4. Structure of the Atom | Class 9 Science | PDF and Web notes

 👉 View PDF

  Chapter 4: STRUCTURE OF THE ATOM

  CHARGED PARTICLES IN MATTER

 Activity to understand the nature of charged particles in matter:

A.    Comb dry hair. It can attract small pieces of paper.

B.    Rub a glass rod with a silk cloth and bring the rod near an inflated balloon. Balloon is attracted to the rod.

-    It means that rubbing two objects together, they become electrically charged.

-    The charge originates due to the fact that an atom is divisible and consists of charged particles.

-    By 1900, it was understood that atoms contain sub-atomic particles.

-    J.J. Thomson identified the electron (e^–).

-    In 1886, E. Goldstein had discovered positively charged radiations (canal rays) in a gas discharge. It led to the discovery of proton (p⁺).

-    The proton has a charge equal in magnitude but opposite in sign to the electron, with a mass approximately 2000 times greater.

-    Mass of a proton is taken as one unit and its charge as +1. Mass of an electron is negligible and its charge is -1.

-    An atom appears to be composed of protons and electrons, mutually balancing their charges.

-    Protons are located in the interior of the atom. So it is harder to remove them compared to electrons.

THE STRUCTURE OF AN ATOM

-    Dalton's atomic theory says that the atom is indivisible and indestructible, but this aspect failed with the discovery of electrons & protons (fundamental particles).

-    J.J. Thomson was the first to propose a model for the structure of an atom.

THOMSON’S MODEL OF AN ATOM

-    Thomson proposed a model of the atom resembling a Christmas pudding. Here, electrons are embedded like currants in a sphere of positive charge (pudding).

-    This model can also be compared to a watermelon. Here, the positive charge spread throughout like the red edible part, and the electrons are like seeds within the positively charged sphere.

Thomson’s model of an atom

Thomson proposed that:

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

(ii)    The negative & positive charges are equal in magnitude. So, the atom is electrically neutral.

However, this model could not explain some results of experiments carried out by other scientists.

RUTHERFORD’S MODEL OF AN ATOM

-    Ernest Rutherford conducted an experiment to know the arrangement of electrons in an atom.

-    He directed fast-moving alpha (α)-particles on a thin gold foil, which was about 1000 atoms thick.

-    α-particles are doubly-charged helium ions. They have a mass of 4 u. This gives them a considerable amount of energy when moving at high speeds.

-    It was expected that α-particles would be deflected by the sub-atomic particles in the gold atoms. Since α-particles are much heavier than the protons, only small deflections were expected.

-    But some unexpected observations were made:

·  Most α-particles passed straight through the gold foil.

·  Some α-particles were deflected by small angles.

·  Surprisingly, about 1 in 12,000 α-particles rebounded.

Scattering of α-particles by a gold foil

(α-particle scattering experiment)

-    Rutherford described the experiment's results as nearly as incredible as firing a 15-inch shell at a piece of tissue paper and having it rebound.

-    To illustrate the implications, imagine a blindfolded child throwing stones at a wall and hearing a sound each time a stone hits. If the child threw stones at a barbed-wire fence with many gaps, most stones would pass through without hitting anything, resulting in no sound.

-    Applying this reasoning to the α-particle scattering experiment, Rutherford concluded:

·   Most of the atom's space is empty because most α-particles passed through gold foil without deflection.

·   Only a few particles were deflected from their path. It indicates that the positive charge of the atom occupies very little space.

·   A tiny fraction of α-particles was deflected by 180°. It indicates that the positive charge and mass of the atom are concentrated in a very small region.

-    Rutherford estimated that the radius of nucleus is about 10⁵ times smaller than the radius of the atom.

-    Based on his findings, Rutherford proposed the nuclear model of the atom:

·   Atoms have a positively charged nucleus where nearly all the mass of an atom resides.

·   Electrons revolve around the nucleus in circular paths.

·   Nucleus is very small compared to the size of atom.

Drawbacks of Rutherford’s model of the atom

-    Electrons in circular orbits are not expected to be stable

because they undergo continuous acceleration.

-    Accelerating charged particles, like electrons, would radiate energy, causing them to lose energy and finally fall into the nucleus. This would result in highly unstable atoms. It is against the fact that atoms are stable.

BOHR’S MODEL OF ATOM

In order to overcome the drawbacks of Rutherford’s model, Neils Bohr put forward the following postulates:

    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.

A few energy levels in an atom

They are represented by the letters K, L, M, N, or the numbers, n=1, 2, 3, 4, ….

NEUTRONS

-    Neutron (n) is a subatomic particle discovered by J. Chadwick (1932). It has no charge.

-    They are seen in the nucleus of all atoms except hydrogen.

-    Neutron has a mass nearly equal to that of a proton. So, the mass of an atom is the sum of the masses of protons and neutrons in the nucleus.

  HOW ARE ELECTRONS DISTRIBUTED IN DIFFERENT ORBITS (SHELLS)?

Bohr and Bury proposed the following rules for the distribution of electrons in an atom's orbits:

·    The maximum number of electrons in a shell is determined by the formula 2n², where n is the orbit number or energy level index:

K-shell (1^st orbit): 2 × 1² = 2 electrons

L-shell (2^nd orbit): 2 × 2² = 8 electrons

M-shell (3^rd orbit): 2 × 3² = 18 electrons

N-shell (4^th orbit): 2 × 4² = 32 and so on.

·    The outermost shell can accommodate a maximum of 8 electrons.

·    Electrons fill inner shells before moving to outer shells. That is, the shells are filled in a step-wise manner.

Schematic atomic structure of the first eighteen elements

Composition of Atoms of the First Eighteen Elements with Electron Distribution in Various Shells

Name of Element	Symbol	Atomic Number	Number of Protons	Number of Neutrons	Number of Electrons	Distribution of electrons				Valency
						K	L	M	N
Hydrogen	H	1	1	-	1	1	-	-	-	1
Helium	He	2	2	2	2	2	-	-	-	0
Lithium	Li	3	3	4	3	2	1	-	-	1
Beryllium	Be	4	4	5	4	2	2	-	-	2
Boron	B	5	5	6	5	2	3	-	-	3
Carbon	C	6	6	6	6	2	4	-	-	4
Nitrogen	N	7	7	7	7	2	5	-	-	3
Oxygen	O	8	8	8	8	2	6	-	-	2
Fluorine	F	9	9	10	9	2	7	-	-	1
Neon	Ne	10	10	10	10	2	8	-	-	0
Sodium	Na	11	11	12	11	2	8	1	-	1
Magnesium	Mg	12	12	12	12	2	8	2	-	2
Aluminium	Al	13	13	14	13	2	8	3	-	3
Silicon	Si	14	14	14	14	2	8	4	-	4
Phosphorus	P	15	15	16	15	2	8	5	-	3, 5
Sulphur	S	16	16	16	16	2	8	6	-	2
Chlorine	Cl	17	17	18	17	2	8	7	-	1
Argon	Ar	18	18	22	18	2	8	8		0

  VALENCY

-    The electrons present in the outermost shell of an atom are called the valence electrons.

-    The combining capacity of an atom of each element is called its valency. i.e., it is the tendency to react and form molecules with atoms of the same or different elements.

-    Atoms of elements, completely filled with 8 electrons in the outermost shell show little chemical activity. They are called inert elements. It means their valency is zero.

-    Of the inert elements, the helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons.

-    Among inert elements, the helium atom has 2 electrons in its outermost shell. Other inert elements have 8 electrons in their outermost shell.

-    Thus the valency of an atom is an attempt to attain a fully-filled outermost shell.

-    An outermost shell with 8 electrons is termed as octet.   

-    Atoms react to achieve octet by sharing, gaining, or losing electrons. The number of electrons gained, lost, or shared to complete the octet directly indicates an element's valency. E.g.,

·   Hydrogen/lithium/sodium atoms have one electron in their outermost shell and can lose that one electron. So its valency = 1.

·   Magnesium has 2 electrons. So its valency = 2

·   Aluminum has 3 electrons. So its valency =  3

-    If an atom's outermost shell is nearly full, its valency is calculated differently. E.g.,

·   Fluorine has 7 electrons in the outermost shell. But it is easier to gain one electron instead of losing 7 electrons. Hence, its valency is 8 – 7 = 1.

·   Oxygen atom has 6 electrons in the outermost shell. It gains 2 electrons instead of losing 6 electrons. Hence its valency is 8 – 6 = 2.

  ATOMIC NUMBER AND MASS NUMBER

ATOMIC NUMBER (Z)

-    It is the total number of protons present in the nucleus of an atom.

-    All atoms of an element have same atomic number (Z).

-    Elements are defined by the number of protons they possess. E.g.,

·    Z= 1 for hydrogen (it has only one proton).

·    Z = 6 for carbon (it has 6 protons).

MASS NUMBER (A)

-    It is 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 & neutrons. Being present in the nucleus, they are also called nucleons. i.e., the mass of an atom resides in its nucleus. E.g.,

·    Mass (A) of carbon is 12 u because it has 6 protons and 6 neutrons (6 u + 6 u = 12 u).

·    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:

  ISOTOPES

-    Isotopes are the atoms of the same element, having the same atomic number but different mass numbers. E.g.,

·     Hydrogen atom has three isotopes, protium (H), deuterium (H or D) and tritium (H or T).

They have same atomic number (Z=1) but the mass number is 1, 2 and 3, respectively.

·     Carbon:C and C.

·     Chlorine: 35 Cl and Cl.

-    Many elements consist of a mixture of isotopes. Each isotope of an element is a pure substance. Isotopes have similar chemical properties but different physical properties.

-    The mass of an atom of a natural element is taken as the average mass of all the naturally occurring atoms of that element.

-    If an element has no isotopes, its atomic mass equals the sum of its protons and neutrons. But if an element occurs in isotopic forms, then we have to know the percentage of each isotope and then the average mass is calculated. E.g.,

Chlorine has 2 natural isotopes with masses 35 u and 37 u in the ratio of 3:1. So the average atomic mass is

-    Since the chemical properties of all the isotopes of an element are the same, normally we are not concerned about taking a mixture.

Applications of some isotopes:

·   An isotope of uranium is used as fuel in nuclear reactors.

·   An isotope of cobalt is used in cancer treatment.

·   An isotope of iodine is used in treatment of goitre.

ISOBARS

-    These are atoms of different elements that have different atomic numbers but the same mass number.

-    E.g., calcium (Z= 20) and argon (Z= 18). They have different number of protons, but the mass number is 40. i.e., total number of nucleons is the same in them.

 👉 View PDF