Chaptemr 4 Structure of the Atom

Topics

• 1. Introduction
• 2. Structure of Atom
• 3. Discovery of Electrons
• 4. Discovery of Proton
• 5. Discovery of Neutron
• 6. Thomson’s Model of The Atom
• 7. Rutherford’s Experiment-Discovery of Nucleus
• 8. Bohr’s Model of Atom
• 9. Atomic Number and Mass Number
• 10. Isotopes and Isobars
• 11. Arrangement of electrons in the atoms
• 12. Valence Electrons

Introduction

• Atoms and molecules are the building blocks of matter.
• The word atom has been derived from the Greek word “a-tomio” which means “un-cuttable” or “non-divisible”.

• Earlier it was believed that atoms are indivisible and don’t have any inner structure.
• At the end of 19^thcentury scientists carried out the experiments which proved that the atoms are divisible.
• Dalton’s atomic theory which was proposed by John Dalton, considered the atom as the ultimate particle of matter.

Structure of Atom

• Atoms are made up three of subatomic particles: electrons, protons and neutrons.
• Electrons have negative charge, protons have positive charge and neutrons are neutral (no charge) in nature.
• Protons and neutrons are present in small nucleus at the centre of the atom.
• Whole of mass of atom is concentrated in the nucleus while the electrons, which are outside the nucleus, have very, very, small mass.
• Electrons are found outside the nucleus. The electrons move around the nucleus in fixed circular orbits called energy levels or shells.

• The atoms of different elements differ in the number of protons, electrons and neutrons.
• Note: – Hydrogen atom is made up of only one electron and one proton.

Discovery of Electrons

• The existence of electrons in atom was shown by J.Thomsonin 1897.

• Streams of particles were produced at cathode (negative electrode) when electricity is passed at a very high voltage.
• As they originated at cathode so they were known as cathode rays.

• Cathode rays consist of small, negatively charged particles called electrons.
• Since all the gases form cathode rays, it was concluded that all the atoms contain negatively charged particles called electrons.
• The electrons are negatively charged particles found in atoms of all elements.
• Electrons are represented as e^–.
• The charge on electron is -1.6 x 10^-19
• The relative charge on electron is (-1).
• The relative mass of electron = (1/1836) u and the absolute mass of electron is 9 x 10^-28

Discovery of Proton

• The existence of protons in the atoms was given by Goldstein.
• When electricity was passed at high voltage through a gas at a very low pressure in a discharge tube, streams of heavy particles were given at anode (positive electrode).These streams of particles are called as anode rays.
• These particles were positively charged and the anode ray particles depend on the nature of the gas taken in the discharge tube.
• Note: – The mass and charge of the anode ray particles depends on the nature of the gas taken in the discharge tube.
• They are named as protons which are positively charged and it is found in all the atoms.
• Protons are represented as p⁺.
• Mass of protonis equal to the mass of hydrogen atom.
  • Therefore relative mass of proton is 1u.
  • Absolute mass of proton is 1840 times that of the electron.
• Charge on electron is equal and opposite to the charge of the electron.
• The charge on proton = 1.6 x 10^-19
• Relative charge is (+1) and the absolute mass of electron is 1.607 x 10^-24

Discovery of Neutron

• James Chadwick in 1932 discovered the neutrons which are the neutral part found in the nucleus of an atom.

• Atoms of all the elements contain neutrons except ordinary hydrogen atom which does not contain any neutron.
• Symbol for neutron is “n”.
• This subatomic part is not present in hydrogen atom.
• Mass of a Neutron:-
  • The relative mass of neutron is 1u. The mass of neutron is equal to the mass of proton.
  • The absolute mass of neutron is 1.6 x 10^-24
• Charge of a Neutron:-
  • Neutron has no charge.
  • It is electrically neutral.
• Atomic Mass of Carbon = Mass of 6 protons + Mass of 6 neutrons
  • = (6 x 1) + ( 6x 1)
  • = 12u

Thomson’s Model of The Atom

• Thomson was the first scientist to prepare a model for the structure of an atom.
• At that time only electrons and protons were discovered.
• Thomson’s model of the atom is similar to that of Christmas pudding. The electrons are embedded in a sphere of positive charge are like currants (dry fruits) in a spherical Christmas pudding.
• Thomson model can also be compared to a watermelon.
• According to this model:-

1. An atom consists of a sphere of positive charge with negatively charged electrons embedded in it.
2. The positive and negative charges in atom are equal in magnitude due to which an atom is electrically neutral. It has no overall positive or negative charge.
3. Thomson’s model explained the electrical neutrality but failed to explain the various experiments carried out by other scientists like Rutherford.

Rutherford’s Experiment-Discovery of Nucleus

Alpha particle

• Alpha particle is a positively charged particle having 2 units of positive charge and 4 units of mass.
• It is actually helium ion, He²⁺.
• Rutherford’s alpha particles scattering experiment led to the discovery of small positively charged nucleus in the atom which contained protons and neutrons.

• When fast moving alpha particles are allowed to strike a thin gold foil of about thousand atoms thick in vacuum, it is found that:-

1. Most of the fast moving alpha particles passed straight through gold foil.
2. Some of the alpha particles were deflected by small angles.
3. About 1 out of every 12000 particles appeared to rebound.

• Conclusion:-

1. Most of the space in the atom is empty.
2. The positive charge of the atom occupies very little space.
3. The electrons revolve around the nucleus in circular paths.
4. All the positive charges and the mass of the atom were concentrated in a very small volume at the centre.
5. It shows the presence of a nucleus in the atom.
   1. Nucleus of an atom is positively charged.
   2. Nucleus of an atom is very dense and hard.
   3. The size of the nucleus is very small compared to the size of the atom.

• Drawbacks:-

1. The major drawback of this model is that it did not explain the stability of the atom.
2. According to the electromagnetic theory if a charged particle undergoes acceleration then it must radiate energy continuously.
3. Thus, the revolving electron loses energy and falls into the nucleus.
4. Then the atom would be highly unstable and hence matter could not exist in a form that is known now but it is known that atoms are stable.

Bohr’s Model of Atom

• According to Neil’s Bohr the electrons could revolve around the nucleus in only certain orbits known as discrete orbits of electrons, each orbit having different radii.

• When an electron is revolving in a particular orbit around the nucleus the electron does not radiate energy.
• Bohr’s model of an atom can be described as follows:-

1. An atom is made up of three sub-atomic particles – protons, electrons, neutrons.
2. The protons and neutrons are located in a small nucleus in the centre of the atom.
3. The electrons revolve around the nucleus rapidly in fixed circular paths called energy levelsor shells.
4. There is limit in the number of electrons that each shell can hold. It is given by the formula 2n², where n = orbit number or energy level index.
5. Each energy level is associated with a fixed amount of energy.
6. There is no change in the energy of the electrons as long as they keep revolving in the same energy level, and the atom remains stable.

A few energy levels in an atom

Atomic Number and Mass Number

• Atomic number: –
  • Atomic number constitutes the total number of protons which are present in the nucleus of that atom.
  • It is denoted by ‘Z’.
• Atomic mass:-
  • Atomic mass is the total number of neutrons and protons which are present inside the nucleus.
  • Mass of electrons is not considered while calculating the mass of the atom and only the mass of neutrons and protons are considered;
  • Since the electrons are the lightest particles their mass is not considered.
  • It is also known as Mass Number.
  • It is denoted by ‘A’.
• Nucleonsà Protons + Neutrons
• General representation of the element: – (^AZ_X) where A = atomic mass and Z = atomic number.
  • For example:- Hydrogen ₁¹H where atomic number=1 and mass number =1
  • Oxygen ¹⁶₈O where atomic number=8 and mass number =16(8 protons and 8 neutrons).

Isotopes and Isobars

• Isotopes: – Two nuclei with the same atomic number and different mass number are isotopes of each other.
  • For example: – There are 3 isotopes of carbon(C) having same atomic number 6 but their mass numbers are different i.e. 12, 13 and 14. (¹²₆C), (¹³₆C), (¹⁴₆C).
  • 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 35u and 37u in the ratio 3:1.

Applications of Isobars

1. Isotope of Uranium is used as a fuel in nuclear reactors.
2. Isotope of Cobalt is used in the treatment of cancer.
3. Isotope of Iodine is used in the treatment of goitre.

• Isobars: – The nuclei which have different atomic number but same mass number are known as isobars.
  • For example: – Nitrogen (¹⁴₇N) and Carbon (¹⁴₆C) are both isobars as their mass numbers are same which is 14 but their atomic numbers are 7 and 6 respectively.
  • The total number of nucleons is the same in the atoms of these pairs of elements.

Arrangement of electrons in the atoms

• Bohr and Bury’s rules for writing the number of electrons in different shells:-

1. The maximum number of electrons present in a shell is given by 2n², where n is the orbit number or energy level index, n=1,2,3,…

The maximum number of electrons in different shells can be given as:-

K- Shell –n =1 -> 2n²=2 x (1)²= 2electrons

L- Shell – n = 2 -> 2n²= 2x (2)²= 8 electrons

M- Shell – n = 3 -> 2n²= 2x (3)²= 18 electrons

N- Shell – n = 4 -> 2n²= 2x (4)²= 32 electrons

And so on

2. The maximum number of electrons that can be accommodated in the outermost orbit is eight.
3. Electrons are not accommodated in a given shell unless the inner shells are filled, i.e. the shells are filled in a stepwise manner.

Schematic atomic structure of the first eighteen elements

Valence Electrons

• The electrons present in the outermost shell of an atom are known as valence electrons (or valency electrons) because they decide the valency (combining capacity) of the atom.
  • For Example: – The atomic number of sodium is 11, which means sodium atom has 11 electrons in it.
  • So the electronic configuration K (2), L (8), M (1).
  • In the sodium atom, M shell is the outermost shell or valence shell.
  • There is 1 electron in the outermost shell of sodium atom; therefore, sodium atom has 1 valence electron.
• Those electrons of an atom which take part in chemical reactions are called valence electrons.
• Valence electrons are located in the outermost shell of an atom.
• Note:-
• In order to find out the number of valence electrons in an atom of the element, we should write down the electronic configuration of the element by using its atomic number.
• The outermost shell will be the valence shell and the number of electrons present in it will give us the number of valence electrons.
• Valency of a metal = No. of valence electrons in the atom
• Valency of a non-metal = (8 – No. of valence electrons in its atom)