# Structure of Atom

The atoms consist of electrons, protons and neutrons. There is an interesting story on how electrons, protons and neutrons were discovered and how the structure of atom has come to be. It’s worth knowing about the fundamental particles of matter. ## 1. Charges Particles in Matter

The particles that carry an electric charge are called charged particles.

On rubbing two objects together, they become electrically charged.  It means that some charged particles are present within the atom i.e., atom is made up of some charged particles. Two such particles are electrons and protons. So, let’s find out more about electrons and protons.

DISCOVERY OF ELECTRONS

In the year 1900, J. J. Thomson performed an experiment called cathode ray experiment also called the discharge tube made up of glass.

In the experiment, a gas at low pressure was taken in a discharge tube. At the ends of the discharge tube, two electrodes (metal plates) were placed, connected to a battery for high voltage supply. The electrode connected to the negative end was known as cathode and that to the positive as anode.

During this experiment, Thomson found a beam originating from the cathode of the discharge tube consisting of negatively charged particles which he called cathode rays. These negatively charged particles were called electrons.

Electrons are negatively charged particles and are denoted by e- ‘. The charge on an electron is equal to -1.6 × 1019  Coulomb. Since, this charge is considered to be the smallest, therefore, charge on eis taken as -1.  The mass of an electron is equal to 9.1 × 10-31kg.

DISCOVERY OF PROTON

We know an atom is electrically neutral. Cathode ray experiment showed that all the atoms contain negatively charged electrons.  So, atoms must also contain some positively charged particles to balance the negative charge of electrons. This was the basis of the discovery of protons.

In the year 1886, before the electrons were discovered by J.J. Thomson, E. Goldstein discovered the presence of a new kind of rays known as canal rays or anode rays.  These rays were positively charged as they were seen moving from the anode towards cathode in a specially designed discharge tube (with a porous cathode), when a high voltage is applied across the electrodes.  Porous cathode is used to provide the path for passing anode rays. It led to the discovery of another sub-atomic particle, the proton.

Protons are positively charged particles and are denoted by ‘p+’.  The charge present on a proton is equal to +1.6 × 10-19 Coulomb and it is considered as +1.  The mass of a proton is equal to 1.6 × 10-27 kg.  The mass of a proton is approximately 2000 times as that of the electron.

## 2. Structure of Atom

According to Dalton’s atomic theory, an atom was indivisible and indestructible. Now, the discovery of two fundamental particles (electrons and protons) inside the atom, led to the failure of this aspect of Dalton’s theory. To explain the arrangement of electrons and protons within an atom, many scientists proposed various atomic models.

THOMSON’S MODEL OF AN ATOM

J.J. Thomson was the first scientist to propose a model for the structure of an atom sometimes called the ‘plum pudding model’.  The electrons are placed in a sphere of positive charges, which was like currants (dry fruits) in a spherical Christmas pudding.

It can also be compared to a watermelon, in which the positive charge in an atom is spread all over like the red edible part, while the electrons are studded in the positively charged sphere, like the seeds in the watermelon. The model was based on the following postulates:

(i) Electrons are embedded in the sphere of positive charges.

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

(iii) The mass of an atom is assumed to be uniformly distributed throughout the atom.

Limitations of Thomson’s model (Plum Pudding model) of an atom are:

(i)  It could not explain the experimental results of other scientists such as Rutherford, as there was no nucleus in the atomic model proposed by Thomson.

(ii) I could not explain the stability of an atom, i.e. how positive and negative charges could remain, so close together.

RUTHERFORD’S MODEL OF AN ATOM

Ernest Rutherford designed an experiment to know how the electrons are arranged within an atom. It is known as the ‘Rutherford’s gold foil experiment’.  He bombarded fast moving a-particles (alpha particles – doubly-charged helium ions having a mass of 4 u) on a thin sheet of gold foil.

(i) Since a-particles have a mass of 4 u, the fast-moving α-particles have a considerable amount of energy.

(ii) He selected a gold foil because he wanted a layer as thin as possible. This gold foil was about 1000 atoms thick.

(iii) Since the α-particles were much heavier than the protons, he did not expect to see large deflections.

The following observations were made by Rutherford during his experiment:

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

(ii) Some of the α-particles were deflected by the foil by small angles.

(iii)  Very few α-particles (one out of 12000) appeared to rebound.

On the basis of his experiment, Rutherford concluded that:

(i) Most of the space inside the atom is empty because most of the a-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 a-particles was deflected by 180° (i.e. they rebound), indicating that all the positive charge and mass of atom were concentrated in a very small volume within the atom. On the basis of his experiment, Rutherford put forward the ” nuclear model of an atom”, having the following features:

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

(ii) The electrons revolve around the nucleus in circular path.

(iii) The size of the nucleus (10-15 m) is very small as compared to the size of the atom (10-10m).

Note   Rutherford suggested that his model of the atom was similar to that of the solar system. In the solar system, the different planets are revolving around the Sun. Similarly, in an atom, the electrons are revolving around the nucleus.  So, these electrons are also called planetary electrons.

Limitations of Rutherford’s Model of an Atom

(i) Any charged particle when accelerated is expected to radiate energy. To remain in a circular orbit, the electron would need to undergo acceleration. Therefore, it 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, which is not the case as atoms are stable. Thus, it could not explain the stability of an atom.

(ii) The model did not mention anything about the arrangement of electrons in the orbit.

BOHR’S MODEL OF AN ATOM

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:

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

(ii) While revolving in discrete orbits the electrons do not radiate energy. (iii) Each of these orbits is associated with a certain value of energy.  Hence, these orbits are called energy shells or energy levels. As the energy of an orbit is fixed (stationary), orbit is also called a stationary state.

(iv) Starting from the nucleus, energy levels (orbits) are represented by numbers (1, 2, 3, 4 etc.)  or by alphabets (K, L, M, N etc.).

(v) The electrons present in first energy level (E1) have the lowest energy.  Energies increases on moving towards outer energy levels.

(vi) Energy of an electron remains same as long as it remains in discrete orbit and it does not radiate energy while revolving.

(vii) When energy is supplied to an electron, it can go to higher energy levels. While an electron falls to a lower energy level when it radiates energy.

NEUTRONS (n)

Neutron was discovered by J. Chadwick in 1932. It is represented by n. Neutrons are electrically neutral particles and are as heavy as protons (i.e. their mass is 1.67493 × 10-27 kg). 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.

DISTRIBUTION OF ELECTRONS IN DIFFERENT ORBITS (Shells)

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

(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, 1, 2, 3,… Therefore, the maximum number of electrons in different shells are as follows:

First orbit or K-shell        = 2 × (1)2 = 2

Second orbit or L-shell = 2 × (2)2 = 8

Third orbit or M-shell    = 2 × (3)2 = 18

Fourth orbit or N-shell = 2 × (4)2 = 32 and so on.

(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 (i.e. the shells are filled in a stepwise manner).

VALENCY

The electrons present in the outermost shell of an atom are known as the valence electrons.  They govern the chemical properties of atoms.

The atoms of elements having completely filled outermost shell means which has eight electrons show little chemical activity, i.e. they are highly stable. Such elements are called inert elements. It means, their valency is zero. Of these inert elements, the helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell.

The tendency to react with atoms of the same or different elements to form molecules is an attempt to attain fully-filled outermost shell.  It means atoms react with other atoms in order to attain fully-filled outermost shell. An outermost shell, which had eight electrons is called an octet. Atoms would thus react, so as to achieve an octet in the outermost shell. This was done by sharing, gaining or the loss of electrons.

The number of electrons lost or gained or shared by an atom to become stable or to achieve an octet in the outermost shell is known as valency of that element. In other words, it is the combining capacity of the atom of an element with the atom(s) of other element(s) in order to complete its octet.

The   valencies of elements of some groups   are described below:

(i) Hydrogen (H), lithium (Li), sodium (Na) and potassium (K) atoms contain one electron each in their outermost shell, therefore, each one of them can lose one electron to become stable. Hence, their valency is 1.

(ii) The valency of each of Mg, Ca and Be is 2 because all of these have 2 valence electrons in the K-Shell and they can lose these 2 electrons to make the octet of electrons in the outermost shell or to become stable.

(iii) The valency of Boron (B)and Aluminium (Al) is 3 because each has 3 valence electrons.

(iv) The valency of carbon and silicon is 4 because each has 4 valence electrons. They can attain octet by sharing of electrons.

(v) Nitrogen and phosphorus each have 5 valence electrons, so their valency is 3 because they can gain 3 electrons (instead of losing five electrons) to become stable.   Hence, their valency is determined by subtracting five electrons from the octet, i.e. 8 – 5 = 3. However, P can also share 5 electrons, hence it shows a valency of 5 along with

(vi) Oxygen and sulphur each have 6 valence electrons; therefore, their valency is 2 because they can gain 2 electrons or share 2 electrons to complete their octet.

(vii) Similarly, fluorine and chlorine each have 7 valence electrons, their valency is l because they can gain 1 electron or share 1 electron to complete their octet.

(viii) All the inert elements, i.e. He, Ne, Ar etc., have completely filled outermost shells. Therefore, their valency is zero.

Note   For metals, valency= Number of valence electrons and for non-metals, valency= 8- number of valence electrons. ## 3. Atomic Number and Mass Number

ATOMIC NUMBER

It is defined as the number of protons present in the nucleus of an atom. It is denoted by Z.  Its written like this Form the figure, atomic number of Helium is 2.

MASS NUMBER

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

Mass number (A) = Number of protons + number of neutrons

From the figure the Mass number(A) of Helium is 4.

Number of neutrons= Mass number (A) ­ – Atomic number (Z)

## 4. Different Atomic Species

ISOTOPES

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

E.g. there are 3 isotopes of hydrogen atom, namely protium, deuterium and tritium.    Isotopes have same number of protons but differ in the number of neutrons. Each isotope of an element is a pure substance.

Chemical properties of elements largely depend on their electronic configuration or outermost electrons and as such the isotopes of an element have similar electronic configuration, therefore, isotopes of an element have same chemical properties. Since, physical properties such as density, light scattering etc., depend on mass, therefore, physical properties are different for isotopes of an element.

Average Atomic Mass of an element

If an element has no isotopes, the mass of its atom would be the same as the sum of masses of protons and neutrons in it. But if an element occurs in isotopic forms, then from the percentage of each isotopic form, the average mass is calculated as:

= [(Atomic mass of isotope I × percentage of isotope I) + (Atomic mass of isotope II  × percentage of isotope 11)+…]

Example:

There are  two isotopic forms of chlorine atom with masses 35u and 37u occurring  in the ratio of 3: 1Therefore, the average atomic mass of chlorine atom, can be calculated as:

The average atomic mass of chlorine atom = Here, 35.5 u is not the atomic mass of any one atom of chlorine but it shows that its given amount contains both the isotopes and their average atomic mass is 35.5 u.

Note: The fractional atomic masses of elements are due to the existence of their isotopes having different masses.

Applications of Isotopes

(i) An isotope of uranium (U-235) is used as a fuel for the production of electricity in nuclear reactors.

(ii) U-238 is used to determine the age of very old rocks and even the age of the earth·.

(iii) An isotope of cobalt (Co-60) is used in the treatment of cancer.

(iv) An isotope of carbon (C-14) is used to determine the age of old specimens of wood or old bones of living organisms.

(v) An isotope of iodine (1-131) is used in the treatment of goiter.

ISOBARS

Atoms of different elements with different atomic numbers but same mass number are known as isobars.

In other words, isobars are the atoms of different elements that have same number of nucleons (Protons +Neutrons) but differ in the number of protons.

Example: are isobars.

Since, isobars have different atomic numbers, therefore different electronic configuration, hence they have different Physical and Chemical properties. This site uses Akismet to reduce spam. Learn how your comment data is processed.