🔬 Class 9 Journey Inside the Atom – Short Notes

These Class 9 Journey Inside the Atom notes explain the development of the atomic theory, subatomic particles, different atomic models, electronic configuration, valency, atomic number, mass number, isotopes and isobars in a simple and exam-friendly manner.

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🧪 1. History of Development of the Atom

🔹 Kanad's Idea

• In ancient times, Kanad proposed that if matter is continuously divided, we would eventually obtain very small particles.
• He suggested that these smallest particles could not be divided further.
• He called these particles Parmanu or atom.

🔹 Pakudha Katyayama

• Pakudha Katyayama stated that different forms of matter exist in nature.
• He suggested that atoms normally exist in a combined form, giving rise to different forms of matter.

🔹 Democritus

• The ancient Greek philosopher Democritus suggested that matter could not be divided indefinitely.
• He proposed that a stage would come when particles could not be divided further.
• He called these indivisible particles atoms.

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⚖️ 2. Laws of Chemical Combination

🔹 Law of Conservation of Mass

The Law of Conservation of Mass states that mass can neither be created nor destroyed during a chemical reaction.

Formula:

Mass of reactants = Mass of products

Example:

Calcium carbonate on heating produces calcium oxide and carbon dioxide.

100 g Calcium Carbonate → 56 g Calcium Oxide + 44 g Carbon Dioxide

🔹 Law of Constant Proportion

A chemical compound always consists of the same elements combined together in a fixed proportion by mass.

Example: Water (H₂O)

• Hydrogen and oxygen are always present in water in the same ratio by mass.
• The mass ratio of hydrogen : oxygen is 1 : 8.
• Therefore, 1 g of hydrogen combines with 8 g of oxygen to form water.

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⚛️ 3. Dalton's Atomic Theory

John Dalton proposed his atomic theory in 1808.

📌 Main Postulates

1. All matter is made up of very tiny particles called atoms.
2. Atoms participate in chemical reactions.
3. Atoms cannot be created or destroyed in a chemical reaction.
4. Atoms of the same element have identical physical and chemical properties.
5. Atoms of different elements have different masses and chemical properties.
6. Atoms combine in simple whole-number ratios to form compounds.
7. The relative number and kinds of atoms in a given compound remain constant.

Examples: H₂O, CO₂ and NH₃ contain atoms combined in definite ratios.

❌ Drawbacks of Dalton's Atomic Theory

• Dalton considered atoms to be indivisible, but atoms are made up of electrons, protons and neutrons.
• Atoms of the same element may have different masses. These are called isotopes.
• Atoms of different elements may have the same mass number. Such atoms are called isobars.

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⚡ 4. Origin of Charge in Matter

Several phenomena, such as the attraction between a rubbed plastic scale and paper, could not be explained by assuming that matter was made only of indivisible particles.

🔹 Electrostatic Force

The force of attraction or repulsion between electric charges is called electrostatic force.

• There are two types of charges: positive and negative.
• Like charges repel each other.
• Unlike charges attract each other.

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🔬 5. Cathode Ray Experiment

The cathode ray experiment was performed by J. J. Thomson. It led to the discovery of the negatively charged particle called the electron.

📌 Important Facts

• Charge of an electron = −1.6 × 10⁻¹⁹ C
• Mass of an electron ≈ 9.1 × 10⁻³¹ kg
• Cathode rays consist of negatively charged particles called electrons.

In the experiment, a gas discharge tube containing electrodes was connected to a high-potential battery. A fluorescent zinc sulphide screen was used to observe the rays. The rays travelled from the cathode towards the anode.

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🔬 6. Canal Ray Experiment

The canal ray experiment was carried out by E. Goldstein. It provided evidence for the presence of positively charged particles.

• Canal rays consist of positively charged particles.
• These rays travel towards the negatively charged cathode.
• The positively charged particle associated with the hydrogen atom is called a proton.

Charge of proton: +1.6 × 10⁻¹⁹ C

Mass of proton: approximately 1.67 × 10⁻²⁷ kg

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⚛️ 7. Discovery of Neutron

The neutron was discovered by James Chadwick in 1932.

• A neutron is a neutral subatomic particle.
• It is present in the nucleus of an atom.
• It has no electric charge.
• The mass of a neutron is approximately equal to the mass of a proton.

📊 Three Main Subatomic Particles

Particle	Charge	Approx. Mass	Location
Electron	−1	9.11 × 10−31 kg	Outside nucleus
Proton	+1	1.67 × 10−27 kg	Nucleus
Neutron	0	1.67 × 10−27 kg	Nucleus

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🍉 8. Thomson's Model of Atom

J. J. Thomson proposed his model of the atom in 1903. It is also known as the watermelon model or plum pudding model.

📌 Postulates

• The atom is a sphere of positive charge.
• Electrons are embedded in the positively charged sphere.
• The total positive charge is equal to the total negative charge.
• Therefore, the atom as a whole is electrically neutral.

Amount of positive charge = Amount of negative charge

❌ Drawback

Thomson's model could not explain the observations of the gold foil experiment performed by Rutherford.

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🥇 9. Rutherford's Gold Foil Experiment

Rutherford performed the famous gold foil experiment to investigate the internal structure of the atom.

🔬 Experimental Setup

• A very thin gold foil was used.
• Gold was chosen because it could be made into an extremely thin foil.
• Alpha (α) particles were directed towards the gold foil.
• A zinc sulphide screen was used to detect the alpha particles.

📌 Observations

1. Most alpha particles passed straight through the gold foil.
2. Some alpha particles were deflected through small angles.
3. A very small number of alpha particles were deflected through large angles, with some being deflected by nearly 180°.

💡 Conclusions

• Most of the atom is empty space.
• Most of the positive charge is concentrated in a very small region.
• This small, dense, positively charged region is called the nucleus.

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⚛️ 10. Rutherford's Model of Atom

• Most of the atom is hollow or empty space.
• Most of the mass of an atom is concentrated in a very small central region called the nucleus.
• The nucleus is positively charged because it contains protons.
• Electrons revolve around the nucleus.
• The nucleus contains protons and neutrons.
• The radius of the nucleus is approximately 10⁻¹⁵ m.

❌ Drawback of Rutherford's Model

According to classical electromagnetic theory, a charged particle moving in a circular orbit should continuously lose energy. Therefore, an electron revolving around the nucleus should lose energy and eventually fall into the nucleus.

This would make the atom unstable, which is not observed in reality.

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🌟 11. Bohr's Model of Atom

Neil Bohr proposed a new model of the atom to overcome the limitations of Rutherford's model.

📌 Main Postulates

• Electrons revolve around the nucleus only in certain specific orbits.
• These specific orbits are called discrete orbits, energy shells or energy levels.
• While moving in a permitted orbit, an electron does not lose its energy.

🪐 Energy Shells

Shell	Principal Energy Level
K-shell	n = 1
L-shell	n = 2
M-shell	n = 3
N-shell	n = 4

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🔢 12. Maximum Number of Electrons in a Shell

The maximum number of electrons that can be accommodated in a shell is given by:

Maximum electrons = 2n²

where n represents the shell number.

• K-shell (n = 1) → 2 × 1² = 2 electrons
• L-shell (n = 2) → 2 × 2² = 8 electrons
• M-shell (n = 3) → 2 × 3² = 18 electrons
• N-shell (n = 4) → 2 × 4² = 32 electrons

The shells are filled stepwise, starting from the inner shell and moving towards the outer shell.

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🧮 13. Electronic Configuration

The distribution of electrons among the different shells of an atom is called electronic configuration.

📌 Examples

• Hydrogen (Z = 1): K = 1 → 1
• Helium (Z = 2): K = 2 → 2
• Lithium (Z = 3): K = 2, L = 1 → 2,1
• Sodium (Z = 11): K = 2, L = 8, M = 1 → 2,8,1
• Silicon (Z = 14): 2,8,4
• Argon (Z = 18): 2,8,8

The outermost shell is especially important because its electrons determine the valency and chemical behaviour of an atom.

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🔗 14. Valence Electrons

The electrons present in the outermost shell of an atom are called valence electrons.

Example: Sodium has electronic configuration 2,8,1. Therefore, it has 1 valence electron.

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⚡ 15. Valency

Valency is the number of electrons an atom gains or loses to achieve a stable electronic configuration.

Atoms generally try to complete their outermost shell with an octet (8 electrons) or a duplet (2 electrons) in the case of the first shell.

📌 Examples

• Sodium (Na): 2,8,1 → loses 1 electron → valency = +1
• Fluorine (F): 2,7 → gains 1 electron → valency = −1
• Magnesium (Mg): 2,8,2 → loses 2 electrons → valency = +2
• Aluminium (Al): 2,8,3 → loses 3 electrons → valency = +3
• Oxygen (O): 2,6 → gains 2 electrons → valency = −2
• Chlorine (Cl): 2,8,7 → gains 1 electron → valency = −1

🌟 Noble Gases

Helium, neon and other noble gases have stable outer electronic configurations. Therefore, they generally do not gain or lose electrons and have valency 0.

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🛡️ 16. Stable Electronic Configuration

Atoms tend to gain or lose electrons to achieve the electronic configuration of the nearest noble gas.

Example:

• Sodium (Na) has atomic number 11 and configuration 2,8,1.
• It loses one electron to form Na⁺.
• Na⁺ has configuration 2,8, which is stable.

Similarly, fluorine gains one electron to form F⁻ and obtains the stable configuration 2,8.

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🔢 17. Atomic Number

The atomic number (Z) of an atom is the total number of protons present in its nucleus.

Z = Number of protons

Examples:

• Hydrogen → 1 proton → Z = 1
• Sodium → 11 protons → Z = 11

📌 Atomic Number of a Neutral Atom

In a neutral atom, the number of protons is equal to the number of electrons. Therefore:

Number of protons = Number of electrons = Atomic number

The number of electrons can change when an atom gains or loses electrons, but the atomic number is determined by the number of protons.

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⚖️ 18. Mass Number

The mass number (A) of an atom is the total number of protons and neutrons present in its nucleus.

A = Number of protons + Number of neutrons

Since:

Number of neutrons = A − Z

📝 Representation of an Atom

^A_ZX

• A = Mass number
• Z = Atomic number
• X = Symbol of the element

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🧬 19. Isotopes

Isotopes are atoms of the same element having the same atomic number but different mass numbers.

The difference in mass number occurs because the number of neutrons is different.

🔹 Hydrogen Isotopes

• Protium: ¹₁H → 1 proton, 0 neutrons
• Deuterium: ²₁H → 1 proton, 1 neutron
• Tritium: ³₁H → 1 proton, 2 neutrons

Protium is the most abundant form of hydrogen.

🔹 Oxygen Isotopes

• Oxygen-16 → 8 protons, 8 neutrons
• Oxygen-18 → 8 protons, 10 neutrons

🔹 Carbon Isotopes

• Carbon-12 → 6 protons, 6 neutrons
• Carbon-14 → 6 protons, 8 neutrons

🔹 Chlorine Isotopes

• Chlorine-35 → 17 protons, 18 neutrons
• Chlorine-37 → 17 protons, 20 neutrons

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🧮 20. Average Atomic Mass

When an element occurs naturally as a mixture of isotopes, its atomic mass is calculated using the abundance of its isotopes.

Average atomic mass = Σ (Abundance × Isotopic mass)

📌 Example: Chlorine

Chlorine occurs mainly as chlorine-35 and chlorine-37. The relative abundance of the two isotopes results in an average atomic mass of approximately 35.5 u.

Therefore, the atomic mass of chlorine is commonly represented as 35.5 u.

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🔄 21. Isobars

Isobars are atoms of different elements having the same mass number but different atomic numbers.

📌 Example

Argon and calcium can be represented as:

• ⁴⁰₁₈Ar
• ⁴⁰₂₀Ca

Both have mass number 40, but their atomic numbers are different.

Important: In isobars, the total number of nucleons (protons + neutrons) remains the same.

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🧠 Quick Revision – Journey Inside the Atom

Concept	Key Point
Electron	Negatively charged particle
Proton	Positively charged particle present in nucleus
Neutron	Neutral particle present in nucleus
Atomic Number	Number of protons
Mass Number	Protons + Neutrons
Valence Electrons	Electrons in the outermost shell
Valency	Electrons gained or lost to attain stability
Isotopes	Same atomic number, different mass numbers
Isobars	Same mass number, different atomic numbers

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🎯 Important Formulas to Remember

• ⚛️ Maximum electrons in shell = 2n²
• 🔢 Atomic number (Z) = Number of protons
• ⚖️ Mass number (A) = Protons + Neutrons
• 🧮 Number of neutrons = A − Z
• ⚡ For a neutral atom: Number of protons = Number of electrons
• 🔗 Valency depends on the number of electrons gained or lost to attain a stable configuration.

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📚 Exam-Focused Points

• ⭐ Remember the three subatomic particles: electron, proton and neutron.
• ⭐ Learn the main postulates and limitations of Dalton's theory.
• ⭐ Understand Thomson's atomic model and its drawback.
• ⭐ Learn the observations and conclusions of Rutherford's gold foil experiment.
• ⭐ Remember Bohr's energy shells and the formula 2n².
• ⭐ Practise electronic configurations of the first 20 elements.
• ⭐ Understand the difference between atomic number and mass number.
• ⭐ Learn how to calculate the number of neutrons.
• ⭐ Clearly distinguish between isotopes and isobars.
• ⭐ Practise examples involving valency and stable noble-gas configurations.