Journey Inside The Atom Class 9 Notes introduces students to the fascinating world of atoms and their structure. This chapter traces the development of atomic theory from ancient ideas to modern scientific models proposed by Dalton, Thomson, Rutherford, and Bohr. Students learn about subatomic particles, atomic number, mass number, electronic configuration, valency, isotopes, and isobars.
Journey Inside The Atom Class 9 Notes
Everything that you see or observe or feel around is matter. You have learnt that matter consists of tiny particles called atoms.
Rediscovering the Roots of Atomic Theory
More than 2,000 years ago there were great thinkers in ancient India and Greece who tried to answer a very important question: “What is everything made of?”
In ancient India, Acharya Kanada was an Indian philosopher. He suggested that if matter (dravya) is divided repeatedly, you will reach a stage where you will encounter the smallest particles that can no longer be divided. He called these particles ‘parmanus’. His ideas are recorded in the Sanskrit text Vaisheshika Sutras.
The Greek philosophers Leucippus and Democritus gave similar ideas. They called the smallest particles “atomos”. They called these indivisible particles atomos (in Greek, atomos means indivisible).
Many centuries later, in 1808, John Dalton gave the first scientific atomic theory. He proposed that all matter is made of tiny particles called atoms.
Atoms are basic building blocks of matter and cannot be broken into smaller parts. His theory was based on the experiments. Dalton’s atomic theory was the first scientific description of how matter is made.
A Short Historical Journey Through Atomic Models
More than 100 years ago, scientists were trying to understand what the atoms looked like. At that time scientists knew that atoms were the smallest particles of matter and could not be divided further.
But later discoveries changed this idea. Scientists found that some substances emit invisible energy and particles, a process called radioactivity. It showed that atoms are not the smallest parts of matter but contain even smaller particles inside them.
In 1897, J. J. Thomson conducted an important experiment using a special tube called a cathode ray tube. He passed electricity through a gas at very low pressure and observed rays moving from the negative electrode (cathode) to the positive electrode (anode).
These rays were called cathode rays. After careful study, he concluded that these rays are made of tiny negatively charged particles, which were later called electrons.
Thomson’s model of an atom
After discovering electrons, J. J. Thomson tried to explain how an atom is still neutral.
He said that an atom is like a ball of positive charge, and tiny negative electrons are spread inside it. Because both positive and negative charges are present, they cancel each other, so the atom becomes neutral.
He compares it with watermelon, where the red part is positive and the seeds are electrons. This model was simple and easy to understand, but later scientists found it was not completely correct.
Testing Thomson’s model: The gold foil experiment
In 1911, scientists Ernest Rutherford and his students tested Thomson’s model using the gold foil experiment.
They fired tiny positively charged particles (known as alpha particles) at a thin gold sheet. They expected that the particles would pass straight through, and most of them did, but some were deflected, and a few even came back.
This was surprising; it shows that the positive charge is not spread everywhere but is concentrated in a small part of the atom. This experiment proved that Thomson’s model was not correct and led to a new understanding of the atom.
A. Rutherford’s model of an atom
After the gold foil experiment, Ernest Rutherford gave a new model of the atom. He said that an atom has a positive charge and an atom is not spread everywhere. An atom has a tiny centre called a nucleus.
Most of the atom is empty space; because of this, most particles passed straight through the gold foil. The nucleus holds almost all the mass of the atom.
Electrons move around the nucleus, like planets around the sun, so that way it’s called the planetary model.
Rutherford’s atomic model was better than Thomson’s atomic model in explaining the results of the gold foil experiment.
B. Limitations of Rutherford’s model
Rutherford’s model explained the presence of a nucleus, but it had a major problem. According to the model, electrons move around the nucleus.
If the electrons are moving around the nucleus, then the electrons should lose energy and fall into the nucleus. If this were to happen, then the atoms would collapse and matter would not exist.
But in reality atoms are stable. This shows that Rutherford’s model was not completely correct.
C. Discovery of the proton
Rutherford showed that the nucleus carries positive charge which comes from the particles called protons. Protons are much heavier than electrons and possess a charge equal and opposite to that of electrons. For an atom to be electrically neutral, the number of protons must be equal to the number of electrons.
For example, a helium atom has 2 protons and 2 electrons, and a sodium atom has 11 protons and 11 electrons. Because the positive and negative charges are equal, atoms remain neutral.
Bohr’s model of the atom
In 1913, Niels Bohr gave a new model to explain why the atoms are stable. He said that electrons move around the nucleus in fixed circular paths called ‘shells’ or ‘energy levels’ (K, L, M, N). Each cell has a fixed energy, and electrons can move only in these shells.
While moving in this fixed shell, the electrons do not lose energy, which helps to keep the atom stable. The cell close to the nucleus has the least energy and energy increases when moving farther away. Electrons can jump from one shell to another by gaining or losing energy.
This model successfully explained the stability of atoms
What Components Contribute to the Mass of an Atom?
Most of an atom’s mass is concentrated in the nucleus and mainly comes from protons and neutrons.
1. Protons
- Positively charged particles
- Each proton has significant mass.
- Found inside the nucleus
2. Neutrons
- No charge (neutral)
- Almost the same mass as protons
- Also present in the nucleus
Why is helium 4 times heavier than hydrogen?
Hydrogen atoms have 1 proton and 0 neutrons; in helium atoms there are 2 protons and 2 neutrons. Hydrogen mass has 1 unit, and helium mass has 4 units. That extra mass comes from neutrons, even though they don’t affect charge.
Discovery of the Neutron
In 1932, the mystery of atomic mass was solved by James Chadwick, a student of Ernest Rutherford. He discovered a new subatomic particle called the neutron.
- Has no electrical charge (neutral)
- Has a mass nearly equal to a proton
- Is represented by the symbol n⁰
Symbols and relative charges of subatomic particles
| S.No. | Subatomic particle | Symbol | Relative charge |
|---|---|---|---|
| 1. | Electron | e– | -1 |
| 2. | Proton | p+ | +1 |
| 3. | Neutron | n0 | 0 |
By 1869, scientists knew about 69 elements, most of which were found naturally on the Earth. Today, we know about 118 unique chemical elements. Some of these are artificially made, and the search for even more continues.
Symbols of Elements
To make chemistry simple and universal, scientists needed a standard way to represent elements.
Early Development
- In 1803, John Dalton introduced pictorial symbols for elements.
- Later, in 1813, Jöns Jakob Berzelius proposed using letters instead
- Today, symbols are approved by the International Union of Pure and Applied Chemistry (IUPAC).
Rules for Writing Symbols
- Symbols are usually the first letter or first two letters, like Hydrogen -> H and Aluminium -> Al.
- The first letter will be capital and the second letter will be small, like Cobalt (Co).
- Sometimes symbols use other letters from the name, like Chlorine -> Cl and Zinc -> Zn.
Names of some common elements and their symbols
Scientists use these symbols instead of full names because they are internationally recognised and allow scientists worldwide to communicate clearly, regardless of language barriers.
Atomic Number
The atomic number of an element is the number of protons present in the nucleus of its atom.
- It is denoted by the symbol ‘Z’.
- It determines the identity of an element.
- It also affects its chemical behaviour.
Why is the atomic number important?
Each element has a unique atomic number; there are no two elements that can have the same atomic number. That’s why the atomic number is like an identity card of an element.
Mass Number
The mass number of an atom is the total number of protons and neutrons present in its nucleus.
- It is denoted by the symbol A
- Protons and neutrons together are called nucleons
Mass Number (A) = Number of Protons + Number of Neutrons
- The mass of an atom mainly comes from protons and neutrons
- The mass of an electron is negligible, so it is ignored
- Since protons and neutrons have nearly equal mass, they together account for atomic mass
Standard Notation of an Atom
An element is written as:
Where:
A = Mass number
Z = Atomic number
X = Symbol of element
For example, the symbol for carbon is C, its atomic number is 6, and its mass number is 12. In notation, it would be written as —
How Are Electrons Distributed in Different Energy Levels?
Scientists like Niels Bohr and Charles Bury gave rules to understand how electrons are arranged in shells (energy levels).
1. Maximum electrons in a shell
- The number of electrons a shell can hold is given by:
- 2n2
- Where n = shell number
| Shell | n | Maximum electrons |
|---|---|---|
| K | 1 | 2 |
| L | 2 | 8 |
| M | 3 | 18 |
| N | 4 | 32 |
2. Outermost shell rule
- The outermost shell can have a maximum of 8 electrons
- Exception: First shell (K) can have only 2 electrons
3. Stepwise filling (order)
- Electrons fill shells from inner to outer:
- K → L → M → N
- A shell is filled only after the previous one is complete
Building up atoms
The distribution of electrons among various shells is known as the electronic configuration of the atom. following lists the symbols, atomic numbers, number of protons, number of neutrons, number of electrons, and the electron distribution in the shells of the first eighteen elements.
Symbols, atomic numbers, number of protons, number of neutrons, number of electrons, and the electronic distribution of atoms of the first eighteen elements
Combining Capacity of an Atom: Valency
The combining capacity of an atom is called its valency. It tells us:
- How many atoms (like hydrogen or chlorine) an element can combine with
- Or how many electrons it can lose, gain, or share
- In H₂O → Oxygen combines with 2 hydrogen atoms
- Valency of oxygen = 2
- In NH₃ → Nitrogen combines with 3 hydrogen atoms
- Valency of nitrogen = 3
- In MgCl₂ → Magnesium combines with 2 chlorine atoms
- Valency of magnesium = 2
The outermost shell containing electrons of an atom is known as its valence shell. The electrons present in it are known as valence electrons.
A Deeper Look into Atomic Structure
Isotopes
Earlier, John Dalton said all atoms of an element are identical. But later, scientists discovered this is not completely true. Atoms of the same element can have:
- Same number of protons (atomic number Z)
- Different number of neutrons
These are called isotopes.
Isotopes = Atoms of the same element with the same atomic number but different mass numbers, Because Mass number = Protons + Neutrons
A. Average atomic mass
When an element has isotopes, its atomic mass is not a whole number. This is because different isotopes exist in different proportions in nature.
We calculate atomic mass using percentage abundance:
Solving:
- (75/100 × 35) = 26.25
- (25/100 × 37) = 9.25
Total = 35.5 u
Isobars
Isobars are atoms of different elements that have:
- Same mass number (A)
- Different atomic number (Z)
| Element | Atomic Number (Z) | Mass Number (A) |
|---|---|---|
| Calcium (Ca) | 20 | 40 |
| Potassium (K) | 19 | 40 |
| Argon (Ar) | 18 | 40 |
Same mass number = 40, But different Number of protons
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