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Journey Inside the Atom Class 9 Science Notes

Journey Inside the Atom

Everything we see, feel, and observe is matter, which is composed of tiny particles called atoms. These atoms are the fundamental building blocks of both living beings and non-living things. For a long time, people wondered if atoms were truly the smallest, indivisible units of matter, or if they were made of even smaller components. This curiosity led scientists on a long journey to uncover the structure of the atom.

Rediscovering the Roots of Atomic Theory

The idea of atoms is not new. Over 2,000 years ago, thinkers in ancient India and Greece pondered the nature of matter.

In India, Acharya Kanada proposed that if you keep dividing matter (dravya), you will eventually reach the smallest, indivisible particles, which he called parmanus.
In ancient Greece, Leucippus and Democritus had a similar idea, calling these indivisible particles atomos (which means "indivisible" in Greek).

These early concepts were philosophical ideas, not based on experiments. The first scientific theory of the atom was proposed by John Dalton in 1808. Based on the scientific experiments of his time, Dalton's atomic theory stated that all matter is made of indivisible particles called atoms. This theory became the starting point for our modern understanding of atomic structure.

A Short Historical Journey Through Atomic Models

The discovery that atoms were not indivisible began with the study of radioactivity, a phenomenon where certain elements emit energy and particles. This proved that atoms must contain smaller particles.

Discovery of the Electron

In 1897, J. J. Thomson conducted experiments by passing an electric current through gases at very low pressure in a glass tube. He observed rays traveling from the negative electrode (cathode) to the positive electrode (anode). These were called cathode rays.

By studying how these rays behaved in electric and magnetic fields, Thomson concluded that they were streams of negatively charged particles. These particles were much smaller than atoms and were later named electrons (symbol $e^{-}$ ). Since these electrons were emitted from atoms and were the same regardless of the gas or electrode material used, Thomson realised that electrons are a fundamental component of all atoms. By convention, the relative charge of an electron is taken as -1.

Thomson's model of an atom

The discovery of the negative electron posed a new puzzle: since atoms are electrically neutral, where is the positive charge? To answer this, Thomson proposed the first model of the atom, often called the plum pudding model.

He imagined the atom as a sphere of uniform positive charge.
The negatively charged electrons were scattered or embedded throughout this positive sphere, like plums in a pudding or seeds in a watermelon.
This arrangement ensured that the total positive charge balanced the total negative charge, making the atom neutral.

Example

A watermelon can be used as an analogy for Thomson's model. The red pulp represents the sphere of positive charge, and the black seeds represent the electrons distributed within it.

Testing Thomson's model: The gold foil experiment

In 1911, Ernest Rutherford directed his colleagues, Geiger and Marsden, to test Thomson's model. They performed the famous gold foil experiment.

Setup: A narrow beam of fast-moving, positively charged alpha particles ( $\alpha$ ) was aimed at an extremely thin sheet of gold foil.
Expectation: According to Thomson's model, where the positive charge is spread out, the alpha particles should have passed straight through the foil with only minor deflections.
Observation: The results were surprising:
- Most alpha particles passed straight through the foil undeflected.
- Some particles were deflected by small angles.
- A very small number of particles were deflected by large angles, and a few even bounced straight back.

This phenomenon of deflection is called scattering. The results could not be explained by Thomson's model, which was therefore rejected.

A. Rutherford's model of an atom

Based on the results of the gold foil experiment, Rutherford proposed a new atomic model, known as the planetary model.

Nucleus: The positive charge and most of the mass of an atom are concentrated in a very small, dense region at the center called the nucleus. The fact that a few alpha particles bounced back meant they had hit something very dense and positively charged.
Empty Space: Most of the atom is empty space. This explained why most alpha particles passed straight through the foil.
Electrons: The negatively charged electrons revolve around the nucleus in circular paths, much like planets orbiting the Sun.

Rutherford calculated that the nucleus is about $10^5$ times smaller than the atom itself. If an atom were the size of a cricket ground, the nucleus would be like a tiny grain of pepper at the center.

B. Limitations of Rutherford's model

While a major step forward, Rutherford's model had a critical flaw: it could not explain why atoms are stable.

According to the laws of physics, a charged particle moving in a circular path is accelerating. An accelerating electron should constantly lose energy by emitting radiation. This loss of energy would cause it to spiral inward and eventually crash into the positively charged nucleus. If this were true, all atoms would collapse. Since matter is stable, this part of Rutherford's model had to be incorrect.

C. Discovery of the proton

Rutherford showed that the positive charge in the nucleus comes from particles called protons (symbol $p^{+}$ ).

A proton is much heavier than an electron.
It has a positive charge that is equal in magnitude but opposite in sign to the charge of an electron. Its relative charge is +1.
For an atom to be electrically neutral, the number of protons must equal the number of electrons.

Bohr's model of the atom

In 1913, Niels Bohr proposed a new model to address the instability of Rutherford's model.

Stationary States: Bohr suggested that electrons do not move randomly but follow fixed circular paths called orbits or shells. These are also known as energy levels because an electron in a particular shell has a definite, constant amount of energy.
No Energy Loss: While revolving in these specific "allowed" shells, an electron does not lose or radiate energy. This was a key postulate that explained the stability of atoms.
Shell Notation: The shells are represented by the letters K, L, M, N,... or by the numbers $n=1, 2, 3, 4,...$ . The K-shell ( $n=1$ ) is closest to the nucleus and has the lowest energy.
Energy Levels: The energy of the shells increases as they get farther from the nucleus (Energy: K < L < M < N...).
Electron Jumps: An electron can move to a higher energy shell by absorbing a specific amount of energy or move to a lower shell by releasing energy.

Bohr's model successfully explained the stability of the atom and was a significant step in understanding atomic structure.

Note

The naming of shells as K, L, M, N... was adopted from early X-ray experiments by physicist Charles Barkla. He started with 'K' to leave room for the possible discovery of shells with even lower energy, though none were ever found.

What Components Contribute to the Mass of an Atom?

Rutherford's work showed that most of an atom's mass is in its nucleus. However, a puzzle remained. A helium atom has two protons, but its mass is about four times that of a hydrogen atom (which has one proton). This suggested there must be another particle in the nucleus that adds mass but has no charge.

Discovery of the Neutron

In 1932, James Chadwick, a student of Rutherford, solved this puzzle. He discovered a new subatomic particle in the nucleus with no electrical charge and a mass nearly equal to that of a proton. This neutral particle was named the neutron (symbol $n$ or $n^0$ ).

Neutrons are found in the nucleus of all atoms except for the most common form of hydrogen. The mass of an atom is therefore mainly the sum of the masses of its protons and neutrons.

The three main subatomic particles are:

S.No.	Subatomic particle	Symbol	Relative charge
1.	Electron	$e^{-}$	-1
2.	Proton	$p^{+}$	+1
3.	Neutron	$n^{0}$	0

Symbols of Elements

To simplify communication in chemistry, a standard system for representing elements was developed.

John Dalton first used pictorial symbols.
In 1813, Berzelius suggested using one or two letters from the element's name.
Today, the International Union of Pure and Applied Chemistry (IUPAC) approves the names and symbols for elements.

Rules for Writing Symbols:

The symbol is often the first one or two letters of the element's English name.
The first letter is always uppercase, and the second letter (if present) is always lowercase. For example, Aluminum is Al (not AL), and Cobalt is Co (not CO).
Some symbols are derived from the element's name in Latin, Greek, or German. For example, Iron is Fe (from ferrum), and Sodium is Na (from natrium).

Atomic Number

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

The atomic number uniquely identifies an element. All atoms of a given element have the same number of protons.
Since a neutral atom has an equal number of protons and electrons, the atomic number also equals the number of electrons in a neutral atom.
Atomic Number (Z) = Number of Protons

Mass Number

The mass number (symbol A) of an atom is the total number of protons and neutrons in its nucleus. Protons and neutrons together are called nucleons.

Mass Number (A) = Number of Protons + Number of Neutrons
The mass of electrons is negligible and is not included in the mass number.

An atom is represented using standard notation:

\text{Mass Number} \atop \text{Atomic Number} \text{Symbol of element} \quad \text{or} \quad { }_{Z}^{A}X

Example

For carbon, which has 6 protons and 6 neutrons, the atomic number (Z) is 6 and the mass number (A) is 12. Its symbol is written as

{ }_{6}^{12}\text{C}

.

How Are Electrons Distributed in Different Energy Levels?

The arrangement of electrons in the various shells of an atom is called its electronic configuration. The rules for filling electrons, proposed by Bohr and Bury, are:

The maximum number of electrons a shell can hold is given by the formula $\boldsymbol{2n^2}$ $2 n^{2}$ , where 'n' is the shell number.
- K-shell (n=1): can hold up to $2 \times 1^2 = 2$ electrons.
- L-shell (n=2): can hold up to $2 \times 2^2 = 8$ electrons.
- M-shell (n=3): can hold up to $2 \times 3^2 = 18$ electrons.
The outermost shell of an atom cannot accommodate more than 8 electrons (this is often called the octet rule).
Electrons fill the shells in a stepwise manner, starting from the innermost shell (K) and moving outwards. A new shell only begins to fill after the inner shell is complete.

Combining Capacity of an Atom: Valency

The valency of an element is its combining capacity. It is determined by the number of electrons in its outermost shell, which is known as the valence shell. The electrons in this shell are called valence electrons.

Atoms tend to achieve a stable configuration, which usually means having 8 electrons in their valence shell (an octet). To do this, they can lose, gain, or share electrons.

Valency is the number of electrons an atom loses, gains, or shares to achieve a stable octet.
If an atom has 1, 2, or 3 valence electrons: It tends to lose them. Its valency is equal to the number of electrons it loses (e.g., Sodium with configuration 2, 8, 1 has a valency of 1).
If an atom has 5, 6, or 7 valence electrons: It tends to gain electrons to complete its octet. Its valency is the number of electrons it gains (e.g., Oxygen with configuration 2, 6 needs 2 electrons, so its valency is 2).
If an atom has 4 valence electrons: It tends to share electrons (e.g., Carbon with configuration 2, 4 shares 4 electrons, so its valency is 4).
If an atom has 8 valence electrons: Its outermost shell is already full. These elements (like Neon and Argon) are very stable, unreactive, and have a valency of 0.

A Deeper Look into Atomic Structure

Isotopes

Dalton's theory claimed all atoms of an element are identical. However, we now know this is not entirely true.

Isotopes are atoms of the same element that have the same atomic number (Z) but different mass numbers (A).

This means isotopes have the same number of protons but a different number of neutrons.
Because they have the same number of electrons, isotopes of an element have similar chemical properties.
They have different physical properties (like mass, boiling point) due to the different number of neutrons.

Example

Hydrogen Isotopes:

Protium ( ${ }_{1}^{1}\text{H}$ ): 1 proton, 0 neutrons
Deuterium ( ${ }_{1}^{2}\text{H}$ ): 1 proton, 1 neutron
Tritium ( ${ }_{1}^{3}\text{H}$ ): 1 proton, 2 neutrons

Applications of Isotopes:

An isotope of uranium ( ${ }_{92}^{235}\text{U}$ ) is used as fuel in nuclear reactors.
An isotope of cobalt ( ${ }_{27}^{60}\text{Co}$ ) is used in cancer treatment.
An isotope of iodine ( ${ }_{53}^{131}\text{I}$ ) is used to treat thyroid disorders.
An isotope of carbon ( ${ }_{6}^{14}\text{C}$ ) is used in radiocarbon dating to find the age of fossils.

Average atomic mass

The atomic mass listed on the periodic table is a decimal because it is the weighted average of the masses of an element's naturally occurring isotopes, taking their abundance into account.

Example

Calculating the Average Atomic Mass of Chlorine Chlorine exists in nature as two isotopes:

{ }^{35}\text{Cl}

(mass 35 u, abundance 75%) and

{ }^{37}\text{Cl}

(mass 37 u, abundance 25%). Calculate the average atomic mass.

Given

Mass of isotope 1 = 35 u
Abundance of isotope 1 = 75% or $\frac{75}{100}$
Mass of isotope 2 = 37 u
Abundance of isotope 2 = 25% or $\frac{25}{100}$

To Find

The average atomic mass of the chlorine atom.

Formula

$\text{Average atomic mass} = (\text{Mass}_1 \times \text{Abundance}_1) + (\text{Mass}_2 \times \text{Abundance}_2)$

Solution

Substitute the given values into the formula: $\text{Average atomic mass} = \left(35 \times \frac{75}{100} + 37 \times \frac{25}{100}\right)$ $= \left(\frac{105}{4} + \frac{37}{4}\right) \text{u}$ $= \frac{142}{4} \text{u}$ $= 35.5 \text{u}$

Final Answer The average atomic mass of chlorine is 35.5 u. This doesn't mean a single chlorine atom has a mass of 35.5 u, but that a sample of chlorine atoms has this average mass.

Isobars

Isobars are atoms of different elements that have different atomic numbers (Z) but the same mass number (A).

This means they have a different number of protons (and are therefore different elements) but the same total number of nucleons (protons + neutrons).

Example

Calcium (atomic number 20) and Argon (atomic number 18) both have an isotope with a mass number of 40. Thus,

{ }_{20}^{40}\text{Ca}

and

{ }_{18}^{40}\text{Ar}

are isobars.

Journey Inside the AtomClass 9 Science Notes

Journey Inside the Atom

Rediscovering the Roots of Atomic Theory

A Short Historical Journey Through Atomic Models

Discovery of the Electron

Thomson's model of an atom

Testing Thomson's model: The gold foil experiment

A. Rutherford's model of an atom

B. Limitations of Rutherford's model

C. Discovery of the proton

Bohr's model of the atom

What Components Contribute to the Mass of an Atom?

Discovery of the Neutron

Symbols of Elements

Atomic Number

Mass Number

How Are Electrons Distributed in Different Energy Levels?

Combining Capacity of an Atom: Valency

A Deeper Look into Atomic Structure

Isotopes

Average atomic mass

Given

To Find

Formula

Solution

Isobars