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CBSE NOTES CLASS 9 SCIENCE CHATER 4

STRUCTURE OF ATOM

Sub-Atomic Particles

Electron

Proton

Neutron

Atomic Models

Thomson’s Model

Rutherford’s Model

The (α)-particle experiment

Bohr’s Atomic Model

Arrangement of electrons in an atom or electronic structure

Valency

Atomic number (Z)

Mass number (A)

Isotopes

Isobars

Average atomic mass of an atom with isotopes

CBSE NOTES CLASS 9 SCIENCE CHATER 4

STRUCTURE OF ATOM

1. Sub-Atomic Particles

Electron was discovered in cathode ray experiment. It is negatively charged and is represented as e-. It was discovered by JJ Thomson in 1869.

Proton was discovered in anode ray experiment. Anode rays are also called positive rays or canal rays. Proton is positively charged. Proton was discovered by Goldstien in 1886.

Neutron was discovered by James Chadwick in 1932. It is neutral (has no charge) and is represented by n. The mass of neutron and proton are almost equal.

2. Atomic Models

Atomic models are structures showing arrangement of different sub atomic particles within the atom.

Various Models of atom

(I) Thomson’s Model or Water Melon Model or Plum Pudding Model

This model was proposed by Joseph James Thomson in 1897. This model is also known as apple pie model.

Thomson proposed that:

(i) An atom is a positively charged sphere and the electrons are embedded in it.

(ii) The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral.

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Merits of Thomson’s Model - Thomson’s model explained that atoms are electrically neutral.

Shortcomings - It could not explain the results of α (alpha) particle scattering experiment carried out by Rutherford.

(II) Rutherford’s Model or Planetary Model

This model was given by Ernest Rutherford in 1911, based on his α particle scattering experiment

The alpha (α)-particle experiment

In this experiment, fast moving alpha (α)-particles were made to strike a thin gold foil.

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Observations

(i) Most of the fast moving α-particles passed straight through the gold foil.

(ii) Some of the α-particles were deflected by the foil by small angles and a very small number by large angles.

(iii) One out of every 12000 particles appeared to rebound

Conclusions

(i) Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected.

(ii) Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.

(iii) A very small fraction of α-particles were deflected by 1800, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.

Rutherford’s Model

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(i) All the positively charged particles are present in a small space in the centre of the atom. This small space is called nucleus.

(ii) Electrons (negatively charged) revolve around the nucleus in circular orbits with a high speed.

(iii) The size of the nucleus is very small compared to the size of atom.

Drawbacks of Rutherford’s Model

(i) It is possible to have infinite number of orbits. In practice it is not the case.

(ii) The electron moving with acceleration must continuously lose energy and eventually will fall into the nucleus, which is contrary to the fact.

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(III) Bohr’s Atomic Model: This model was given by Neils Bohr in 1913.

According to this model,

(i) Only certain special orbits known as discrete orbits (or energy levels) of electrons, are allowed inside the atom.

(ii) While revolving in discrete orbits the electrons do not radiate energy.

The first orbit (i.e., for n = 1) is denoted by letter K. Similarly, for n = 2, it is L-shell, for n = 3, it is M-shell and for n = 4, it is N-shell. These orbits or shells are also called energy levels.

3. Arrangement of electrons in an atom

The arrangement of electrons in various shells (energy levels) of an atom of the element is known as electronic configuration.

The Maximum number of electrons that could be put in a particular shell (i.e., energy levels) was given by Bohr and Bury.

According to Bohr-Bury Scheme

4. Valency

5. Atomic number (Z)

Atomic number of an element is equal to the number of protons present in the nucleus of an atom. Atomic number (Z) = number of protons = number of electrons.

6. Mass number (A)

It refers to the total number of neutrons and protons present in an atom.

Mass number (A) = number of protons + number of neutron

An element is represent as

Symbol ofthe elementAtomic NumberMass Number.

For example, nitrogen is written as N7 14.

7. Isotopes

Atoms of the same element having same atomic number but different mass numbers are known as Isotopes eg.

(i) C612, C613 and C614

(ii) H11 (Protium), H12 (Deutrium), H13 (Tritrium)

(iii) Cl1735 and Cl1737

Applications of Isotopes

8. Isobars

Atoms of different elements having same mass number are known as Isobars, e.g Ca2040 and Ar1840

9. Average atomic mass of an atom with isotopes

Average Atomic Mass= 

AM of Isoptope1 ×percentage+AM of Isoptope2×percentage100 

Q: Find Average mass of chlorine if abundance of Cl-35 in nature is 75% and that of Cl-37 is 25%