basic_nuclear_physics, Class for CBRN.pptx

Learning basics, radioactivity

1. BASIC NUCLEAR PHYSICS SI-DIWAKER YADAV 1 2 BN NDRF

2. The objective of this chapter is to introduce the concept of atom, atomic structure, elements and isotopes, Isobar, Isotones, Radioactivity, Half Life. OBJECTIVE

3. Structure of an Atom

4. HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA ( greek for indivisible )

5. HISTORY OF THE ATOM 1808 John Dalton suggested that all matter was made up of tiny spheres that were able to bounce around with perfect elasticity and called them ATOMS

6. HISTORY OF THE ATOM 1898 Joseph John Thompson found that atoms could sometimes eject a far smaller negative particle which he called an ELECTRON

7. HISTORY OF THE ATOM Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge 1904 like plums surrounded by pudding. PLUM PUDDING MODEL

8. HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them passed through. About 1 in 10,000 hit

9. HISTORY OF THE ATOM gold foil helium nuclei They found that while most of the helium nuclei passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back. helium nuclei

10. HISTORY OF THE ATOM Rutherford’s new evidence allowed him to propose a more detailed model with a central nucleus . He suggested that the positive charge was all in a central nucleus. With this holding the electrons in place by electrical attraction However, this was not the end of the story.

11. HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits . Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

12. Bohr’s Atom electrons in orbits nucleus

13. The Atom

14. Nucleus Protons and neutrons together form the nucleus of the atom. The nucleus determines the identity of the element and its atomic mass.

15. Protons Protons are positively charged particles found inside the nucleus of an atom. Each element has a unique atomic number (a unique number of protons).

16. Neutrons Neutrons are the other particle found in the nucleus of an atom. Unlike protons and electrons, however, neutrons carry no electrical charge and are thus "neutral."

17. Electrons Electrons are negatively charged particles that surround the nucleus in “orbits” similar to moons orbiting a planet.

18. Atom positively charged (+) protons, uncharged neutrons and negatively charged (-) electrons The atom is composed of:

19. Atom Thomson’s Model Rutherford’s Model

20. Atom Bohr’s Model

21. Particle Symbol Mass (kg) Energy (MeV) Charge Proton p 1.672E-27 938.2 +1 Neutron n 1.675E-27 939.2 Electron e 0.911E-30 0.511 -1 Summary of the Atom

22. Xy Z A Xy = element symbol A = mass number (neutron + protons) Z = atomic number (protons) Symbol of Element

23. Elements An element is a substance made up of a single type of atom For an uncharged atom, the number of electrons equals the number of protons.

24. In 1869, Russian chemist Dmitri Mendeleev first described an arrangement of the chemical elements now known as the periodic table. The periodic table displays all chemical elements systematically in order of increasing atomic number (the number of protons in the nucleus). Periodic Table of the Elements

25. Rare Earth Elements Actinide Series Lanthanide Series Periodic Table of the Elements

26. Atomic Number (No. of Protons) Element Symbol 1 Hydrogen H 2 Helium He 6 Carbon C 7 Nitrogen N 8 Oxygen O 11 Sodium Na 14 Silicon Si 15 Phosphorous P 26 Iron Fe 27 Cobalt Co 38 Strontium Sr 47 Silver Ag 53 Iodine I 55 Cesium Cs 79 Gold Au 82 Lead Pb 88 Radium Ra 92 Uranium U 94 Plutonium Pu

27. Atomic Number (Z) The number of protons (p) in the nucleus of an atom S 35 16

28. Atomic Mass Number (A) Sum of the number of protons and neutrons (p + n) in the nucleus of an atom P 32 15 I 125 53

29. ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shell a maximum of 2 electrons second shell a maximum of 8 electrons third shell a maximum of 1 8 electrons

30. ELECTRONIC CONFIGURATION With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example; N Nitrogen 7 14 2 in 1 st shell 5 in 2 nd shell configuration = 2 , 5 2 + 5 = 7

31. ELECTRONIC CONFIGURATION Write the electronic configuration for the following elements; Ca O Cl Si Na 20 40 11 23 8 17 16 35 14 28 B 11 5 a) b) c) d) e) f) 2,8,8,2 2,8,1 2,8,7 2,8,4 2,3 2,6

32. SUMMARY The Atomic Number of an atom = number of protons in the nucleus. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. The number of Protons = Number of Electrons. Electrons orbit the nucleus in shells . Each shell can only carry a set number of electrons.

33. Atomic Mass Unit (amu) Where 1 amu is approximately equal to 1.6605 x 10 -24 grams

34. Atomic Mass Unit (amu) The atomic mass of the proton and the neutron is approximately: Proton = 1.6726 x 10 -24 grams = 1.0073 amu Neutron = 1.6749 x 10 -24 grams = 1.0087 amu Thus, the neutron is just a little heavier than the proton.

35. Atomic Mass Unit (amu) The difference in the mass of the neutron and the proton can be understood if we assume that the neutron is merely a proton combined with an electron forming a neutral particle slightly more massive than a proton alone.

36. Atomic Mass Unit (amu) The atomic mass of the electron is approximately: Electron = 9.1094 x 10 -28 grams = 0.00055 amu Thus, the electron has a much smaller mass than either the proton or the neutron, 1837 times smaller or about 2000 times smaller.

37. Isotopes Atoms of an element that have a different number of neutrons in the nucleus are called isotopes of each other. Xy Z A Xy = element symbol A = atomic mass (neutron + protons) Z = atomic number (protons) isotope notation typically written as:

38. Isotopes The number of protons and electrons remain the same. But the number of neutrons varies.

39. Isobars Nuclides that have same mass number with different atomic and neutron number

40. Isotones Nuclides that have same neutron number with different atomic and mass numbers

41. Stable Nuclides long range electrostatic forces short range nuclear forces p p n Line of stability Unstable isotopes undergo disintegration by emitting radiations to achieve stability. This disintegration phenomenon is called ‘Radioactivity’

42. Summary

43. RADIOACTIVITY

44. RADIOACTIVITY NUCLEAR FORCES : Nucleus contains neutrons & positively charged protons. It is expected that closely packed, positively charged protons would tend to move apart. But this does not happen. This is due to Nuclear Force , which is even greater than the force of repulsion between positive charges. The nuclear force is also called ‘ Binding Force ’.

45. STABLE NUCLIDES : If the binding force is greater than the electrostatic force, the nucleus will remain intact and unchanging. All together about 280 stable nuclides have been identified. For light nuclei, the numbers of neutrons and protons are approximately equal, whereas for heavy nuclei there might be about 1½ times as many neutrons as protons. In the higher atomic numbered elements electrostatic repulsive forces becoming dominant as the binding nuclear forces are becoming saturated.

46. UNSTABLE NUCLIDES An unstable nucleus is one, which has too many or too few neutrons for its particular number of protons. Unstable nucleus undergo disintegration by emitting particles to achieve stability. This disintegration phenomenon is called Radioactivity and the unstable nuclide is called radionuclide.

47. RADIOACTIVE TRANSFORMATION The process, whereby a radioactive nucleus emits particles, is called ‘ Radioactive Decay ’ or ‘Radioactive Transformation ’. A unstable atom emits Alpha, Beta, Neutron, positron, Gamma etc.

48. Ionisation Radiation  which passes through matter creates ion pairs  along its path. It takes energy to ionise atoms , so the radiation will give up energy as it travels in the matter. When the matter is human tissue this energy deposition is what causes biological damage. These energies are usually expressed in units known as electron volts  ( eV ). Alpha particles produce about 50,000 ion pairs per centimeter as they travel through air, whereas beta particles produce about 100 ion pairs in the same distance . The average energy loss per ion pair  formed in air is 33.5 eV

49. Activity In radiation protection, it is more important to know how much radiation is being emitted rather than how many radioactive atoms remain in the sample. Activity is define as the number of disintegrations of the radioactive material taking place in a unit time (usually seconds). SI unit is the Becquerel (Bq) 1Bq = 1disintegration per second 1 Ci = 3.7 x 10 10 Bq

50. HALF LIFE TIME TAKEN FOR ACTIVITY OF A RADIOACTIVE SUBSTANCE TO DECAY TO HALF ITS ORIGINAL VALUE VARIES FROM LESS THAN A MILLIONTH OF A SECOND TO MILLIONS OF YEARS ENABLES TO PREDICT THE REMAINING ACTIVITY

51. Fractions Multiples Symbol Prefix Factor Symbol Prefix Factor d deci 10 -1 k kilo 10 3 c centi 10 -2 M mega 10 6 m milli 10 -3 G giga 10 9  micro 10 -6 T tera 10 12 n nano 10 -9 p pico 10 -12

52. THANK YOU