One atomic mass unit is defined as:

1 u = 1/12th of the mass of one Carbon-12 atom

The Carbon-12 scale was adopted internationally in 1961.

Important Values

• Mass of one Carbon-12 atom = 1.9926 × 10⁻²³ g
• 1 u = 1.6605 × 10⁻²⁴ g

Average Atomic Mass

Most elements occur naturally as mixtures of isotopes.

Therefore, atomic mass is taken as the weighted average of isotopic masses.

This is called average atomic mass.

Molecule

A molecule is the smallest particle of a substance capable of independent existence and showing all the properties of that substance.

Molecules Of Elements

Molecules may contain atoms of the same element.

Examples:

• H₂
• O₂
• N₂
• Cl₂
• O₃

Atomicity

The number of atoms present in one molecule of an element is called atomicity.

Examples:

Molecule	Atomicity
H₂	2
O₂	2
O₃	3
N₂	2

Molecules Of Compounds

Atoms of different elements combine in fixed proportions to form molecules of compounds.

Examples:

• H₂O
• NH₃
• CO₂
• SO₂
• CH₄

Ion

An ion is a charged particle carrying either positive or negative charge.

Example:

NaCl consists of:

• Na⁺ (cation)
• Cl⁻ (anion)

Types Of Ions

Cation

Positively charged ion.

Examples:

• Na⁺
• Ca²⁺
• Ba²⁺

Anion

Negatively charged ion.

Examples:

• Cl⁻
• S²⁻
• H⁻

Polyatomic Ion

A group of atoms carrying a charge is called a polyatomic ion.

Example:

• NH₄⁺

Chemical Formulae

The formula of an ionic compound is obtained from the charges of positive and negative ions.

Valency

The combining capacity of an element is called its valency.

Examples:

Element	Valency
H	1
O	2
Na	1
Mg	2
Al	3
Cl	1

Molecular Mass

Molecular mass is the relative mass of a molecule compared to Carbon-12.

It is calculated by adding the atomic masses of all atoms present in the molecule.

Examples:

Compound	Molecular Mass
HNO₃	63 u
C₁₂H₂₂O₁₁	342 u
O₂	32 u

Gram Atomic Mass

Gram atomic mass is the quantity of an element whose mass in grams is numerically equal to its atomic mass.

Example:

• Atomic mass of hydrogen = 1 u
• Gram atomic mass of hydrogen = 1 g

It is also called gram atom.

Gram Molecular Mass

Gram molecular mass is the quantity of a substance whose mass in grams is numerically equal to its molecular mass.

Example:

• Molecular mass of O₂ = 32 u
• Gram molecular mass of O₂ = 32 g

It is also called gram molecule.

Mole Concept

The mole concept is one of the most convenient ways of expressing the amount of reactants and products in a chemical reaction.

A mole represents a fixed number of particles.

Avogadro Number

The number of particles present in one mole of any substance is:

6.022 × 10²³

This number is called:

• Avogadro Constant
• Avogadro Number

It is represented by Nₐ.

It is named after Italian scientist Amedeo Avogadro.

One Mole Of Atoms

1 mole of atoms of an element contains:

6.022 × 10²³ atoms

Examples:

Element	Atomic Mass	1 Mole Contains
Hydrogen	1 u	6.022 × 10²³ atoms
Carbon	12 u	6.022 × 10²³ atoms
Oxygen	16 u	6.022 × 10²³ atoms

One Mole Of Molecules

1 mole of molecules contains:

6.022 × 10²³ molecules

Important Conceptual Points

• Matter is made up of atoms.
• Molecules may consist of same or different kinds of atoms.
• Isotopes have same atomic number but different mass numbers.
• Isobars have same mass number but different atomic numbers.
• Carbon-12 is the standard for atomic mass scale.
• 1 u = 1/12th mass of Carbon-12 atom.
• Atomicity means number of atoms in a molecule.
• Valency is the combining capacity of an element.
• One mole always contains 6.022 × 10²³ particles.
• Gram atomic mass and gram molecular mass are numerically equal to atomic and molecular masses respectively.

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