Everything in the universe that has mass and occupies space (volume) is called matter. Air, water, plants, animals, rocks, and even our own body are all made of matter.

The smallest building block of matter is called an atom. Atoms are extremely tiny particles that combine to form all substances around us. Although atoms are too small to be seen with the naked eye, they are the basic units from which every element and compound is made.

The idea of the atom is not new. Around 2000 years ago, ancient Indian philosophers described the smallest indivisible particle of matter as "Parmanu" in Sanskrit texts. Around the same time, the Greek philosopher Democritus used the word "Atomos", meaning "indivisible".

In 1808, the English scientist John Dalton proposed the modern Atomic Theory. According to Dalton, all matter is made up of extremely small, indivisible particles called atoms. His theory laid the foundation of modern atomic science.

After John Dalton proposed his Atomic Theory in 1808, scientists believed that an atom was the smallest and indivisible particle of matter. According to Dalton, atoms could not be divided into smaller parts.

For many years, scientists accepted Dalton's idea. However, towards the end of the 19th century, two important discoveries completely changed this belief.

1. Discovery of Radioactivity

Scientists discovered that some elements, such as uranium, naturally emitted invisible rays. This phenomenon was called radioactivity.

If atoms were truly indivisible, they should have remained unchanged. But the emission of these rays suggested that something was happening inside the atom.

2. J. J. Thomson's Cathode Ray Experiment

Around the same time, the English scientist J. J. Thomson performed the famous Cathode Ray Experiment. His experiment led to the discovery of the electron, the first subatomic particle.

This proved that atoms were not indivisible. Instead, they contained smaller particles within them.

The discovery of subatomic particles completely changed our understanding of the atom. Scientists proved that atoms are not indivisible but are made up of smaller particles called subatomic particles. The three fundamental subatomic particles are the electron, proton and neutron.

1. Discovery of Electron – J. J. Thomson (1897)

In 1897, the English scientist J. J. Thomson studied the conduction of electricity through gases at a very low pressure. He used a glass discharge tube fitted with two electrodes—a cathode (−) and an anode (+).

When a very high voltage was applied across the electrodes, a beam of invisible rays travelled from the cathode towards the anode. These rays were called cathode rays.

Observations

• Cathode rays travelled from the cathode towards the anode.
• They were deflected towards the positive plate, showing that they carried a negative charge.
• The nature of cathode rays was the same irrespective of the gas taken inside the tube.
• The rays were also independent of the material used for the cathode.

Conclusion

J. J. Thomson concluded that cathode rays are streams of tiny negatively charged particles called electrons. Since every atom produced these particles, electrons are present in all atoms.

2. Discovery of Proton – Anode Rays / Canal Rays (Goldstein, Rutherford)

After the discovery of electrons, scientists wondered whether atoms also contained positively charged particles.

In 1886, Eugen Goldstein modified the discharge tube by using a perforated cathode (a cathode having tiny holes). During the experiment, he observed another type of rays moving in the opposite direction. These rays passed through the holes in the cathode and were called canal rays or anode rays.

Observations

• Canal rays travelled from the anode towards the cathode.
• They carried a positive charge.
• Their properties depended on the gas present inside the tube.
• They were much heavier than electrons.

Conclusion

Goldstein's experiment proved the existence of positively charged particles in atoms. Later, Ernest Rutherford identified these particles as protons.

3. Discovery of Neutron – James Chadwick (1932)

Although electrons and protons had been discovered, scientists could not explain the total mass of the atom.

In 1932, the English physicist James Chadwick discovered a new particle called the neutron.

Neutrons have no electric charge and their mass is almost equal to that of a proton. They are present inside the nucleus of the atom along with protons.

After the discovery of the electron, proton and neutron, a new question arose:

To answer this question, scientists proposed different atomic models. Each new model tried to explain the structure of the atom more accurately than the previous one.

1. Thomson's Model of Atom (1904)

After J. J. Thomson discovered the negatively charged electron, he faced an important question: atoms are electrically neutral, so where is the positive charge present?

To explain this, Thomson proposed that an atom is a sphere of positive charge in which electrons are embedded throughout the atom.

This model is commonly known as the Plum Pudding Model.

Main Features of Thomson's Model

• The atom is a sphere of positive charge.
• Electrons are embedded throughout the positively charged sphere.
• The positive and negative charges balance each other, making the atom electrically neutral.
• There is no separate nucleus in Thomson's model.

2. Rutherford's Nuclear Model of Atom (1911)

Ernest Rutherford proposed the nuclear model of the atom based on the results of his famous α-particle scattering experiment.

α-Particle Scattering Experiment

Rutherford directed a beam of high-energy α-particles, which are positively charged, towards a very thin sheet of gold foil.

A fluorescent screen was placed around the gold foil to detect the paths of the α-particles after they passed through or were deflected by the foil.

Observations

• About 99.9% of α-particles passed straight through the gold foil without any deflection.
• A small number of α-particles were deflected through small angles.
• A very small number of α-particles were deflected backwards or rebounded.

Conclusions

1. Most of the atom is empty space.
   Since most α-particles passed straight through the foil, most of the space inside an atom must be empty.
2. There is a small positively charged nucleus.
   The deflection of some α-particles showed that positive charge is concentrated in a very small region at the centre of the atom.
3. The nucleus is very small and dense.
   The few α-particles that bounced back showed that the nucleus is extremely small, dense and contains almost all the mass of the atom.

Main Features of Rutherford's Model

1. Nuclear Region

The atom has a very small, dense region at its centre called the nucleus. It contains the positive charge and almost all the mass of the atom.

2. Extranuclear Region

Most of the atom is empty space. Electrons revolve around the nucleus in this region. This arrangement was compared with planets revolving around the Sun, so Rutherford's model is also called the Planetary Model.

3. Bohr's Model of the Atom (1913)

Rutherford's model successfully explained the presence of the nucleus and showed that most of the atom is empty space. However, it could not explain the stability of the atom.

To overcome this problem, Danish physicist Niels Bohr proposed a new model of the atom in 1913.

Main Postulates of Bohr's Model

1. Fixed Circular Paths

Electrons do not move randomly around the nucleus. They revolve only in certain fixed circular paths called orbits, shells, energy levels or stationary states. Each shell has a definite amount of energy.

2. Representation of Energy Levels

The shells are represented by the letters: K, L, M, N, ...

They are also represented by numbers: n = 1, 2, 3, 4, ...

3. No Energy Loss in a Fixed Shell

Electrons can move only in the allowed fixed shells. While revolving in a fixed shell, an electron does not lose energy. This explains the stability of atoms.

4. Movement of Electrons Between Shells

An electron can move from one energy level to another by absorbing or releasing a definite amount of energy.

⚡ Electron Transition

Absorption of energy → Electron moves to a higher energy level.

Release of energy → Electron moves to a lower energy level.

The energy absorbed or released is equal to the difference between the energies of the two shells.

5. Electron Capacity of Shells

Each shell can accommodate only a certain maximum number of electrons. The maximum number of electrons that can be accommodated in a shell is given by:

Maximum number of electrons = 2n²

Here, n is the shell number.

📊 Electron Capacity of Shells

Shell	n	2n²	Maximum Electrons
K	1	2 × 1²	2
L	2	2 × 2²	8
M	3	2 × 3²	18
N	4	2 × 4²	32

How Does Bohr's Model Explain the Stability of Atom?

According to Rutherford's model, an electron revolving around the nucleus should continuously lose energy and eventually fall into the nucleus. This would make the atom unstable.

Bohr introduced the idea of stationary states. Electrons in these permitted energy levels have definite energy and do not lose energy while remaining in the same shell.

⭐ Key Idea

Electrons in fixed energy levels do not continuously lose energy. Therefore, the atom remains stable.

🌟 Remember

Thomson → Positive sphere + electrons embedded in it

Rutherford → Tiny positive nucleus + mostly empty space + electrons around the nucleus

Bohr → Electrons occupy fixed energy levels or shells and do not lose energy while remaining in the same shell

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