CLASS 9 Science
CHAPTER 4
Structure Of The Atom (NCERT Notes)
CHAPTER 4
Structure Of The Atom (NCERT Notes)
Charged Particles in Matter
On rubbing two objects together, they become electrically charged. Where does this charge come from? This question can be answered by knowing that an atom is divisible and consists of charged particles.
Many scientists contributed in revealing the presence of charged particles in an atom.
It was known by 1900 that the atom was not a simple, 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 peeled off but not protons.
The Structure of an Atom
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’S MODEL 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. Or watermelon, 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.
Thomson proposed that:
(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. Although Thomson’s model explained that atoms are electrically neutral, the results of experiments carried out by other scientists could not be explained by this model.
RUTHERFORD’S MODEL OF AN ATOM
Rutherford's model of an atom. In this experiment, fast moving alpham(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.
• 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.
The results α-particle scattering experiment gave totally unexpected . The following observations were made:
(i) Most of the fast moving α-particles passed straight through the gold foil.
(ii) Some of the α-particles were deflected by the foil by small angles.
(iii) Surprisingly one out of every 12000 particles appeared to rebound.
Rutherford conclusion from the α-particle scattering experiment that–
(i) Most of the space inside the atom is empty because most of the α-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 α-particles were deflected by 1800,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 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:
(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 of the atom
The orbital revolution of the electron 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
When Rutherford’s model of the atom was failed then 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.
These orbits or shells are called energy levels. Energy levels in an atom are shown in
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. In general, a neutron is represented as ‘n’. The mass of an atom is defined as the sum of the masses of protons and neutrons present in the nucleus.
How are Electrons Distributed in Different Orbits (Shells)?
According to Bohr and Bury theory: 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 index, 1,2,3,…. Hence the maximum number of electrons in different shells are as follows:
First orbit or K-shell will be = 2 × 12 = 2,
Second orbit or L-shell will be = 2 × 22 = 8,
Third orbit or M-shell will be = 2 × 32 = 18,
Fourth orbit or N-shell will be = 2 × 42= 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.
Valency
The electrons in an atom are arranged in different shells/orbits. The electrons present in the outermost shell of an atom are known as the valence electrons. According to Bohr-Bury the outermost shell of an atom can accommodate a maximum of 8 electrons. It was observed that the atoms of elements, having a completely filled outermost shell show little chemical activity. In other words, their combining capacity or valency is zero.
Inert elements, that is helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell. The combining capacity of the atoms of other elements, that is, their tendency to react and form molecules with atoms of the same or different elements, was thus explained as an attempt to attain a fully-filled outermost shell. An outermost-shell, which had eight electrons was said to possess an octet. Atoms would thus react, so as 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 us directly the combining capacity of the element, that is called the 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 are said to have valency of one. Valency of magnesium and aluminium is two and three, respectively, because magnesium has two electrons in its outermost shell and aluminium has three electrons in its outermost shell.
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 you a valency of one for fluorine.
Atomic Number and Mass Number
ATOMIC NUMBER
The is defined as the total number of protons present in the nucleus of an atom. Protons are present in the nucleus of an atom. It is the number of protons of an atom, which determines its atomic number. It is denoted by ‘Z’. All atoms of an element have the same atomic number, Z. Elements are defined by the number of protons they possess. 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. Mass of an atom is practically due to protons and neutrons alone. These are present in the nucleus of an atom. Hence protons and neutrons are also called nucleons. Therefore, the mass of an atom resides in its nucleus. 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).
In the notation for an atom, the atomic number, mass number and symbol of the element are to be written as:
An elements have the same atomic number but different mass numbers are called isotopes. For example Hydrogen atom, has three atomic species, protium (11H), deuterium (21H or D) and tritium ( 31H or T). The atomic number of each one is 1, but the mass number is 1, 2 and 3, respectively. Other such examples are (i) carbon, 12 6 C and 14 6 C, (ii) chlorine, 35 17 Cl and 37 17 Cl, etc.
Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance. The chemical properties of isotopes are similar but their physical properties are different. Chlorine occurs in nature in two isotopic forms, with masses 35 u and 37 u in the ratio of 3:1.
What should we take as the mass of chlorine atom?
The mass of an atom of any natural element is taken as the average mass of all the naturally occuring atoms of that element. If an element has no isotopes, then the mass of its atom would be the same as the sum of protons and neutrons in it. But if an element occurs in isotopic forms, then we have to know the percentage of each isotopic form and then the average mass is calculated. The average atomic mass of chlorine atom, on the basis of above data, will be
Applications
The chemical properties of all the isotopes of an element are the same, normally we are not concerned about taking a mixture. Some of them are :
(i) An isotope of uranium is used as a fuel in nuclear reactors.
(ii) An isotope of cobalt 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. The number of electrons in these atoms is different, the total number of nucleons is the same in the atoms of this pair of elements. For example calcium, atomic number 20, and argon, atomic number 18. The number of electrons in these atoms is different, but the mass number of both these elements is 40. That is, the total number of nucleons is the same in the atoms of this pair of elements.
( From NCERT Book )