The document discusses the structure of atoms, defining them as the basic building blocks of matter made up of protons, neutrons, and electrons, with historical insights into their discovery. It explores various atomic models proposed by scientists like Dalton, Thomson, Rutherford, and Bohr, highlighting key features and limitations of each. Additionally, it explains concepts like electronic configuration, valency, isotopes, and isobars, emphasizing the applications of isotopes in fields such as nuclear energy and medicine.

1. STRUCTURE OF THE
CLASS – 9

2. WHAT IS AN ATOM ?
Atoms are defined as “Basic Building Blocks of Matter”. It is the Smallest
Constituent Unit of Matter that posses the properties of the chemical element.
Atoms are very small. Atomic Radius is measured in Nanometers(1 m =109
nm).
Atom is made up of three Sub – Atomic Particles :-
 Protons – It is found in the Nucleus with a Positive Charge and mass of 1 A.M.U.
 Neutrons – It is found in the Nucleus with a Neutral Charge and mass of 1
A.M.U.
 Electrons – It is a Sub – Atomic Particle with Electrically Negative Charge and of
Negligible Mass ( Mass of the Proton).

3. The Sub – Atomic
Particles

4. PROTONS
In 1886 E. Goldstein discovered evidence for
the existence of this Positively Charged
Particle (Protons) . He used a Cathode Ray
Tube with Holes. Then he noticed that these
rays were travelling in the Opposite Direction
from the Cathode Rays. He called them as
the Canal Rays and showed that they were
composed of Positively Charged Particles.
The mass of the Proton is approximately
2000 times as that of the Electron. In general,
Proton is represented as ‘p+
’. Its mass is taken
as 1 Atomic Mass Unit and the Charge of +1.

5. NEUTRONS
In 1932, J. Chadwick discovered the Sub –
Atomic Particle which had Neutral Charge
and a Mass of 1 A.M.U. which is equal to
the mass of the Protons. It was named as
Neutron. Neutrons are present in the
nucleus of all atoms, except Hydrogen. In
general, a Neutron is represented as ‘n’.
(Atomic Mass = No. of Protons + No. of
Neutrons)

6. ELECTRONS
In 1897, J.J. Thompson discovered
Electrons which had a Negative Charge
and mass which is than that of the
Protons. In general, Electron is
represented as ‘e-
’.
(Atomic Number = No. of Protons = No. of
Electrons)

7. DALTON’S ATOMIC THEORY
The Atomic Theory of John Dalton suggested that the
atom was Indivisible and Indestructible. But the
discovery of Sub – Atomic Particles led to the failure of
Dalton’s aspect. Then, it was considered necessary to
know how electrons and protons are arranged within
an atom. For explaining this, many scientists proposed
various atomic models. J.J. Thomson was the first one
to propose a model for the structure of an atom. Then,
Ernest Rutherford proposed a model for the structure
of an atom by doing Certain Experiments. After that,
Neils Bohr postulated a model of an atom which
overcome the Objections raised by Rutherford.

8. J.J. THOMSON’S MODEL
J.J. Thomson proposed the model of an atom to be
similar to that of a Christmas pudding. The
Electrons, in a sphere of Positive Charge, were like
Currants (dry fruits) in a Spherical Christmas
pudding. It was just like a Watermelon where the
Red Part was the Positive Sphere and the Seeds are
the Electrons which are embedded in it.
Thomson proposed that :-
 An atom consists of a positively charged sphere
and the electrons are embedded in it.
 The Negative and Positive Charges are equal in
Magnitude. So, the atom as a whole is Electrically
Neutral.

9. RUTHERFORD’S MODEL
Ernest Rutherford was interested in knowing how the electrons are arranged
within an atom. So, he carried out an experiment in which fast moving Alpha
(α)-Particles were made to fall on a thin Gold Foil.
 He selected a Gold Foil because he wanted as thin a layer as possible. This
gold foil was about 1000 Atoms Thick.
 α-Particles are Doubly-Charged Helium Ions. As, they have a mass of 4 u, the
fast-moving α-particles have a considerable amount of energy.
 It was expected that α-particles would be deflected by the sub-atomic
particles in the gold atoms. Since the α-particles were much heavier than the
protons, he did not expect to see large deflections.

10. RUTHERFORD’S OBSERVATIONS
From the – particles Experiment Rutherford concluded
that :-
 Most of the fast moving – particles passed straight
through the Gold Foil because most of the space
inside the atom is Empty.
 Some of the α-particles were deflected by the foil by
Small Angles, indicating that the Positive Charge of
the atom occupies Very Little Space.
 Surprisingly one out of every 12000 Particles
appeared to Rebound by 180º, indicating that all the
Positive Charge and Mass of the Gold Atom were
concentrated in a Very Small Volume within the atom.

11. DRAWBACKS OF RUTHERFORD’S MODEL
On the basis of his Experiment, Rutherford put forward the Nuclear Model of An
Atom, which had the following features:
 There is a Positively Charged Center in an atom called the Nucleus. Nearly all
the Mass of An Atom resides in the Nucleus.
 The Electrons revolve around the Nucleus in Circular Paths.
 The Size of the Nucleus is very small as compared to the Size of the Atom.
Drawbacks of Rutherford’s Model of the Atom
The Revolution of the Electron in a Circular Orbit is not expected to be Stable.
Any particle in a Circular Orbit would Undergo Acceleration. During
Acceleration, charged particles would Radiate Energy. Thus, the Revolving
Electron would Loses Energy and finally fall into the Nucleus. If this were so, the
atom should be Highly Unstable. But atoms are Quite Stable.
On the basis of his Experiment, Rutherford put forward the Nuclear Model of An
Atom, which had the following features:
 There is a Positively Charged Center in an atom called the Nucleus. Nearly all
the Mass of An Atom resides in the Nucleus.
 The Electrons revolve around the Nucleus in Circular Paths.
 The Size of the Nucleus is very small as compared to the Size of the Atom.
Drawbacks of Rutherford’s Model of the Atom
The Revolution of the Electron in a Circular Orbit is not expected to be Stable.
Any particle in a Circular Orbit would Undergo Acceleration. During
Acceleration, charged particles would Radiate Energy. Thus, the Revolving
Electron would Loses Energy and finally fall into the Nucleus. If this were so, the
atom should be Highly Unstable. But atoms are Quite Stable.

12. BOHR’S MODEL OF ATOM
In order to overcome the objections raised
against Rutherford’s Model of the Atom, Neils
Bohr put forward the following postulates about
the model of an atom:
 Only Certain Special Orbits known as Discrete
Orbits of Electrons, are allowed inside the atom.
 While revolving in Discrete Orbits the Electrons
Don’t Radiate Energy.
 These orbits or shells are also known as Energy
Levels.
 These orbits or shells are represented by the
Letters K,L,M,N or by the Numbers, n=1,2,3,4.

13. ELECTRONIC CONFIGURATION
The Electronic Configuration was suggested by Neils Bohr and Charles Rugeley Bury.
The following rules are followed for writing the number of electrons in different Discrete
Orbits:
 The maximum number of electrons present in a shell is given by the formula , where ‘n’ is
the orbit number with 1, 2, 3, 4 increases according to the:
• 1st
Orbit (K – Shell) : 2 * 12
= 2 Electrons
• 2nd
Orbit (L – Shell) : 2 * 2
2
= 8 Electrons
• 3rd
Orbit (M – Shell) : 2 * 32
= 18 Electrons
• 4th
Orbit (N – Shell) : 2 * 42
= 32 Electrons
 The maximum number of electrons that can be accommodated in the Outermost Orbit is
8.
 Electrons are not filled in the next shell unless previous shells are filled.

14. VALENCY
The combining capacity of the atoms of other elements, that is, their tendency
to react and form molecules with atoms of the same or different elements is
called Valency.
Valency is determined by the outermost shell, which has to posses an Octet(8
Electrons). This is done by sharing, gaining or losing electrons. The number of
electrons gained, lost or shared so as to make the octet of electrons in the
outermost shell, gives us directly the Valency.
• If the element is a Non – Metal having 5,6,7- Electrons in the Outermost shell
then the valency is determined by subtracting No. of Electrons in Outermost
Shell from Octet.
• If the element is a Metal having 1,2,3 – Electrons in the Outermost shell then
the valency is said to be the No. of electrons in the Outermost Shell.

15. Composition of Atoms of the First Eighteen Elements
with Electron Distribution in Various Shells

16. ISOTOPES
Isotopes are atoms of the same element that have the same number of protons
(i.e., atomic number, "Z") but a different number of neutrons, meaning that
their mass number, "A", varies. For example, take the case of hydrogen atom, it
has three atomic species, Protium (1
1 H), Deuterium (2
1 H), Tritium (3
1 H).
Applications
Since the chemical properties of all the isotopes of an element are the same,
normally we are not concerned about taking a mixture. But some isotopes have
special properties which find them useful in various fields. Some of them are :
(i) An isotope of uranium is used as a fuel in nuclear reactors.
(ii) An isotope of cobalt is used in the treatment of cancer.
(iii) An isotope of iodine is used in the treatment of goitre.

17. ISOBARS
Let us consider two elements — calcium, atomic number 20, and argon,
atomic number 18. The number of electrons in these atoms is different, but
the mass number of both these elements is 40. That is, the total number of
nucleons is the same in the atoms of this pair of elements. Atoms of different
elements with different atomic numbers, which have the same mass number,
are known as Isobars.