Structure of the Atom

Class 09 Science

Atoms and molecules are the fundamental building blocks of matter. The existence of different kinds of matter is due to different atoms constituting them.

Charged Particles in Matter

On rubbing two objects together, they become electrically charged. Many scientists contributed in revealing the presence of charged particles in an atom.

It was known by 1900 that the atom was indivisible particle but contained at least one sub-atomic particle - the electron identified by J.J. Thomson. 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 the discovery of another sub-atomic particle. This sub-atomic particle had a charge, equal in magnitude but opposite in sign to that of the electron. Its mass was approximately 2000 times as that of the electron. It was given the name of proton.

In general, an electron is represented as ‘e^–’ and a proton as ‘p⁺’. The mass of a proton is taken as one unit and its charge as plus one. The mass of an electron is considered to be negligible and its charge is minus one.

It seemed that an atom was composed of protons and electrons, mutually balancing their charges. It also appeared that the protons were in the interior of the atom, for whereas electrons could easily be removed off but not protons.

Thomson's Model of an Atom

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

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. 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.

Rutherford's Model of Atom

Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. Rutherford designed an experiment for this. In this experiment, 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. The α-particles are doubly-charged helium ions. Since 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. But, the α-particle scattering experiment gave totally unexpected results. The following observations were made:

Most of the fast moving α-particles passed straight through the gold foil.
Some of the α-particles were deflected by the foil by small angles.
Surprisingly one out of every 12000 particles appeared to rebound.

Rutherford concluded from the α-particle scattering experiment that:

Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected.
Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.
A very small fraction of α-particles were deflected by 180°, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.

He also calculated that the radius of the nucleus is about 10⁵ times less than the radius of the atom.

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

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 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, ....

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 eventually named as neutron. Neutrons are present in the nucleus of all atoms, except hydrogen. The mass of an atom is given by the sum of the masses of protons and neutrons present in the nucleus.

How are Electrons Distributed in Different Orbits (Shells)?

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

(i) The maximum number of electrons present in a shell is given by the formula 2n², where n is the orbit number or energy level. Hence the maximum number of electrons in different shells are:

K-shell = 2 × 1² = 2
L-shell = 2 × 2² = 8
M-shell = 2 × 3² = 18
N-shell = 2 × 4² = 32

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

(iii) Electrons are not accommodated in a given shell, unless the inner shells are filled. The shells are filled in a step-wise manner.

Valency

The electrons present in the outermost shell of an atom are known as the valence electrons. The outermost shell of an atom can accommodate a maximum of 8 electrons. It was observed that the atoms of elements, completely filled with 8 electrons in the outermost shell show little chemical activity. Their combining capacity or valency is zero.

An outermost-shell, which had eight electrons was said to possess an octet. Atoms would react to achieve an octet in the outermost shell. This was 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 the combining capacity of the element or valency.

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 its valency could be 7. But it is easier for fluorine to gain one electron instead of losing seven electrons. Hence, its valency is determined by subtracting seven electrons from the octet and this gives a valency of one for fluorine.

Atomic Number

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

For hydrogen, Z = 1, because in hydrogen atom, only one proton is present in the nucleus. Similarly, for carbon, Z = 6.

Mass Number

The mass number is defined as the sum of the total number of protons and neutrons present in the nucleus of an atom. It is denoted by A. Protons and neutrons are also called nucleons.

For example, mass of carbon is 12 u because it has 6 protons and 6 neutrons, 6 u + 6 u = 12 u. Similarly, the mass of aluminium is 27 u (13 protons + 14 neutrons).

$$ {}^{A}_{Z}\mathrm{X} $$

Isotopes

In nature, a number of atoms of some elements have been identified, which have the same atomic number but different mass numbers.

For example, hydrogen atom has three atomic species, namely protium, deuterium and tritium. The atomic number of each one is 1, but the mass number is 1, 2 and 3, respectively.

Isobars

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

For example, calcium (atomic number 20) and argon (atomic number 18) The number of protons in these atoms is different, but the mass number of both these elements is 40.

First 18 Elements

Element	Atomic Number	Mass Number	K L M N	Valency
Hydrogen (H)	1	1	1 - - -	1
Helium (He)	2	4	2 - - -	0
Lithium (Li)	3	7	2 1 - -	1
Beryllium (Be)	4	9	2 2 - -	2
Boron (B)	5	11	2 3 - -	3
Carbon (C)	6	12	2 4 - -	4
Nitrogen (N)	7	14	2 5 - -	3
Oxygen (O)	8	16	2 6 - -	2
Fluorine (F)	9	19	2 7 - -	1
Neon (Ne)	10	20	2 8 - -	0
Sodium (Na)	11	23	2 8 1 -	1
Magnesium (Mg)	12	24	2 8 2 -	2
Aluminum (Al)	13	27	2 8 3 -	3
Silicon (Si)	14	28	2 8 4 -	4
Phosphorus (P)	15	31	2 8 5 -	3 or 5
Sulfur (S)	16	32	2 8 6 -	2
Chlorine (Cl)	17	35.5	2 8 7 -	1
Argon (Ar)	18	40	2 8 8 -	0