1. Rediscovering the Roots of Atomic Theory

Everything that you see or observe, or feel around is matter. Matter consists of tiny particles called atoms. Since ancient times, philosophers and scientists have pondered the fundamental building blocks of matter.

1.1 Ancient Perspectives

• Acharya Kanada (India): Suggested that if matter (dravya) is divided repeatedly, a stage is reached where the smallest particles cannot be divided further. He called these indivisible particles parmanus. His ideas are recorded in the Vaisesika Sutras.
• Leucippus and Democritus (Greece): Proposed a similar idea, calling these indivisible particles atomos (a Greek word meaning indivisible).

These early concepts were imaginary ideas rather than being based on experimental observations.

1.2 Dalton's Atomic Theory (1808)

John Dalton proposed the first scientific description of how matter is made. He suggested that all matter is composed of indivisible particles called atoms. Atoms are the fundamental building blocks of matter that cannot be broken down into smaller parts.

2. A Short Historical Journey Through Atomic Models

More than a hundred years ago, scientists tried to imagine what atoms might look like by proposing simple models. Driven by curiosity and experimentation, these models were continuously refined.

2.1 Discovery of the Electron (J.J. Thomson, 1897)

J.J. Thomson studied the conduction of electric current through gases at very low pressure using a glass tube with two electrodes and a high voltage. He observed rays moving from the cathode (negative electrode) to the anode (positive electrode), called cathode rays.

Thomson concluded that cathode rays are streams of negatively charged particles with a much smaller mass than atoms. These particles, later called electrons ($e^-$), are a fundamental component of all atoms. The charge of an electron is approximately $-1.602 \times 10^{-19} \text{ C}$ but is taken as $-1$ by convention.

2.2 Thomson’s Model of an Atom

Since atoms are electrically neutral and contain negative electrons, there must be a positive charge to balance them. Thomson proposed that an atom is a sphere of positive charge with electrons distributed throughout it.

• Plum Pudding Model: This model was compared to a pudding with plums (electrons) embedded in it.
• Watermelon Analogy: The red pulp represents the positively charged matter, and the seeds represent electrons.

2.3 Testing Thomson’s Model: The Gold Foil Experiment (1911)

Geiger and Marsden, working under Ernest Rutherford, aimed a narrow beam of alpha ($\alpha$) particles at an extremely thin sheet of gold foil. Alpha particles are positively charged particles (helium nuclei, containing two protons and two neutrons).

Expectation: If Thomson's model was correct, the positive charge would be spread evenly, so the alpha particles should pass straight through or deflect only slightly.

Observation: While most particles passed straight through, some were sharply deflected, and a few even bounced back (scattering). This proved Thomson's model wrong.

2.4 Rutherford’s Planetary Model of an Atom

Based on the gold foil experiment, Rutherford concluded that the positive charge is not spread out but concentrated in an extremely small region called the nucleus. His model proposed:

1. Most of an atom is empty space (since most $\alpha$-particles passed undeflected).
2. The nucleus is dense, containing all the positive charge and most of the mass.
3. The electrons revolve around the nucleus, somewhat like planets orbiting the Sun (hence the name planetary model).

2.5 Limitations of Rutherford’s Model

A particle moving in a circular path is constantly accelerating. A negatively charged electron accelerating around a nucleus should continuously lose energy and spiral inward, eventually crashing into the positively charged nucleus. If this happened, atoms would collapse. Since atoms are stable, a new explanation was needed.

2.6 Discovery of the Proton

Rutherford showed that the nucleus carries positive charge due to particles called protons ($p^+$). Protons are much heavier than electrons and possess a charge equal and opposite to electrons ($+1$). For an atom to be electrically neutral, the number of protons must equal the number of electrons.

3. Bohr’s Model of the Atom (1913)

To explain why atoms are stable, Niels Bohr proposed a new model with specific rules:

• Electrons do not move randomly; they follow fixed circular paths called stationary states, orbits, or shells.
• In each shell, an electron has a definite amount of energy, so these shells are also called energy levels.
• Shells are represented by letters (K, L, M, N...) or numbers ($n = 1, 2, 3, 4...$).
• Electrons can revolve only in these allowed shells, not between them. While in a fixed shell, an electron does not lose energy.
• The first energy level K ($n = 1$) is closest to the nucleus and has the least energy. Energy increases as we move further from the nucleus.

4. Components that Contribute to Mass: Discovery of the Neutron

A hydrogen atom has one proton. A helium atom has two protons. However, the mass of a helium atom is about four times that of hydrogen, not double! This puzzle led to the discovery of another particle.

In 1932, James Chadwick discovered a neutral subatomic particle with a mass nearly equal to that of a proton. He named it the neutron ($n^0$). Neutrons are found in the nucleus of all atoms (except hydrogen). The mass of an atom comes mainly from protons and neutrons packed tightly in the nucleus.

Subatomic Particles Summary:

Subatomic particle	Symbol	Relative charge
Electron	$e^-$	$-1$
Proton	$p^+$	$+1$
Neutron	$n^0$	$0$

5. Symbols of Elements

John Dalton was the first to use pictorial symbols. Later, Berzelius suggested using letters from the elements' Latin names. Today, IUPAC approves chemical symbols based on the following rules:

• Many symbols use the first letter or the first two letters of the name.
• The first letter is always uppercase, and the second is lowercase (e.g., Hydrogen: H, Aluminium: Al, Cobalt: Co).
• Some symbols come from Latin, Greek, or German names (e.g., Iron is Fe from ferrum, Sodium is Na from natrium, Potassium is K from kalium, Gold is Au from aurum).

6. Atomic Number and Mass Number

6.1 Atomic Number ($Z$)

The number of protons in the nucleus of an atom is known as its atomic number. It determines the identity of an element. In a neutral atom, the number of protons equals the number of electrons.

6.2 Mass Number ($A$)

The total number of protons and neutrons present in the nucleus is called the mass number. Protons and neutrons together are called nucleons.

Standard Notation: An element's symbol, mass number, and atomic number are written as:

7. Distribution of Electrons in Different Energy Levels

Bohr and Bury suggested rules for filling electrons in orbits:

1. The maximum number of electrons in a shell is given by the formula $2n^2$, where $n$ is the shell number.
   • K-shell ($n=1$): $2(1)^2 = 2$ electrons
   • L-shell ($n=2$): $2(2)^2 = 8$ electrons
   • M-shell ($n=3$): $2(3)^2 = 18$ electrons
2. The maximum number of electrons accommodated in the outermost shell is always 8 (except the first shell, which holds 2).
3. Electrons are filled in a stepwise manner. The inner shells must be filled before electrons occupy the outer shells.

8. Valency: Combining Capacity of an Atom

The outermost shell of an atom is known as its valence shell, and the electrons in it are valence electrons.

An outermost shell with 8 electrons is called an octet. Elements with a complete octet (like noble gases) are largely unreactive and stable. Atoms with incomplete valence shells tend to lose, gain, or share electrons to complete their octet.

The number of electrons gained, lost, or shared to complete the octet is the valency (combining capacity) of the element.

• If valence electrons $\le 4$: Valency is generally equal to the number of valence electrons (it loses or shares them). e.g., Sodium (2,8,1) has valency 1. Carbon (2,4) has valency 4.
• If valence electrons $> 4$: Valency is generally $8 - (\text{number of valence electrons})$ (it gains electrons). e.g., Oxygen (2,6) has valency $8 - 6 = 2$.

9. A Deeper Look: Isotopes and Isobars

9.1 Isotopes

Atoms of the same element that have the same atomic number ($Z$) but different mass numbers ($A$) are called isotopes. This means they have the same number of protons but a different number of neutrons.

Examples:

• Hydrogen: Protium ($_{1}^{1}\text{H}$), Deuterium ($_{1}^{2}\text{H}$), Tritium ($_{1}^{3}\text{H}$).
• Carbon: $_{6}^{12}\text{C}$, $_{6}^{13}\text{C}$, $_{6}^{14}\text{C}$.
• Chlorine: $_{17}^{35}\text{Cl}$, $_{17}^{37}\text{Cl}$.

Applications of Radioactive Isotopes:

• Uranium-235 ($^{235}\text{U}$) is used as fuel in nuclear reactors.
• Cobalt-60 ($^{60}\text{Co}$) is used in radiation treatment for cancer.
• Iodine-131 ($^{131}\text{I}$) is used to treat goitre and thyroid cancer.
• Carbon-14 ($^{14}\text{C}$) is used in carbon dating for archaeology and geology.

9.2 Average Atomic Mass

Because elements exist in isotopic forms in nature, we use a weighted average atomic mass based on the relative abundance of each isotope.

For example, Chlorine exists as $75\%$ $^{35}\text{Cl}$ and $25\%$ $^{37}\text{Cl}$.

9.3 Isobars

Atoms of different elements that have different atomic numbers but the same mass number are called isobars. This means the total number of nucleons (protons + neutrons) is the same, but the elements themselves are entirely different.

Example: Calcium ($Z=20$), Potassium ($Z=19$), and Argon ($Z=18$) can all have a mass number of $40$.