Journey Inside the Atom

Class 09 Science

Matter consists of tiny particles called atoms. In 1808, John Dalton proposed his atomic theory. It was based on scientific experiments of that time. He proposed that all matter is composed of indivisible particles called atoms.

The atoms are the fundamental building blocks of matter that cannot be broken down into smaller parts. Dalton’s atomic theory was the first scientific description of how matter is made. It became the starting point for the current understanding of atomic structure.

Until the late 19th century, atoms were thought to be the smallest, indivisible units of matter. However, scientists discovered that certain elements emit invisible energy and particles called radiation, a phenomenon known as radioactivity. This showed that atoms must be composed of smaller particles, proving that they were not indivisible as previously believed.

In 1897, J. J. Thomson studied the conduction of electric current through gases at a very low pressure. He used a glass tube with two electrodes and applied a high voltage. He observed rays moving from the cathode (negative electrode) to the anode (positive electrode). These were called cathode rays.

By studying these cathode rays in electric and magnetic fields, he concluded that they are streams of negatively charged particles, with a much smaller mass than atoms. These particles, later called electrons, were emitted from atoms, indicating that atoms are composed of smaller subatomic components.

It was found that the nature of cathode rays was independent of the material of the cathode and the gas filled in the cathode ray tube. It showed that electrons are a fundamental component of all atoms, present in every element. The charge of an electron (-1.602 × 10^-19 C) is taken as -1 as a matter of convention and convenience.

Thomson's Model of an Atom

When J. J. Thomson discovered tiny negatively charged particles called electrons, he faced a puzzle - atoms are neutral, so where is the positive charge present? To solve this, Thomson proposed the atom to be a sphere of positive charge with electrons distributed throughout it. This model was compared to a pudding with plums embedded in it, called the plum pudding model.

The Gold Foil Experiment

In 1911, Geiger and Marsden, working under Ernest Rutherford, tested Thomson’s model of the atom through what became famous as the gold foil experiment. They aimed a narrow beam of alpha particles at an extremely thin sheet of gold foil. Alpha (α) particles are tiny, positively charged particles emitted from certain radioactive elements.

According to Thomson’s model, the positive charge in the atom was spread out evenly. So they expected the alpha particles to pass straight through the gold foil or be deflected only slightly. But to their surprise, while most particles passed through undeflected, some were sharply deflected, and a few even bounced back. This deflection from the straight path is called scattering. Hence, the gold foil experiment is also called an α-ray scattering experiment.

Thomson’s model failed to explain the results of the gold foil experiment, particularly the deflection of some α-particles through large angles and that most of the α-particles passed undeflected.

Rutherford's Model of an Atom

From the gold foil experiment, Rutherford concluded that the positive charge of an atom is not spread throughout but remains concentrated in an extremely small region called the nucleus. He proposed that:

Most of an atom is empty space, as most α-particles passed through the gold foil without any deflection.
The nucleus is dense, contains all the positive charge and most of the mass of an atom.
The electrons revolve around the nucleus, somewhat like planets orbiting the Sun. Hence, this model is called the planetary model of the atom.

Limitations of Rutherford’s Model

Rutherford’s atomic model was better than Thomson's atomic model in explaining the results of the gold foil experiment. However, it had a limitation. It could not explain the stability of the atom.

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, while a sodium atom has 11 protons and 11 electrons. Since the total positive charge equals the total negative charge, the helium or the sodium atoms are also electrically neutral. Similarly, all the atoms are electrically neutral.

Bohr's Model of the Atom

To explain why atoms are stable, Niels Bohr proposed a new model of the atom in 1913. According to Bohr:

Electrons do not move randomly around the nucleus but 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.
These shells are represented by the letters K, L, M, N, ... or by the numbers n = 1, 2, 3, 4, ...
Electrons can revolve only in these allowed shells, and not in between them. While moving in a fixed shell, an electron does not lose energy.
The first energy level K (n = 1) is the one closest to the nucleus and has the least energy.
The energy of these levels increases as we move away from the nucleus. That is, the energy of an electron in the L-shell (n = 2) is more than that of an electron in the K-shell (n = 1). The farther away a shell is from the nucleus, the higher is its energy.
An electron can move to another shell by absorbing or releasing a fixed amount of energy equal to the difference between the energies of the two levels.
Each shell can hold only a certain number of electrons.

Discovery of the Neutron

Rutherford’s model showed that most of the mass of an atom is concentrated in its nucleus. Electrons that revolve around the nucleus are so light that their mass can be ignored.

However, something puzzling appeared early in the 20th century. For instance, a hydrogen atom has one proton, whereas a helium atom has two protons, yet the mass of a helium atom is about four times that of a hydrogen atom, not double. This led scientists to wonder whether besides protons, is there something else in the nucleus adding mass without affecting its charge?

In 1932, this problem was solved by James Chadwick (a student of Ernest Rutherford). He discovered a new subatomic particle with a mass nearly equal to that of a proton but no electrical charge. This neutral particle was named as neutron.

The neutrons are found in the nucleus of all atoms except hydrogen. Thus, the mass of an atom comes mainly from its protons and neutrons packed tightly in the nucleus. This also explains why atoms are heavier than the mass of their total number of protons.

Atomic Number (Z)

The atoms of an element are all alike but different from the atoms of other elements. They differ in the number of electrons and protons in them.

The number of protons in the nucleus of an atom of an element is known as its atomic number. It is designated by the symbol Z, this number determines the identity of an element and its chemical behaviour.

Since the atom as a whole is neutral, the number of protons in it equals the number of electrons orbiting the nucleus. For example, hydrogen has one proton and one electron, so its atomic number is 1. Helium, with atomic number 2, has 2 protons and 2 electrons.

Mass Number (A)

The total number of protons and neutrons present in the nucleus of an atom is called its mass number, and is denoted by A. The protons and neutrons present in the nucleus are called nucleons.

Mass number = Number of protons + Number of neutrons

Since, the mass of a neutron is roughly equal to that of a proton, the mass of a helium atom can be accounted by two protons and two neutrons.

$$ {}^{A}_{Z}\mathrm{X} $$

How are Electrons Distributed in Different Energy Levels?

The maximum number of electrons present in a shell is given by the formula 2n², where n is the number of the shell.

Hence, K-shell (n = 1) 2 × 1² = 2, can accommodate 2 electrons, L-shell (n = 2) 2 × 2² = 8 electrons, and M-shell (n = 3) 2 × 3² = 18 electrons.

The maximum number of electrons that can be accommodated in the outermost shell is 8 (the first shell can accommodate a maximum of two electrons).

Electrons are filled in these shells in a step-wise manner, starting from the one closest to the nucleus and moving outward, i.e., in the order K, L, M, N, ... The L-shell will be filled only after the K-shell is complete, and so on.

The distribution of electrons among various shells is known as the electronic configuration of the atom.

Valency

Atoms of the same or different elements can combine to form molecules. The number of atoms of hydrogen or chlorine with which one atom of an element can combine to form a compound is called its combining capacity.

The outermost shell containing electrons of an atom is known as its valence shell. The electrons present in it are known as valence electrons.

If the outermost shell of an atom has 8 electrons, it is called an octet. Elements with complete octet of electrons (8 electrons), or 2 electrons in the case of helium in their valence shell are largely unreactive and more stable. On the other hand, atoms with incomplete valence shells are usually more reactive. Such elements lose, gain, or share electrons to complete their octet.

The number of electrons gained, lost, or shared to complete the octet is called the valency of the element.

Generally, if the element has fewer than four electrons in its valence shell, it tends to lose electrons to complete its octet and become stable. On the other hand, if the number of valence electrons is more than four, it tends to gain electrons to complete its octet. For example, the electronic configuration of sodium is 2, 8, 1. It can get an octet by losing one electron. Therefore, its valency or combining capacity is 1. On the other hand, oxygen has an electronic configuration of 2, 6. It has six valence electrons. Therefore, it can gain two electrons to attain an octet, its valency is 2. The electronic configuration of carbon is 2, 4. Carbon has four valence electrons and cannot easily gain or lose them. Therefore, it can share four electrons with other atoms to complete its octet. Thus, the valency of carbon is 4.

Isotopes

Dalton proposed that all atoms of an element are identical and have the same mass. But scientists later discovered that there were atoms of the same element that could have the same number of protons (atomic number, Z) yet could have different numbers of neutrons, and thus, different mass numbers.

These atoms with the same atomic number but different mass numbers are called isotopes.

For example, naturally occurring hydrogen is a mixture of three different isotopes: protium, deuterium, and tritium. All of these contain one proton each, whereas deuterium contains one neutron, and tritium contains two neutrons.

Isobars

When atoms of different elements have the same mass number, but different atomic numbers, they are called isobars.

For example, calcium (atomic number 20), potassium (atomic number 19), and argon (atomic number 18). The number of protons in these atoms is different, but the mass number is 40.

First 18 Elements

Element	Atomic Number	Mass Number	K L M N	Valency
Hydrogen (H)	1	1	1 - - -	1
Helium (He)	2	4	2 - - -	0
Lithium (Li)	3	7	2 1 - -	1
Beryllium (Be)	4	9	2 2 - -	2
Boron (B)	5	11	2 3 - -	3
Carbon (C)	6	12	2 4 - -	4
Nitrogen (N)	7	14	2 5 - -	3
Oxygen (O)	8	16	2 6 - -	2
Fluorine (F)	9	19	2 7 - -	1
Neon (Ne)	10	20	2 8 - -	0
Sodium (Na)	11	23	2 8 1 -	1
Magnesium (Mg)	12	24	2 8 2 -	2
Aluminum (Al)	13	27	2 8 3 -	3
Silicon (Si)	14	28	2 8 4 -	4
Phosphorus (P)	15	31	2 8 5 -	3 or 5
Sulfur (S)	16	32	2 8 6 -	2
Chlorine (Cl)	17	35.5	2 8 7 -	1
Argon (Ar)	18	40	2 8 8 -	0