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Structure of Atoms - Class 9 Science Notes

Introduction


Matter is made up of fundamental building blocks, that is atoms and molecules. The existence of different kinds of matter is due to different atoms constituting them.


Are atoms really indivisible, as proposed by Dalton?


One of the first indication that atoms are not indivisible (means divisible), comes from studying static electricity and the conditions under which electricity is conducted by different substances.


Charged Particles in Matter

From different experiments, it was discovered that an atom is divisible and consists of charged particles.


By 1900, J. J. Thomson discovered a sub-atomic particle called an electron.


Even before the electron was identified, E. Goldstein in 1886 discovered the presence of new radiations in a gas discharge and called them canal rays. These rays were positively charged radiations which ultimately led to discovery of another sub-atomic particle called proton.


The electron is represented as ‘e’ while the proton as ‘p’.

The mass of proton is taken as one unit and its charge as plus one, whereas the mass of an electron is considered to be negligible and its charge is minus one.



Neutrons


In 1932, J. Chadwick discovered another sub-atomic particle which had no charge and a mass nearly equal to that of a proton. It was named as Neutrons. In general neutron is represented as ‘n’.


Neutrons are present in the nucleus of all the atoms, expect hydrogen.


The mass of an atom is given by the sum of the masses of protons and neutrons present in the nucleus.



Particle

Electron

Proton

Neutron

Symbol

e

p

n

Relative Charge

-1

+1

0

Nature

Negatively Charged

Positively Charged

Neutral

Discovery

JJ Thomson

E. Goldstein

Chadwick



The Structure of an Atom


Dalton’s atomic theory suggested that the atom is indivisible and indestructible. But the discovery of the two fundamental particles (electrons and protons) inside the atom, led to the failure of this aspect of Dalton’s atomic theory.


So, it became necessary to know how the electrons and protons are arranged within an atom. For explaining this, many model for the structure of atom were proposed.



(A) Thomson’s Model of an Atom


  • Thomson proposed the model of an atom to be similar to that of a Plum pudding (Christmas pudding) or Watermelon.


Thomson model is compared with watermelon as shown in the figure:

  • The positive charge in the atom is spread all over like the red edible part of the watermelon.

  • While the electrons are studded in the positively charged sphere, like the seeds in the watermelon.

Postulates of Thomson’ model of an Atom:

  1. An atom consists of a positively charged sphere and the electrons are embedded in it.

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


Although, Thomson’s model explained that atoms are electrically neutral, but it failed to explain the stability of an atom. This theory also fails to accounts for the position of neutrons in the atom.


There are no experimental evidences in its supports.



(B) Rutherford’s Model of an Atom


Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. For this Rutherford designed an experiment, in which fast moving alpha – 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.

  • Alpha–particles are doubly charged helium ions. Since they have a mass of 4 u, the fast moving alpha – particles have a considerable amount of energy.

  • It was expected that alpha – particles would be deflected by the sub-atomic particles in the gold atoms. Since the alpha – particles were much heavier than the protons, he did not expect to see large deflections.

But, the alpha – particles scattering experiments gave totally unexpected results. The following were the observations:

(a) Most of the fast moving alpha – particles passed straight through the gold foil.

(b) Some of the alpha – particles were deflected by the foil by small angles.

(c) Surprisingly one out of every 12000 particles appeared to rebound.








Conclusion from alpha – particles scattering experiment are:


  1. Most of the space inside the atom is empty because most of the alpha – particles passed through the gold foil without getting deflected.

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

  3. A very small fraction of alpha – particles were deflected by 180o, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.


From the data, it is was concluded that the radius of the nucleus is 105 times less than the radius of the atom.



On the basis of experiment, Rutherford put forward the nuclear model of an atom, which had the following features:

  1. There is positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus.

  2. The electrons revolve around the nucleus in circular paths.

  3. The size of the nucleus is very small as compared to the size of the atoms.


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 radiating energy and would lose energy and finally fall into the nucleus. If this were so, the atom should be highly unstable and hence the matter would not exist in the form that we know.

  • Fails to explain the arrangement of the electrons in an atom.


(C) Bohr’s Model of an Atom


In order to overcome the objective raised against Rutherford’s model of an atom, Neil Bohr put forward the following postulates of the model of an atom:

  1. Only certain special orbits known as discrete orbit of electrons are allowed inside the atom.

  2. While revolving in discrete orbits the electrons do not radiate energy.


These orbits or shells are called energy levels. These orbits or shells are represented by the letters K, L, M, N, … or the numbers, n = 1, 2, 3, 4, …


Summary of three different models of an Atom:



How are electrons distributed in different orbits (shells)?


The distribution of electrons into different orbits was suggested by Bohr and Bury. The following rules are followed for writing the number of electrons in different energy levels or shells:


  • The maximum number of electrons present in a shell is given by the formula

2n2, where n = orbit number or energy level index,

n = 1, 2, 3, ……

Hence maximum number of electrons in different shells are as follows:

First orbit or K-shell = 2 × 12 = 2

Second orbit or L-shell = 2 × 22 = 8

Third orbit or M-shell = 2 × 23 = 18

Fourth orbit or N-shell = 2 × 24 = 32 and so on …



  • The maximum number of electrons that can be accommodated in the outermost orbit is 8.


  • Electrons are not accommodated in a given shell, unless the inner shells are filled. That is, the shells are filled in a step-wise manner.


Atomic structure of the first eighteen elements:



Valency


  • Valency is the combining capacity of an atom.

  • The electrons present in the outermost shell of an electrons are known as the valence electrons.

  • The outermost shell of an atom can be accommodated with a maximum of 8 electrons; hence the shell is completely filled and show little chemical activity. In other words, their combining capacity is zero.

  • Of these inert elements, the helium atom has 2 electrons in its outermost shell and all other elements (neon, argon, krypton, xenon and radon) have atoms with 8 electrons in the outermost shell.

  • The combining capacity of the atoms of elements, ie., their tendency to react and form molecules with atoms of the same or different elements, can be explained based on their attempt to attain a fully-filled outermost shell



  • An outermost shell, which had eight electrons was said to possess an octet. This is done by sharing, gaining or losing electrons.

  • The number of electrons shared, gained or lost so as to make the octet of electrons in the outermost shell, gives the combining capacity of the element, i.e., valency.

  • For example, hydrogen/ lithium/ sodium atoms contain one electron each in their outermost shell, therefore each one of them can lose one electron. So, they have valence of one.

  • If the number of electrons in the outermost shell of an atom is close to its full capacity, then valency is determined in a different way.

  • For example, the fluorine atom has 7 electrons in the outermost shell and so its valency could be 7. But for fluorine to gain one electron is easier instead of losing 7 electrons. Hence, its valency is determined by subtracting seven electrons from the octet and thus its valency would be one.



Atomic Number and Mass Number


Atomic Number:


Atomic number is defined as the total number of protons present in the nucleus of an atom. It is denoted by ‘Z’.


All atoms of an element have the same atomic number, Z. in fact the elements are defined by the number of protons they possess.


For example, for Hydrogen, Z = 1, because it possesses only one electron.

Similarly, for Carbon, Z = 6, as it possesses six protons.



Mass Number:

The mass number is defined as the sum of total number of protons and neutrons present in the nucleus of an atom. It is denoted by ‘A’.



The protons and neutrons together are called nucleons.


For example, mass of carbon is 12 u, i.e., 6 protons + 6 neutrons, 6 u + 6 u = 12 u

Similarly, mass of aluminium is 27 u, i.e., 13 protons + 14 neutrons.


In the notation for an atom, the atomic number, mass number and symbol of the element are to be written as:

For example, nitrogen is written as





Isotopes:


Isotopes are defined as the atoms of the same element, having the same atomic number but different mass numbers.


The chemical properties of isotopes are similar but their physical properties are different.



Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance.

For example, Chlorine occurs in two isotopic forms in nature, with masses 35 u and 37 u in the ratio of 3:1.



The average mass is taken of all the naturally occurring atoms of that element. If an element has no isotope, then the mass of its atom is the sum of protons and neutrons in it. But if an element occurs in isotopic form, then percentage of each isotopic from is to be known and then average mass is calculated.



So, for certain amount of chlorine taken it will contain both the isotopes of chlorine and the average mass is 35.5 u.



Application:

(i) An isotope of Uranium is used as a fuel in nuclear reactors.

(ii) An isotope of cobalt (Co60) is used in the treatment of cancer.

(iii) An isotope of iodine is used in the treatment of goitre.



Isobars:


Atoms of different elements with different atomic numbers, which have the same mass number, are known as isobars. Or


The total number of nucleons (protons + neutrons) is the same in the atoms.





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