Chemistry class 9 Structure of the atom

Class 9 chemistry chapter 4 - structure of the Atom 

       NOTESWALA 2020

Structure of the Atom

Atoms

•the building blocks of matter.
•the smallest unit of matter. 

Cathode ray experiment

•J. J. Thomson discovered the existence of electrons.
•He did this using a cathode ray tube, which is a vacuum-sealed tube with a cathode and anode on one end that created a beam of electrons travelling towards the other end of the tube.

•The air inside the chamber is subjected to high voltage and electricity flows through the air from the negative electrode to the positive electrode.

•The characteristics of cathode rays (electrons) do not depend upon the material of electrodes and the nature of the gas present in the cathode ray tube.

•The experiment showed that the atom was not a simple, indivisible particle and contained at least one subatomic particle – the electron.

Electrons

•Electrons are the negatively charged sub-atomic particles of an atom.

•The mass of an electron is considered to be negligible, and its charge is -1.

•The symbol for an electron is e^–

•Electrons are extremely small.

•They are found outside the nucleus.

Thomson’s model of an atom

•According to Thomson,(i) An atom consists of 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

•The first model of an atom to be put forward and taken into consideration.

•He proposed a model of the atom be similar to that of a Christmas pudding/watermelon.

•The red edible part of the watermelon is compared with the positive charge in the atom.

•The black seeds in the watermelon are compared with the electrons which are embedded on it.

Rutherford’s experiment and observations

• In this experiment, fast-moving alpha (α)-particles were made to fall on a thin gold foil. His observations were:

• A major fraction of the α-particles bombarded towards the gold sheet passed through it without any deflection, and hence most of the space in an atom is empty.

• Some of the α-particles were deflected by the gold sheet by very small angles, and hence the positive charge in an atom is not uniformly distributed.

• The positive charge in an atom is concentrated in a very small volume.

• Very few of the α-particles were deflected back, that is only a few α-particles had nearly 180o angle of deflection. So the volume occupied by the positively charged particles in an atom is very small as compared to the total volume of an atom.

Rutherford’s model of an atom

Rutherford concluded the model of the atom from the α-particle scattering experiment as:

(i) There is a positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus.

(ii) The electrons revolve around the nucleus in well-defined orbits.

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

Drawbacks of Rutherford’s model

• He explained that the electrons in an atom revolve around the nucleus in well-defined orbits. Particles in a circular orbit would experience acceleration.

• Thus, the revolving electron would lose energy and finally fall into the nucleus.

• But this cannot take place as the atom would be unstable and matter would not exist in the form we know.

Bohr’s Model of an atom

• Bohr came up with these postulates to overcome the objections raised against Rutherford’s model:

• Electrons revolve around the nucleus in stable orbits without emission of radiant energy. Each orbit has a definite energy and is called an energy shell or energy level.

• An orbit or energy level is designated as K, L, M, N shells. When the electron is in the lowest energy level, it is said to be in the ground state.

• An electron emits or absorbs energy when it jumps from one orbit or energy level to another.

• When it jumps from a higher energy level to lower energy level, it emits energy while it absorbs energy when it jumps from lower energy level to higher energy level.

Atomic Number

The number of protons present in the nucleus of an atom is called atomic number. It is denoted by Z.

Atomic number of some common elements

For atoms:

Atomic number = number of proton = number of electron

For ions:

Atomic number = number of proton ≠ number of electron

For example, in aluminium atom number of electrons is equal to atomic number but in aluminium ion it is not so.

Mass number

Mass number is equal to the number of nucleons present inside the nucleus of an atom. It means it is the sum of number of protons and neutrons present in the nucleus of an atom. It is denoted by letter A.

Mass number of element = Atomic mass of element = number of protons + number of neutrons

For example -

Carbon   : Mass number  =  12  (no. of p⁺ = 6, no. of  n = 6 )
Nitrogen : Mass number  =  14  (no. of p⁺ = 7, no. of  n  = 7 )
Fluorine  : Mass number  =  19  (no. of p⁺ = 9, no. of  n = 10 )

Isotopes

Isotopes are atoms of same element having same number of protons but different number of neutrons. Isotopes have similar chemical properties but different physical properties.

For example -

Uses of Isotopes:

1. An Isotope of uranium (23592Ur) is used in nuclear power plants to generate electricity.

2. Used for medical purposes:

An Isotope of cobalt is used in the treatment of cancer.

An isotope of iodine is used in the treatment of goiter.

Isobars

The atoms of several elements can have a similar mass number but distinct atomic masses. Such elements are called Isobars.