Chapter 4 Structure of the Atom Science Notes Class 9 

Table of Contents-

1. Introduction

2. Charged particles in matter ( Subatomic particles )

    2.1 Electron

    2.2 Proton

    2.3 Neutron

3. The structure of the atom

    3.1 Thomson's Model of an Atom

    3.2 Rutherford's Model of an Atom

          3.21 Drawbacks of Rutherford's model of the atom

    3.3 Bohr's Model of an Atom

4. Arrangement of electrons in the atoms

5. Valence electrons

6. Valency

    6.1 Types of Valency

7. Atomic number and Mass number

8. Isotopes and Isobars and their applications

Introduction- 

As we know. that atoms and molecules are the building blocks of matter and everything around us is made up of atoms and molecules.

So, the question arises of what actually an atom is? or what it looks like?

A significant challenge faced by scientists at the end of the 19th century was uncovering the structure of the atom.

Charged particles in matter ( Subatomic particles )

The nature of charged particles in matter can be explained through simple activity.

Activity- Comb your dry hair, then put this comb towards the small pieces of paper and see. Did you see the comb attract small pieces of paper?

Electron-

• J.J. Thomson discovered the electron in 1897.
• He took a discharge tube which is a long glass tube sealed at both ends and fitted with two metal electrodes. The electricity is passed through a gas present in a discharge tube.
• Streams of minute particles were ejected out by the cathode. These streams of particles were called cathode rays.
• These cathode rays are small negatively charged particles called electrons.

Protons-

• E. Goldstein discovered the proton in 1886.
• The proton is a positively charged particle.

Neutron-

• James Chadwick in 1932 discovered the neutron.
• The neutron is a neutral particle found in the nucleus of an atom present in the center.

The structure of the atom-

Dalton's atomic theory suggested that the atom was not divisible. But, now the discovery of two fundamental particles ( electrons and protons) inside the atom, led to the failure of this aspect of Dalton's atomic theory.

Thomson's Model of an Atom-

• Thomson proposed the model of an atom to be similar to that of Christmas pudding.
• The spherical portion of Christmas pudding is equivalent to the positive charge and the electrons embedded in it are equivalent to the negative charge.
• Like that of watermelon, the red edible part represents the sphere of positive charge whereas the black seeds embedded in the watermelon are like the elctrons.

Rutherford's Model of an Atom-

• In his experiment, when the fast-moving alpha particles were allowed to strike a very thin gold foil in a vacuum it was found that,

1. Most of the alpha particles pass straight through a gold foil.
2. A few of the alpha particles were deflected by the foil at small angles.
3. A very few, one out of every 1200 particles rebound completely on hitting the gold foil and turn back on their path.

The following 3 conclusions are made by the Rutherford -

1. Most of the alpha particles pass through straight means that most of the space inside the atom is empty.
2. Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space and is present in the center.
3. Rarely, one particle was deflected by 180 degrees, and all the positive charge and the mass of the gold atom were concentrated in the center.

Drawbacks of Rutherford's model of an atom-

According to him, the electrons that revolve around the nucleus in the circular orbit would undergo acceleration and due to acceleration, charged particles would radiate energy and lose energy thereby falling into the nucleus. So, the atom should be highly unstable and should not exist at all.

Bohr's Model of Atom-

Bohr put forward the following postulates about Bohr's model of an atom-

• An atom consists of  3 particles electrons, protons, and neutrons.
• The protons and neutrons are located in the center of the nucleus of an atom.
• The electrons revolve rapidly around the nucleus in a fixed orbit ( known as discrete orbits) called energy levels or shells.
• These orbits or shells are represented by the letters K, L, M, N......... and so on or the numbers n= 1,2,3,4.......
• There is no change in the energy of electrons as long as they keep revolving at the same energy level, and the atom remains stable.

Arrangement of electrons in different orbits-

The following rules are followed when writing the number of electrons in different energy levels or orbits.

• The maximum number of electrons that can occupy a given shell is given by the formula 2n square where 'n' is the orbit number or energy level.
• For example,   For K shell ( 1st orbit ) = 2(1) square = 2

                                   For L shell ( 2 nd orbit ) = 2(2) square =4

• Electrons are filled in the increasing order of the inner shells i.e. first inner shell is filled up then the outer shell.

Valence electrons-

Those electrons that are present in the outermost shell of an atom are known as valence electrons.

Valency- 

The combining capacity of an atom of an element to form a chemical bond is known as its valency.

For example- The valency of carbon is 4 and that of hydrogen is 1. So, one carbon atom combines with 4 hydrogen atoms to form a methane molecule. ( CH4).

Types of Valency-

1. Electrovalency- The number of electrons lost or gained by one atom of an element to achieve the nearest gas electronic configurations known as its electrovalency.

For example, Sodium ( Na ) atomic number is 11 so its electronic configuration is 2, 8,1. The outermost valence electron after losing 1 electron achieves the nearest inert gas electron configuration.

2. Covalency- The number of electrons shared by one atom of an element to achieve the nearest inert gas electron configuration is known as its covalency.

For example, a Chlorine( Cl ) atom shares 1 electron to achieve the nearest gas electron arrangement therefore the covalency of chlorine is 1.

Atomic number and Mass number-

Atomic number- The number of protons in one atom of an element is known as the atomic number of that element.

                          Atomic number = Number of protons 

Mass number - The sum total of the number of protons and neutrons is called mass number.

                          Mass number = number of protons + number of neutrons

Isotopes and Isobars -

Isobars - Isotopes are atoms of the same element having the same atomic number but different mass numbers.

Isobars - Isobars are the atoms of different elements that have different atomic numbers but the same mass number.

Applications-

• Isotopes of uranium 235 atoms are used in nuclear power plants to generate electricity.
• Radioactive isotopes are used as a tracer in medicine to detect the presence of tumors and blood clots, in the human body.
• Iodine 131 radioactive isotopes are used in the treatment of goiter 

 

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