Unit2 Presentation

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Unit 2 - Nuclear Chemistry and the Structure of the Atom

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Unit2 Presentation

  1. 1. Atomic Structure Early Atomic Theory Democritus (400 BC) suggested that the world was made of 2 thing empty space and tiny, invisible particles called atoms .
  2. 2. Aristotle Aristotle believed matter was composed of 4 basic elements, water, fire, earth, and air. He was a well respected so everyone discounted Democritus’s atomic theory and believed Aristotle’s theory without question.
  3. 3. John Dalton (early 1800s) using the experimental observations of others, Lavoisier, and Proust proposed his own atomic theory.
  4. 4. Dalton’s Atomic Theory <ul><li>All elements are composed of tiny </li></ul><ul><li>particles called atoms. </li></ul><ul><li>Atoms of the same element are identical . The atoms of any one element are different from those of another element. </li></ul><ul><li>Isotopes of an element are not identical . </li></ul>
  5. 5. <ul><li>Atoms cannot be divided , </li></ul><ul><li>created, or destroyed. Atoms are divisible by a nuclear charge. </li></ul><ul><li>Atoms of different elements can combine with one another in simple, whole number ratios to form compounds . </li></ul><ul><li>Chemical reactions occur when atoms are joined , separated , or rearranged . However , atoms of one element are not changed into atoms of another element. </li></ul>
  6. 6. Early Research on Atomic Particles Thomson (1867) used a cathode ray tube to prove the existence of negatively charged subatomic particles called electrons .
  7. 7. Cathode Rays and the Electron Thomson investigated cathode rays using a devise called a cathode ray tube or CRT. Currents of electricity were pumped into vacuum tubes causing the tube to fluoresce. Thomason used magnets to determine the identity of the particles making up the rays. He found them to be tiny, negatively charged particles.
  8. 9. Robert Millikan (1909) used his oil drop experiment to prove that the charge on of an electron is -1.
  9. 11. Rutherford Using his gold foil experiment, Rutherford proposed that the atoms is composed of a lot of empty space with, a small dense, positively charged nucleus.
  10. 13. Thomson – used a modified cathode ray tube to prove the existence of a positively charged subatomic particle called a proton . It has equal but opposite charge to the electron (+1) and a mass 1840 times heavier than an electron.
  11. 14. Chadwick Chadwick used a device to prove that the nucleus contained neutral particles of the same mass as the proton called neutron .
  12. 16. The atomic number is the number of protons in an atom. This number is unique for all elements and the atomic number is used to identify each element. Since atoms are electrically neutral, THE NUMBER OF PROTONS EQUALS THE NUMBER OF ELECTRONS.
  13. 17. Isotopes are atoms of the same element that differ in the number of neutrons in the nucleus. Isotopes of the same element have the different chemical properties. The number of neutrons determines the particular isotope of the element.
  14. 19. Subatomic Particle Chart Particle Charge Relative mass Location Electron -1 1/1840 Outside Nucleus Proton +1 1 Nucleus Neutron 0 1 Nucleus
  15. 20. The mass number is the total number of protons and neutrons in an atom. Mass Number = protons + neutrons Mass number – protons = Number of Neutrons
  16. 23. Calculating Atomic Mass Atomic Mass is the weighted average of the masses of the isotopes of that element. A weighted average takes into consideration both the mass and the abundance of each isotope. The correct unit for atomic mass is amu.
  17. 24. To calculate relative atomic mass, multiply the mass number of each isotope by its percent abundance changed to a decimal and total. ( Mass #)(isotope’s relative abundance) + (Mass #)(another isotopes Rel. abund) = Relative atomic mass of the element
  18. 25. EXAMPLE: 32.09384 amu Symbol Abundance Calculation Average Atomic Mass 32 S 95.00% 32 X 0.95 30.4 33 S 0.76% 33 X 0.0076 0.2508 34 S 4.22% 34 X 0.0422 1.4348 36 S 0.014% 36 X 0.00014 0.00504
  19. 26. Example 2: Neon has 2 isotopes, Ne-20 with an abundance of 90% and Ne-22 with an abundance of 10%. Calculate the average atomic mass of neon.
  20. 27. Example 3: Carbon occurs in nature as a mixture of atoms of which 98.89% have a mass of 12.00 u and 1.11% have a mass of 13.00335 u. Calculate the atomic mass of carbon.
  21. 28. Radioactivity Radioactivity was 1 st discovered by Antoine Becquerel, when a photographic plate never exposed to Sunlight in his lab had become exposed. The only possible culprit was a nearby uranium salt sitting on the bench top.
  22. 29. History of Radioactivity The term radioactivity was 1 st used by Marie Curie in 1898. Curie and her husband, Pierre, found that radioactive particles were emitted as either electrically negative which were called beta particles ( ß) or positive particles called alpha particles ( α ).
  23. 30. Nuclear Chemistry Nuclear reactions are reactions that affect the nucleus of the atom.
  24. 31. Radioactivity is the phenomenon of radiation (particles and/or energy) being ejected spontaneously by an unstable nucleus until it reaches a more stable arrangement.
  25. 32. Nuclear Stability is determined by the ratio of protons to neutrons in the nucleus.
  26. 33. There are forces in the nucleus that oppose each other, the &quot; Strong &quot; force holding Protons and Neutrons to each other and the electrostatic force of protons repelling other protons.  Under certain arrangements of protons and neutrons the electrostatic force can cause instability in the nucleus causing it to decay.   It will continue to decay until it reaches a stable combination.
  27. 34. Radioactive decay is the process by which the unstable nuclei lose mass and/or energy by emitting radiation. Eventually unstable nuclei achieve a more stable state when they are transformed into atoms of a different element .
  28. 35. This graph shows the stable nuclei in red .  There are several things to notice:
  29. 36. <ul><li>There are no stable nuclei with an atomic number higher than 83 or a  neutron number higher than 126 . </li></ul>
  30. 37. <ul><li>The more protons in the nuclei, the more neutrons are needed for stability.  Notice how the stability band pulls away from the P=N line. </li></ul>
  31. 38. <ul><li>Stability is favored by even numbers of protons and even numbers of neutrons. 168 of the stable nuclei are even-even while only 4 of the stable nuclei are odd-odd. </li></ul>
  32. 39. <ul><li>Types of Radioactive Decay  </li></ul><ul><ul><li>When unstable nuclei decay, the reactions generally involve the emission of a particle and or energy . </li></ul></ul><ul><li>For each type of decay, the equation is balanced with regard to atomic number and atomic mass. In other words, the total atomic number before and after the reaction are equal.  And the total atomic mass before and after the reaction are also equal . </li></ul>
  33. 40. Transmutation When particles break down in the nucleus in an atom of an element (radioactive decay), the element changes into another element. This is called transmutation .
  34. 41. TYPES OF RADIATION Gamma emission is the high energy electromagnetic radiation given off in most nuclear reactions. GAMMA RAYS ARE NOT MATTER, THEY ARE ENERGY. Therefore, they are not involved in balancing the nuclear equation. They are very damaging and difficult to shield against.
  35. 42. Gamma Emission ( λ ) Generally accompanies other radioactive radiation because it is the energy lost from settling within the nucleus after a change.
  36. 43. Alpha Emission ( α ) Happens when the atomic number is greater than 83 The 2 p+ 2n ( ) loss brings the atom down and to the left toward the belt of stable nuclei.
  37. 44. Alpha Particle Emission ( α ) Uranium - 238 Thorium - 238 Alpha Particle ( α )
  38. 45. BETA EMISSION A beta particle (a high energy electron, charge of -1) is generated in the nucleus as a neutron is converted into a proton. Carbon - 14 Nitrogen - 14 Beta Particle
  39. 46. Beta Particle Emission ( β ) Happens to nuclei with high neutron:proton ratio A neutron becomes a proton causing a shift down and to the right on the stability graph   
  40. 47. Positron Emission A positron is an antimatter particle that is like an electron but has a positive charge. A positron is generated as a proton is converted to a neutron. Fluorine - 18 Oxygen - 18 Positron
  41. 48. Positron Emission Happens to nuclei with a low neutron:proton ratio A proton becomes a neutron causing a shift up and to the left                                               
  42. 49. This graph shows all the trends of decay and the band of stable nuclei
  43. 50. Penetrating Power of Radiation
  44. 51. Nuclear Fission Fission is the breaking apart of a very heavy nucleus into parts.
  45. 52. A nuclear reactor is a device for controlling nuclear fission to produce energy for practical use. The main fuels for nuclear reactors are U-235 and plutonium-239 .
  46. 53. Fusion is the combining of 2 small nuclei into 1 larger one. Fusion of hydrogen into helium occurs in the sun. Fusions reactions should produce much more energy than fission and use much more accessible fuels. However, currently many problems exist in fusion reactions such as the extremely high temperature needed for the reaction.
  47. 55. Decay Series for Uranium - 238
  48. 56. Half-Life Half-Life (t 1/2 ) is the time required for half of the atoms of a radioisotope to emit radiation and to decay to products .
  49. 57. Half-Life Example It takes 4.5 X 10 9 years for one half of a sample of uranium-238 to decay to lead-206. Therefore, it would take another 4.5 X 10 9 years for one half of the remaining uranium to decay, et cetera, et cetera, et cetera. 100g 50g 50g 1 half-life 25g 25g 12.5g 12.5g 6.25g 2nd half-life 3rd half-life 4th half life
  50. 58. How many atoms of a 2.97g. sample of molybdenum-91 would remain after 62 min. if the half-life of molybdenum-91 is 15.49 min.? How many ½ lives is this? Answer = 0.19 g 4 half-lives # Half-Lives Time Spent (min) Amount Remaining (g) 0 0 2.97 1 15.49 1.49 2 30.98 0.74 3 46.47 0.37 4 61.96 0.19
  51. 59. Radiation Detection Film badges are used to monitor the amount of radiation exposure people have received.
  52. 60. Geiger Counter Instrument that detects radiation by measuring current produced by gas particles ionized by radioactivity
  53. 62. Scintillation Counter Instrument that converts light to an electric signal for detecting radiation.
  54. 63. Uses for Nuclear Radiation Since the physical and chemical properties of radioisotopes of an element are the same as stable ones, many uses for radioactive nuclides are possible.
  55. 64. In medicine radioactive nuclides are used to destroy cancer cells and as tracers to tract substances through the body or identify cancer and other diseases. Cobalt - 60 Radioactive Tracer
  56. 66. In agriculture, radioactive nuclides are used as tracers in fertilizer to determine the effectiveness or to prolong shelf life of food by irradiating to destroy microorganisms.
  57. 67. In dating radioactive nuclides are used to determine the age of objects. Example: Carbon -14 is used to date organic materials.
  58. 68. In energy production, currently nuclear fission is used to create energy. Example: Comanche Peak nuclear power plant in Glen Rose produces energy that is used by TXU.
  59. 69. Nuclear Waste Nuclear fission produces radioactive wastes that must be contained and stored on-site (temporary) or disposed of (permanent).

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