Class 9 Atoms and Molecules NCERT Full Chapter

Atoms and Molecules

Imagine you have a huge boulder in front of you, and you keep breaking it into smaller and smaller pieces. At some point, you’ll reach a stage where the particle becomes so small that it can no longer be divided further without losing its properties. This smallest indivisible particle is what we call an atom. Today, we’ll discuss the concept of atoms, molecules, laws governing their combinations, and how these tiny particles form the basis of all matter around us.

Laws of Chemical Combination

Before we dive into the concepts of atoms and molecules, let’s first understand the rules that govern how substances combine chemically. These rules are called the Laws of Chemical Combination.

1. Law of Conservation of Mass

This law states that mass can neither be created nor destroyed during a chemical reaction.
Explanation:
For example, if hydrogen reacts with oxygen to form water, the total mass of hydrogen and oxygen before the reaction will be equal to the mass of water after the reaction.

Mathematically:

Mass of reactants = Mass of Products

Law of Definite Proportions

This law states that, ” a chemical compound always contains the same elements in the same proportion by mass, irrespective of the source or method of preparation.”
Example:
Water (H₂O) will always contain 2 parts hydrogen and 16 parts oxygen by mass, no matter where you get the water from.

What Are Atoms?

Now, let’s dive deeper into atoms. Atoms are the basic building blocks of matter. Everything around us, from the air we breathe to the food we eat, is made up of atoms. But remember, atoms are indivisible in chemical processes—meaning you can’t break them down further by chemical means.

Key Features of Atoms:

The atom is the fundamental unit of matter, retaining the characteristics of an element. Below we will know about its key features:

1. Smallest Unit of Matter:
An atom represents the tiniest division of matter that still maintains the properties of a specific element.


2. Subatomic Particles:

Protons: Positively charged particles found in the nucleus.

Neutrons: Neutral particles also located in the nucleus.

Electrons: Negatively charged particles that orbit the nucleus in specific energy levels.



3. Nucleus:

The dense center of an atom containing protons and neutrons.

It accounts for most of the atom’s mass.



4. Electron Arrangement:
Electrons are arranged in defined shells or energy levels around the nucleus, labeled as K, L, M, etc.


5. Neutrality:
In a neutral atom, the number of protons equals the number of electrons, resulting in no overall charge.


6. Atomic Number (Z):
This is the total number of protons in an atom’s nucleus, uniquely identifying each element.


7. Mass Number (A):
The mass number is the sum of protons and neutrons, representing the atom’s approximate mass.


8. Isotopes:
Atoms of the same element with identical atomic numbers but varying mass numbers due to differences in the number of neutrons.


9. Chemical Behavior:
The chemical properties of an atom are influenced by the electrons in its outermost shell, also known as valence electrons.


10. Quantum Mechanics:
The behavior and position of electrons follow quantum principles, described as regions of probability called orbitals.


11. Conservation in Reactions:
During chemical reactions, atoms are neither created nor destroyed but are simply rearranged to form new substances.

We can summarise these above like as follows below :-

1. Atoms are extremely small and cannot be seen with the naked eye.


2. They combine to form molecules or compounds.


3. Each element has its unique type of atom.

Symbols of Atoms

Each atom is represented by a specific symbol. For example:

Hydrogen: H

Oxygen: O

Sodium: Na

Chlorine: Cl

Atoms & their Symbols

Here is a list of elements with their atomic number and symbols from 1 to 118:

1. H – Hydrogen


2. He – Helium


3. Li – Lithium


4. Be – Beryllium


5. B – Boron


6. C – Carbon


7. N – Nitrogen


8. O – Oxygen


9. F – Fluorine


10. Ne – Neon


11. Na – Sodium


12. Mg – Magnesium


13. Al – Aluminium


14. Si – Silicon


15. P – Phosphorus


16. S – Sulfur


17. Cl – Chlorine


18. Ar – Argon


19. K – Potassium


20. Ca – Calcium


21. Sc – Scandium


22. Ti – Titanium


23. V – Vanadium


24. Cr – Chromium


25. Mn – Manganese


26. Fe – Iron


27. Co – Cobalt


28. Ni – Nickel


29. Cu – Copper


30. Zn – Zinc


31. Ga – Gallium


32. Ge – Germanium


33. As – Arsenic


34. Se – Selenium


35. Br – Bromine


36. Kr – Krypton


37. Rb – Rubidium


38. Sr – Strontium


39. Y – Yttrium


40. Zr – Zirconium


41. Nb – Niobium


42. Mo – Molybdenum


43. Tc – Technetium


44. Ru – Ruthenium


45. Rh – Rhodium


46. Pd – Palladium


47. Ag – Silver


48. Cd – Cadmium


49. In – Indium


50. Sn – Tin


51. Sb – Antimony


52. I – Iodine


53. Xe – Xenon


54. Cs – Cesium


55. Ba – Barium


56. La – Lanthanum


57. Ce – Cerium


58. Pr – Praseodymium


59. Nd – Neodymium


60. Pm – Promethium


61. Sm – Samarium


62. Eu – Europium


63. Gd – Gadolinium


64. Tb – Terbium


65. Dy – Dysprosium


66. Ho – Holmium


67. Er – Erbium


68. Tm – Thulium


69. Yb – Ytterbium


70. Lu – Lutetium


71. Hf – Hafnium


72. Ta – Tantalum


73. W – Tungsten


74. Re – Rhenium


75. Os – Osmium


76. Ir – Iridium


77. Pt – Platinum


78. Au – Gold


79. Hg – Mercury


80. Tl – Thallium


81. Pb – Lead


82. Bi – Bismuth


83. Po – Polonium


84. At – Astatine


85. Rn – Radon


86. Fr – Francium


87. Ra – Radium


88. Ac – Actinium


89. Th – Thorium


90. Pa – Protactinium


91. U – Uranium


92. Np – Neptunium


93. Pu – Plutonium


94. Am – Americium


95. Cm – Curium


96. Bk – Berkelium


97. Cf – Californium


98. Es – Einsteinium


99. Fm – Fermium


100. Md – Mendelevium


101. No – Nobelium


102. Lr – Lawrencium


103. Rf – Rutherfordium


104. Db – Dubnium


105. Sg – Seaborgium


106. Bh – Bohrium


107. Hs – Hassium

108. Mt – Meitnerium


109. Ds – Darmstadtium


110. Rg – Roentgenium


111. Cn – Copernicium


112. Nh – Nihonium


113. Fl – Flerovium


114. Mc – Moscovium


115. Lv – Livermorium


116. Ts – Tennessine


117. Og – Oganesson


118. Uuo – Ununoctium (now known as Oganesson, Og)



These elements make up the current periodic table from Hydrogen (1) to Oganesson (118).

The symbols are derived from either the English names or the Latin names of the elements (e.g., Sodium is , from its Latin name Natrium).

Atomic Mass

Let’s discuss atomic mass, which is one of the fundamental properties of an atom. It tells us how heavy an atom is compared to the mass of a hydrogen atom.

Atomic mass is expressed in atomic mass units (amu).

1 amu is defined as {1}/{12}th of the mass of a carbon-12 atom.
For example:

Atomic mass of oxygen = 16 amu

Atomic mass of hydrogen = 1 amu

What Are Molecules?

When two or more than two  similar  type of  atoms combine together so it forms a molecule

Atoms rarely exist independently. Instead, they combine to form molecules, which are groups of two or more atoms chemically bonded together. Molecules can be of two types:

1. Molecules of Elements:

Molecules consisting of the same type of atoms.
Examples:

O₂ (oxygen gas)

H2 (hydrogen gas)



2. Molecules of Compounds:
Molecules consisting of different types of atoms.
Examples:

H₂O (water)

CO₂ (carbon dioxide)

Chemical Formulae

To represent molecules, we use chemical formulae. A chemical formula is a symbolic representation of a molecule, showing the types and numbers of atoms involved.

Rules for Writing a Chemical Formula:

1. The valency of each element is considered. Valency is the combining capacity of an element.


2. The symbols of elements are arranged in a specific sequence (usually metals before non-metals).
For example:

Water: H₂O

Ammonia:NH3

Mole Concept

One of the most powerful tools in chemistry is the mole concept. It helps us count atoms and molecules, which are incredibly small in size.

What is a Mole?

A mole is a quantity that contains  particles (atoms, molecules, or ions). This number is called Avogadro’s number.


Molar Mass:

The mass of 1 mole of a substance is its molar mass.
For example:

Molar mass of oxygen (O₂) = 32 g/mol

Molar mass of water (H₂O) = 18 g/mol

What is Empirical Formula

The empirical formula of a compound shows the simplest ratio of atoms of each element in that compound. It does not give the exact number of atoms but tells us the basic proportion of elements.

For example:

If a compound has 2 hydrogen atoms for every 1 oxygen atom, its empirical formula is H₂O.

Key Points:

1. Simplest Ratio: It only tells the simplest whole number ratio of elements, not the actual number of atoms.
2. Different from Molecular Formula: The molecular formula shows the actual number of atoms, while the empirical formula simplifies it.

Examples:

1. Glucose

Molecular formula: C₆H₁₂O₆

Empirical formula: CH₂O (because 6:12:6 simplifies to 1:2:1).

2. Hydrogen Peroxide

Molecular formula: H₂O₂

Empirical formula: HO (because 2:2 simplifies to 1:1).

3. Ethene (C₂H₄)

Molecular formula: C₂H₄

Empirical formula: CH₂ (because 2:4 simplifies to 1:2).

4. Water

Molecular formula: H₂O

Empirical formula: H₂O (already in simplest ratio).

How to Find the Empirical Formula:

1. Find the ratio of elements in the compound.
2. Simplify the ratio to the smallest whole numbers.

Example:
If a compound contains 4 grams of hydrogen and 32 grams of oxygen:

Find the ratio using their atomic masses:

Hydrogen (H) = 1 g/mol, Oxygen (O) = 16 g/mol.

Moles of H = 4/1 = 4, Moles of O = 32/16 = 2.

Ratio of H:O = 4:2 = 2:1.

Empirical formula = H₂O.

Here are some practice problems related to empirical formula, along with solutions to help students understand:

Problem 1:

A compound contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen by mass. Find its empirical formula.
(Atomic masses: C = 12, H = 1, O = 16)

Solution:

1. Convert percentages to moles:

Moles of C = 40/12 = 3.33

Moles of H = 6.7/1 = 6.7

Moles of O = 53.3/16 = 3.33

2. Divide by the smallest value (3.33):

C: 3.33/3.33 = 1

H: 6.7/3.33 = 2

O: 3.33/3.33 = 1

3. Empirical formula: CH₂O

Problem 2:

A compound is found to contain 70% iron (Fe) and 30% oxygen (O) by mass. Find its empirical formula.
(Atomic masses: Fe = 56, O = 16)

Solution:

1. Convert percentages to moles:

Moles of Fe = 70/56 = 1.25

Moles of O = 30/16 = 1.875

2. Divide by the smallest value (1.25):

Fe: 1.25/1.25 = 1

O: 1.875/1.25 = 1.5

3. Multiply to get whole numbers:

Fe: 1×2 = 2

O: 1.5×2 = 3

4. Empirical formula: Fe₂O₃

Problem 3:

A compound contains 2.7 g of aluminum (Al) and 2.4 g of oxygen (O). Find its empirical formula.
(Atomic masses: Al = 27, O = 16)

Solution:

1. Convert mass to moles:

Moles of Al = 2.7/27 = 0.1

Moles of O = 2.4/16 = 0.15

2. Divide by the smallest value (0.1):

Al: 0.1/0.1 = 1

O: 0.15/0.1 = 1.5

3. Multiply to get whole numbers:

Al: 1×2 = 2

O: 1.5×2 = 3

4. Empirical formula: Al₂O₃

Problem 4:

A compound is 87.5% nitrogen (N) and 12.5% hydrogen (H) by mass. Find its empirical formula.
(Atomic masses: N = 14, H = 1)

Solution:

1. Convert percentages to moles:

Moles of N = 87.5/14 = 6.25

Moles of H = 12.5/1 = 12.5

2. Divide by the smallest value (6.25):

N: 6.25/6.25 = 1

H: 12.5/6.25 = 2

3. Empirical formula: NH₂

What is AMU?

AMU stands for Atomic Mass Unit. It is a small unit of mass used to measure the mass of atoms and molecules.

1. Definition:
   1 AMU is defined as one-twelfth (1/12) the mass of a carbon-12 atom.

In simple terms, it’s a way to express how heavy an atom is compared to a carbon atom.

2. Why is AMU Used?
   Atoms are extremely small, and their masses are too tiny to measure in grams. So, scientists use AMU to make it easier to compare atomic masses.

Key Points:

1. Mass of Proton: ~1 AMU
2. Mass of Neutron: ~1 AMU
3. Mass of Electron: ~0.00055 AMU (negligible compared to protons and neutrons).

Examples of Atomic Mass in AMU:

Hydrogen (H): 1 AMU

Carbon (C): 12 AMU

Oxygen (O): 16 AMU

When we say “Oxygen has an atomic mass of 16 AMU,” it means it is 16 times heavier than 1/12 of a carbon-12 atom.

Fun Fact:

AMU is now officially called the Dalton (Da) in modern science, but the term AMU is still widely used in classrooms.

Example Problem:

How many molecules are present in 18 g of water?
Solution:
1 mole of water =  molecules
Mass of 1 mole of water = 18 g
So, 18 g of water contains  molecules.

Dalton’s Atomic Theory

Finally, let’s revisit Dalton’s Atomic Theory, which laid the foundation for our understanding of atoms:

1. Matter is made up of tiny indivisible particles called atoms.


2. Atoms of a given element are identical in mass and properties.


3. Atoms combine in fixed, whole-number ratios to form compounds.


4. Atoms are neither created nor destroyed in chemical reactions.



While Dalton’s theory had limitations (e.g., atoms are divisible into subatomic particles), it remains a cornerstone of chemistry.



Short Notes

Atoms are the smallest units of matter, and molecules are groups of atoms bonded together.

Chemical reactions follow the Laws of Conservation of Mass and Definite Proportions.

The mole concept simplifies the counting of atoms and molecules in bulk quantities.

Understanding chemical formulae and atomic mass is essential for representing and calculating chemical substances.



I hope you now have a clear understanding of how atoms and molecules form the basis of everything in the universe. If you have any questions, feel free to ask!