Basic Atomic structure

Introduction to atomic structure(History ,components,Theories,Models Atomic mass etc)

1. ATOMIC STRUCTURE Vinay Desai M.Sc Radiation Physics KIDWAI MEMORIAL INSTITUTE OF ONCOLOGY

2. Every matter consists of basic entities called elements. Each element is composed of smallest particles called ‘ ATOM ’. Atom- the name is derived from Greek language Atomos means ‘ Not to be cut’. ATOM:

3. Atom and ‘DEMOCRITUS’ DEMOCRITUS was a Greek philosopher who began search for description of matter more than 2400 years ago (4 th century B.C.) Democritus was the person who first suggested the existence of ATOM & coined the name ÁTOMOS ’ means Not to be cut or Indivisible.

4. Atomic structure Atom consists of positively charged NUCLEUS at the centre and negatively charged Electrons revolving around it. Radius of an atom -- 10 -10 m. Radius of the nucleus -- 10 -15 m. N ucleus consists of Protons and Neutrons together called Nucleons. Most of the mass of an Atom is possessed by Nucleus.

5. Representation of an Atom : E xample : 235 U 92 X-Uranium A-235 Z-92

6. Theories of Atomic models: John Dalton Matter is made of indivisible atoms, they are indestructible. All atoms of a given (same) element are identical in their physica l and chemical properties. Atoms of different elements differ in their physical and chemical properties. Atoms of different elements combine in simple whole-numbers ratios to form Molecules Chemical reactions consist of the combination separation or rearrangement of atoms Limitations : It could not explain Why and how do atoms combine together to form compound atoms ( molecules) T he nature of forces which hold atoms together in compound atoms Why atoms cannot exist in free state and why compound atoms can exist freely.

7. Theories of Atomic models: J.J.THOMPSON Plum pudding model(1904) N egative particles are evenly scattered throughout an atom with a positively charged mass of matter. S imilar to that of chocolate chip icecream L ater proved to be incorrect. Limitations: It could not explain the result of scattering experiment explained by rhe Rutherford experiment. It did not have any experimental evidence in its support . T hermionic emission, photoelectric emission and ionization were explained on this basis.

8. Theories of Atomic models: RUTHERFORD Rutherford Gold foil experiment setup: If a thin foil metal is placed in the path of the beam, the image becomes diffuse. This due to the scattering of α - particles by the atoms of the foil. The particles scattered in various directions were counted by scintillation counter It was found that although most of the particles scattered through angles of the order of 1 or less But a small number say about 1 in every 10,000 scattered through 90 or even 180 The experiment is known as Rutherford's alpha particle scattering experiment .

9. Results of the Rutherford experiment (a ) The results that the metal foil experiment would have yielded if the plum pudding model had been correct . (b) Actual results

10. The whole of the positive charge of atom must be concentrated in a very small space Atom is mostly hollow inside Since α - particles are positively charged, the part of the atom deflecting them must also be positive Most of the mass of the atom is concentrated in the nucleus In this model, the mass of the atom (leaving the mass of its electrons) and its whole positive charge are concentrated at the centre of the atom in a nucleus of radius 10 -15 m The electrons are distributed around the nucleus in a hollow sphere of radius 10 -10 m Conclusion of Rutherford experiment

11. Regarding stability of atom Electrons revolving around the nucleus have centripetal acceleration According to electrodynamics, accelerated charged particles radiate energy in the form of electromagnetic waves Hence electromagnetic waves should be continuously radiated by the revolving electrons Due to this continuous loss of energy of the electrons, the radii of their orbits should be continuously decreasing and ultimately the electron should fall into the nucleus Thus atom cannot remain stable Rutherford’s model also failed to explain the Line spectrum. Drawback’s of Rutherford’s Model

12. Niels David Bohr Theories of Atomic models: Bohr’s theory (Postulates ) Fixed circular orbits :The electrons move around the nucleus in concentric circular orbits . While revolving in stable orbits, the electrons do not radiate energy in spite of their acceleration towards the centre of the orbit. Each of the fixed orbits is associated with a definite amount of energy called stationary energy. The energy levels are numbered as 1, 2, 3, 4… or designated as K,L,M , N … 4. Jumping of an electron from one energy level to the other (ground state and excited state) . 5 . Principle of quantization of angular momentum of the moving electron an electron can move only in that orbit in which the angular momentum of the electron around the nucleus is an integral multiple of h/2 π . NIELS DAVID BOHR A Danish physicist who developed Bohr model of atomic structure, in which he introduced the theory of electrons orbiting around the nucleus.

13. No explanation for the spectra of multi electron systems : Eg : He, Li No explanation of fine spectrum of atoms : No explanation for Zeeman and Stark effect : effect of electric and magnetic fields on the spectral atoms. When a magnetic field is applied on an atom, its usually observed spectral lines split. This effect is known as Zeeman’s effect Spectral lines also get split in the presence of electric field. This effect is known as Stark effect . Limitations of Bohr’s Postulates

14. Isotopes - Elements having same atomic number (protons) , but different mass numbers (nucleons). eg: 12 C, 13 C and 14 C are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of all carbon isotopes is 6. Isobars - Elements having same mass number , but different atomic numbers. eg: 40 S , 40 Cl , 40 Ar , 40 K, and 40 Ca are isobars containing 40 nucleons; however, they contain varying atomic number. Classification of atoms:

15. Isomers - Molecules with the same molecular formula but different chemical structures. That is, isomers contain the same number of atoms of each element, but have different arrangements of their atoms in space. eg: 131m Xe 54 is an isomer of 131 Xe 54 m stands for meta-stable state isomers represents identical atoms but they differ in nuclear energy states . Isotones - Nuclides having very same neutron number N , but different proton number Z. eg:boron-12 Carbon-13 nuclei both contain 7 neutrons

16. NUCLEAR STABILTY: The Strong Force is exerted by anything with mass to attract other masses together and works within a very short distance . Neutrons has no charge, but have the strong force to bring other nucleons together. As a general rule, a nucleus will need a neutron/proton ratio of 3:2 (or 1.5:1) in order to stay together. This rule is more precise for larger nuclei. Of all known isotopes of natural elements (about 1500), only 250 of them are stable

17. .All of these stable isotopes have an atomic number in between 1 and 83. The mass of a nucleus will be less than the mass of all of the protons and neutrons making it up. The difference is called the mass defect , which is converted into energy if the nucleus is broken up. The amount of energy that keeps a nucleons together is called the Binding Energy . This amount of energy is higher for nuclei that are stable than it would be for unstable nuclei. (Joules)

18. Binding energy can be calculated by the formula E=mc 2 ( Einstein,s principle of equivalence of mass & energy relation) Where, c= speed of light ,m=mass & E = Energy Eg: If a mass of 1 kg is converted into energy, m=1kg speed of light is 3x10 8 m/s E=1 kg x (3x10 8 m/s) 2 E= 9 x 10 16 J Also, Mass of electron at rest in terms of energy equivalent is given by, m=9.1x10 -31 kg and speed of light is 3x10 8 m/s E = 9.1x10 -31 x (3x10 8 ) 2 E=8.19x10 -14 J E=0.511 MeV

19. Masses of atoms & atomic particles are conveniently given in terms of amu. An amu is defined as 1/12 of the mass of the 12 C 6 Carbon atom. 1 amu= 1.66x10- 27 kg Mass expressed in terms of amu is known as atomic mass or atomic weight. Atomic mass and Energy Units

20. Charge and Mass of sub-atomic particles Particle Charge of particle Mass of the particle Proton 1.602176 x 10 -19 Coulomb 1.00727 amu Neutron Electrically neutral 1.00866 amu Electron 1.602176 x 10 -19 Coulomb 0.00548 amu Number of Protons = Number of Neutrons 1 amu = 1/12 th of mass of an carbon atom. 1 amu = 1.66 x10 -27 kg

21. Avagadro's law- Every gram atomic weight of a substance contains same number of atoms, the number is referred as Avagadro's number Value of Avagadro's number 6.0221x10 23 atoms per gram atomic weight Gram atomic weight

22. Basic unit of energy is joule (J) 1 Joule is the work done when force of 1 Newton is acting through a distance of 1 meter. Another energy unit in Nuclear physics is electron volt (eV) 1 eV is defined as the Kinetic energy acquired by an electron in passing through a potential difference of 1V. 1eV=1Vx1.602x10 -19 C 1eV = 1.602 x 10 -19 J Product of Potential difference and Charge of the electron Atomic energy unit - Joule

23. DISTRIBUTION OF ORBITAL ELECTRONS: Postulates of Bohrs theory : a) Electrons can exist in only those orbits for which angular momentum of electron is an integral multiple of h/2 Π (h= plancks constant 6.62x10 -34 ) b) No energy is gained or lost while an electron remains in any one of the permissible orbits. According to the model proposed by Niels Bohr in 1913 Electrons revolve around the nucleus in specific orbits. They are prevented from leaving the atom by the necessary centripetal force of attraction between the positively charged nucleus and negatively charged electron. Bohr’s Atomic model

24. Arrangement of electrons in orbitals Innermost orbit is called as K-shell. Followed by orbital’s called L-shell, M-shell and N - shell. The maximum no. of electrons in an orbital is given by the formula 2n 2 . Eg: 1) Hydrogen atom has 1 electron in K-shell 2)Helium atom has 2 electrons in K-shell 3)Oxygen atom has 8 electrons (2 in K-shell, 6 in L-shell)

25. NUCLEAR FORCES: There are four types of forces in nature 1)Strong nuclear force 2)Electromagnetic force 3)Weak nuclear force 4)Gravitational force Strong nuclear force -Short range force(~10 -15 m )

26. Energy level diagram of a particle in a nucleus A)Particle with no charge B)Particle with positive charge U(r) - Potential energy r - distance from centre of nucleus B- barrier height R- Nuclear radius

27. Energy level diagram of decay of 60 Co 27 nucleus. Nuclear energy levels: 60 Co 27 nucleus firstly emits β - particle with emission of photons 60 Co 27 60 Ni 28 + e - + ν Due to Nuclear transformation an neutron disintegrates in to a proton,an electron & a neutrino

28. Thank you… Vinay Desai M.Sc Radiation Physics Radiation Physics Department KIDWAI MEMORIAL INSTITUTE OF ONCOLOGY Bengaluru E-mail:- [email protected]