Topics to be learn : Charged Particles in Matter The Structure of an Atom How are Electrons Distributed in Different Orbits (Shells)? Valency Atomic Number and Mass Number Isotopes

Introduction :

• Atoms and molecules are the fundamental building blocks of matter, with different types of matter around us due to the presence of different types of atoms and molecules.
• Dalton initially assumed that atoms were indivisible, but experimental evidence revealed sub-atomic particles like electrons, protons, and neutrons. The number of these particles varies among different elements.

Charged Particles in Matter

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

Rubbing two objects together causes them to become electrically charged, indicating the presence of charged particles within an atom. Two such particles are electrons and protons.

Discovery of Electrons : The electron, was identified by J. J. Thomson. These are

negatively charged particles and are denoted by symbol 'e⁻'. Its charge is taken as −1. The mass of an electron is considered to be negligible.

Discovery of Protons :

• Before the identification of electron, E. Goldstein in 1886, discovered the presence of new radiations in a gas discharge and called them canal rays or anode rays.
• These positively charged radiations move from the anode towards the cathode in a specially designed discharge tube with a porous cathode.
• This led to the discovery of the proton denoted by symbol 'p⁺', a sub-atomic particle with a charge of +1 and a mass approximately 2000 times larger than the electron. The porous cathode provided the path for passing anode rays.

The Structure of an Atom

According to Dalton's atomic theory, atom was indivisible and indestructible. The discovery of two fundamental particles (electrons and protons) inside the atom, led to the failure of this aspect of Dalton's theory.

To know the arrangement of electrons and protons within an atom, many scientists proposed various atomic models. These models are as follows

Thomson’s Model of an Atom :

• J. Thomson was the first scientist to propose a model for the structure of an atom. Thomson's model of an atom was similar to Christmas pudding. The electrons in a sphere of positive charge, were 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 studded in the positively charged sphere, like the seeds in the watermelon.

Thomson proposed that :

• Electrons are embedded in the sphere of positive charge of an atom.
• The negative and positive charges are equal in magnitude. Therefore, the atom as a whole is electrically neutral.
• The mass of an atom is assumed to be uniformly distributed throughout the atom.

Limitations of Thomson's Model of an Atom :

• Thomson's model could not explain the experimental results of other scientists such as Rutherford, as there is no nucleus in the atomic model proposed by Thomson.
• It 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 :

Rutherford's Atom Model Experiment :

• Ernest Rutherford designed an experiment to understand electron arrangement within an atom.
• He bombarded fast-moving a-particles on a thin gold foil.
• A-particles are double-charged helium ions with a mass of 4u.
• Rutherford selected a gold foil about 1000 atoms thick for thinness.
• He anticipated a-particles to be deflected by sub-atomic particles in gold atoms, causing small deflections due to their heavier weight.

Observations of Rutherford Experiment :

∝-scattering experiment and gave totally unexpected results which are as follows

• Most of the fast moving a-particles passed straight through the gold foil.
• Some of the ∝-particles were deflected by the foil by small angles.
• Very few a-particles (one out of 12000) appeared to rebound.

Conclusion of Rutherford Experiment :

On the basis of his experiment, Rutherford concluded that

• Most of the space inside the atom is empty. This is because most of the ∝-particles passed through the gold foil without getting deflected.
• Very few particles were deflected from their path. This indicate that the positive charge of the atom occupies very little space.
• A very small fraction of ∝-particles were deflected by 180° (i.e. they rebound). This indicate that all the positive charge and mass of gold atom were concentrated in a very small volume within the atom.
• From the data he also calculated that the radius of the nucleus is about 10⁵ times less than the radius of the atom.

Rutherford's Nuclear Model of an Atom :

On the basis of his experiment, Rutherford put forward the nuclear model of an atom, having the following features

• There is a positively charged, highly densed centre in an atom, called nucleus. Nearly, the whole mass of the atom resides in the nucleus.
• The electrons revolve around the nucleus in circular paths.
• The size of the nucleus (10-15 m) is very small as compared to the size of the atom (10-10 m).

Note : Rutherford suggested that his model of atom was similar to that of 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 :

The limitations of Rutherford's model of an atom are

• 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. Therefore, matter would not exist, but this is not so. It means that atoms are quite stable. Thus, it could not explain the stability of an atom when charged electrons are moving under the attractive force of positively charged nucleus.
• Rutherford's model could not explain the distribution of electrons in the extra nuclear portion of the atom.

Bohr’s Model of 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

• Atom consists of positively charged nucleus around which electrons revolve in discrete orbits, i.e. electrons revolve in certain permissible orbits and not just in any orbit.
• Each of these orbits are associated with 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 stationary state.
• Starting from nucleus, energy levels (orbits) are represented by numbers (1, 2, 3, 4 etc.) or by alphabets (K, L, M, N etc.).
• The electrons present in first energy level (E₁) have lowest energy. Energies increases on moving towards outer energy levels.
• Energy of an electron remains same as long as it remains in discrete orbit and it does not radiate energy while revolving.

Neutrons :

In 1932, J. Chadwick discovered a neutron, a subatomic particle with no charge and mass nearly equal to a proton.

Neutrons are present in all atoms' nuclei, except hydrogen. They are represented as 'n' and the mass of an atom is determined by the sum of protons and neutrons.

How are Electrons Distributed 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 as follows

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 x (1)² = 2

Second orbit or L-shell = 2 x (2⁾²= 8

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

Fourth orbit or N-shell = 2 x (4)² =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 (i.e. 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. 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 combining capacity or 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. It means, atom react with other atoms in order to attain fully-filled outermost shell.
• An outermost shell, which had eight electrons is called an octet. This was done by sharing, gaining or losing 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, valency is the combining capacity of the atom of an element with the atom(s) of other element(s) in order to complete its octet.
• 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.
• Similarly, 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.

Atomic Number and Mass Number

Atomic Number :

It is defined as the number of protons present in the nucleus of an atom.

All the atoms of the same element have the same number of protons in their nuclei and hence, they have the same atomic number.

It is denoted by Z and written as a subscript to the left of the symbol.

∴ Atomic number = number of protons

Examples :

• Atomic number of hydrogen is 1 (Z = 1) as there is only one proton present in the nucleus of its atom.
• Similarly, for carbon, Z = 6 as there are 6 protons present in the nucleus of its atom.

Mass Number :

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

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

For example, carbon has 6 protons and 6 neutrons, thus, its mass number is 6 + 6 = 12 u.

Number of neutrons = Mass number - Atomic number

(∵ Atomic number = Number of protons)

Mass number is written as a superscript to the left of the symbol. 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 \(_7^{14}\)N

Example : An atom of an element A may be written as \(_{12}^{24}\)A.

The superscript 24 indicate, 24 is the mass number of atom A.

The subscript 12 indicate, 12 is the atomic number of atom A.

Number of protons = 12

Number of neutrons = Mass number – Atomic number = 24 – 12 = 12

Number of electrons =12

Electronic configuration of given atom = K-2, L-8, M-2

It can lose two electrons (and attain stable configuration), therefore its symbol of its ion is A²⁺.

Isotopes :

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

e.g. There are 3 isotopes of hydrogen atom, namely protium (\(_1^1\)H), deuterium (\(_1^2\)H or D) and tritium (\(_1^3\)H or T) and 2 isotopes of carbon, \(_{6}^{12}\)C, \(_{6}^{14}\)C.

• In other words, it can be said that isotopes have same number of protons but differ in the number of neutrons.
• Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance.
• The chemical properties of isotopes of an element are similar as they contain same number of valence electrons.
• They differ in their physical properties because of having different atomic masses.

Average Atomic Mass : The mass of an atom of any natural element is taken as

the average mass of all the naturally occurring atoms of that 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

Average atomic mass of an element = [(Atomic mass of isotope I x percentage of isotope I) + (Atomic mass of isotope II x percentage of isotope II) + ... ]

Example : The two isotopic forms of chlorine atom with masses 35u and 37u occur in the ratio of 3 : 1.

∴ The average atomic mass of chlorine atom

= \((35×\frac{75}{100}+37×\frac{25}{100})\) = \((\frac{105}{4}+\frac{37}{4}) = \frac{142}{4}\) = 35.5u

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

Applications of Isotopes :

• Uranium-235 (U-235) : Used as fuel for the production of electricity in nuclear reactions.
• Uranium-238 (U-238) : Used in determining the age of very old rocks and even the age of the earth.
• Cobalt-60 (Co-60) : Used in treatment of cancer.
• Cobalt-14 (C-14) : Used in determining the age of old specimens of wood or old bones of living organisms.
• Iodine-131 (1-131) : Used in treatment of goitre.

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. e.g. \(_{18}^{40}\)Ar and \(_{20}^{40}\)Ca are isobars have same mass number, i.e. 40 but different atomic number, i.e. 18 and 20 for Ca.

Since, isobars have different atomic number as well as different electronic configuration. Thus, they also have different chemical properties.