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Nuclear Chemistry

Nuclear Chemistry. Brief history of nuclear related discoveries Electron, proton, neutron Nuclear transformations Natural radioactivity Half Life, carbon dating Nuclear chemistry equations Chain reaction, atom bomb Applications Nuclear reactors Radioisotopes Personl Exposure

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Nuclear Chemistry

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  1. Nuclear Chemistry • Brief history of nuclear related discoveries • Electron, proton, neutron • Nuclear transformations • Natural radioactivity • Half Life, carbon dating • Nuclear chemistry equations • Chain reaction, atom bomb • Applications • Nuclear reactors • Radioisotopes • Personl Exposure • Radon, other natural sources • 3-mile Island (USA), Chernobyl (USSR), Japan • Potassium Iodide protection

  2. Atomic versus NuclearWhat’s the Difference? • Atomic properties are those of an atom • Chemical reactions (gain/loss of electrons) • Emission of light (electron orbit jumps) • Bonds between elements (covalent, ionic) • Nuclear properties are within the atom • Construction of nucleus (electrons, protons) • Radioactive disintegration • Fission/fusion of nuclei, element conversion

  3. Atomic number • Atomic Number (and element number) • Number protons = number electrons = “Z” • Atomic Mass Number • Total Number of Protons defining the element = Z • Total number of Neutrons in element nucleus = N • Total mass of nucleus = A = Z + N • Electron mass is ignored 1/1836 =0.054% (negligible) • Isotopes • Same atomic number, different number of neutrons • Large variations of isotopes between elements • Isotope significance (e.g. U-235 vs U-238, C-14 vs C-12) • Atomic Weight (most often used) • Weighted average of isotope masses • What’s on the periodic chart

  4. Weighted Atomic Mass (natural values) • Lots of isotopes of Uranium exist, all different masses • Most prevalent isotope is U-238 at 99.283% • Most useful isotope is U-235 at 0.711% • Other isotopes rather rare, relatively insignificant • “Weighted Average” recognizes quantity • Sum of mass contributions = the weighted average value • This is what we see on the periodic charts.

  5. Nuclear reaction: A reaction that changes an atomic nucleus, usually causing the change of one element into another. A chemical reaction never changes the nucleus. Different isotopes of an element have essentially the same behavior in chemical reactions but often have completely different behavior in nuclear reactions. The rate of a nuclear reaction is unaffected by a change in temperature or pressure (within the range found on earth) or by the addition of a catalyst. The nuclear reaction of an atom is essentially the same whether it is in a chemical compound or in an uncombined, elemental form. The energy change accompanying a nuclear reaction can be up to several million times greater than that accompanying a chemical reaction. Nuclear Changes

  6. Nuclear Nomenclature • The atomic number, written below and to the left of the element symbol, gives the number of protons in the nucleus and identifies the element. • The mass number, written above and to the left of the element symbol, gives the total number of nucleons, a general term for both protons (p) and neutrons (n). • The most common isotope of carbon, for example, has 12 nucleons: 6 protons and 6 neutrons:

  7. Writing Nuclear Reactions • “stacked” numbers difficult to write • Possible but difficult using a word processor • Alternative is “front and back” values • Carbon 14 = 6C14 • Front value is atomic number Z • Back value is atomic mass A

  8. 60 years of discoveries • 1884 - Chemistry theories, 56 theses, Arrhenius • 1888 - Proton Discovery, Goldstein • 1895 – X-Ray discovery, Roentgen • 1896 - Radioactivity Discovery, Baquerel • 1897 - Electron Discovery, J.J. Thompson • 1905 - Radioactive Element separation, Curie • 1905 - Equivalence of mass & energy E=mC2, relativity, photoelectric effect, A. Einstein • 1916 - general relativity, proven in 1919, A. Einstein • 1932 - Neutron Discovery, Chadwick • 1933 - Nuclear chain reaction proposed, Szilard • 1938 - Nuclear fission discovered • 1942 - First operational nuclear reactor, Fermi • 1945 - First warfare use of nuclear energy

  9. Cathode Rays (electrons)first demonstrated by Crookes in1895Early investigator of radiation inside electrical discharge tubes, eventually leading to CRT (Television) tubes. He was one of the first to experiment with radioactivity and its ability to make certain minerals glow. He also invented the ”radiometer” still in use as an educational toy.

  10. ELECTRON - J.J. Thompson 1897found a new particle “boiling off” a heated filament which had <1/1000 mass of hydrogen. It had a negative charge by its magnetic and/or electrostatic deflection. Using similar apparatus he discovered isotopes of the same element with different mass, which led to science of mass spectrometry

  11. PROTON, Goldstein in1888Used high voltage to ionize gases, accelerating particles through holes in cathode, causing “canal rays” (trails looked like canals). Particles were positive. Hydrogen particles later identified as protons by Rutherford in 1919

  12. NEUTRON, Chadwick 1932observed a new form of penetrating radiation, which had no charge (not protons or electrons)

  13. NEUTRINOPredicted by Wolfgang Pauli in 1930, based on conservation discrepancies. the “little neutron” Indirectly observed in 1942 and1946 via interactions with other particles, directly observed in 1972 “bubble chamber”

  14. Discovery and Nature of Radioactivity • In 1896, the French physicist Henri Becquerel noticed a uranium-containing mineral exposed a photographic plate that had been wrapped in paper. • Marie and Pierre Curie investigated this new phenomenon, which they termed radioactivity: The spontaneous emission of radiation from a nucleus. • Ernest Rutherford established that there were at least two types of radiation, which he named alphaand beta. Shortly thereafter, a third type of radiation was found and named for the third Greek letter, gamma.

  15. Baquerel – observed radioactivity 1896Photographic plate accidentally exposed by Uranium Baquerel is SI unit of radiation, Bq = disintegrations/sec1 Curie = radiation from 1 gram of Radium = 3.7*10^10 Bq

  16. Radioactivity Emissions from Pitchblende (uranium ore) Found to expose photographic film We will measure pitchblende today 3 common types of nuclear radiation Alpha (α), helium nuclei particle, 2He4 = 2p+2n Very strong but not very penetrating Beta (β), an electron particle, -1e0 Mildly penetrating, stopped by thick paper Gamma (γ), radiation similar to X-Ray Very penetrating, used for imaging

  17. Stable and Unstable Isotopes • Every element in the periodic table has at least one radioactive isotope, or radioisotope, and more than 3300radioactive isotopes are known. • Their radioactivity is the result of having unstable nuclei. Radiation is emitted when an unstable radioactive nucleus, or radionuclide, spontaneously changes into a more stable one. • There are only 264 stable isotopes among all the elements. • All isotopes of elements with atomic numbers higher than that of bismuth (83) are radioactive.

  18. Radioactivity • Results from unstable elements • Heavy elements formed inside stars • Formed and stable at extreme temperatures • Unstable and disintegrate at earth temperatures • Half-Life • Time it takes for ½ of material to disintegrate • Uranium 238 is 4.5 billion years, same as earth’s age • There was twice as much uranium when earth formed • A non-linear scale (e.g ½ of ½, etc) • A natural nuclear reactor happened in Okla, Africa

  19. Marie & Pierre Curie, 1905Separated tons of mineral pitchblende to discover and isolate Radium, and polonium (named for Poland). The standard measure for radioactivity is the Curie = Ci

  20. For elements in the first few rows of the periodic table, stability is associated with a roughly equal number of neutrons and protons. • As elements get heavier, the number of neutrons relative to protons in stable nuclei increases. • Lead-208, for example, the most abundant stable isotope of lead, has 126 neutrons and 82 protons in its nuclei.

  21. Alpha rays move about ~0.1c and can be stopped by a few sheets of paper or by the top layer of skin. • Beta rays move at up to 0.9c and have about 100 times the penetrating power of a particles. A block of wood or heavy clothing is necessary to stop b rays. • Gamma rays move at c and have about 1000 times the penetrating power of a rays. A lead block several inches thick is needed to stop g rays.

  22. When passed between two charged plates: • Alpha rays, helium nuclei (He+2 ), bend toward the negative plate because they have a positive charge. • Beta rays, electrons (e- ), bend toward the positive plate because they have a negative charge. • Gamma rays, photons (g), do not bend toward either plate because they have no charge.

  23. Nuclear Decay • Nuclear decay: The spontaneous emission of a particle from an unstable nucleus. • Transmutation: The change of one element into another. • The equation for a nuclear reaction is not balanced in the usual chemical sense because the kinds of atoms are not the same on both sides of the arrow. A nuclear equation is balanced when the number of nucleons and the sums of the charges are the same on both sides.

  24. Nuclear Decay • Decrease in atomic number • Loss of alpha particle with positive charge • Mass of 2 protons and 2 neutrons • Loss of 2 protons reduces element # by 2 • Uranium 92U238 into Thorium 90Th234 + 2He4 • Increase in atomic number • Loss of electron with negative charge • Charge of nucleus increases by 1 • Next higher element formed • Thorium 90Th234 Proactinium 91Pa234 + -1 e0 • Proactinium 91Pa234 Uranium 92U234 + -1 e0

  25. During alpha emission, the nucleus loses two protons and two neutrons. • Emission of an a particle from an atom of uranium-238 produces an atom of thorium-234.

  26. Beta emission involves the decompositionof a neutron to yield an electron and a proton. • Iodine-131, a radioisotope used in detecting thyroid problems, undergoes nuclear decay by b emission to yield xenon-131.

  27. Anti-Particles • In 1928 Paul Dirac predicted that all particles should have opposites called anti-particles. • The first of these was discovered in 1932 by Carl Anderson. This was an electron with a positive electric charge (+1). This particle is the anti-electron (also called a positron). It is identical in every respect to the electron apart from its electric charge. • When an electron and positron come into contact, they mutually annihilate each other producing a flood of energy in accordance with Einstein's famous equation, E=mC2

  28. Anti-matter • Anti-particles making up atoms are same as conventional particles except for charge • Electron (e-)  Positron (e+), 2008-LNL • Proton (p+)  Anti-Proton (p-), 1955-UCB • Hydrogen  Anti-Hydrogen; 1995-CERN • Matter+Antimatter  pure energy • Antimatter engines in Star Trek ! • Parts of universe could be “anti-matter” • Same properties, light emission, etc. • Matter + Antimatter  pure energy • Something to avoid ….

  29. Positron emission involves the conversion of a proton in the nucleus into a neutron plus an ejected positron. • A positron has the same mass as an electron but a positive charge. • Potassium-40 undergoes positron emission to yield argon-40, this hppens in your body which contains tiny amounts of K-40.

  30. Radioactive Half-Life • Rates of nuclear decay are measured in units of half-life , defined as the amount of time required for one-half of the radioactive sample to decay. • Each passage of a half-life causes the decay of one half of whatever sample remains. The half-life is the same no matter what the size of the sample, the temperature, or any other external conditions.

  31. All nuclear decays follow the same curve, 50% of the sample remains after one half-life, 25% after two half-lives, 12.5% after three half-lives, and so on.

  32. Carbon-14 • Carbon14 discovered 1940 at UC Berkeley • Formed in upper atmosphere from nitrogen • Solar neutrons convert nitrogen to carbon-14 • Carbon-14 becomes carbon dioxide • Carbon-14 dioxide absorbed by plant life • Animal life eats the plant life, absorbs C-14 • Steady state evolves … until object dies • After death, the decay rate reduces C-14

  33. Decay of Carbon-14Neutron turns into proton + electron (beta particle)Mass remains at 14, but carbon becomes nitrogenintensity of electron emission indicates object’s age

  34. 14C is continually produced in the atmosphere and incorporated into life cycles, so 14C amount in living things is constant. Upon death, no more 14C is absorbed, so concentration decreases. Measuring the remaining radioactivity provides an age estimate. The half life of 14C is 5730 years. The method is good for estimating age of objects 500 to 50,000 years old.

  35. Carbon dating methodologycounts/min versus age of sample, assumes C-14 formation at a constant rate

  36. Half Life calculations • Reduction in material is non-linear • original½ at half-life¼ at 2 half-lives • General equation: • n1/n0 = (0.5)^(#of half lives) • n1/n0 = (0.5)^1  50% remaining • n1/n0 = (0.5)^2  25% remaining • n1/n0 = (0.5)^3  12.5% remaining

  37. Half Life calculations • Carbon dating example: (reverse example) • old object has 12.79% normal C14 • Half life of C14 is 5730 years • n1/n0 = (0.5)^(number of half lives) • Take natural logs to eliminate exponents: • Ln(n1/n0)= Ln(0.5)*(t / 5730) • Ln(0.1279/1.00)= Ln(0.5)*(t / 5730) • Ln(0.1279)= -.6931(t / 5730) • -2.0565 = -0.6931 * (t / 5730) • t = (-2.0565)*(5730 years)/-0.69314 = 17,002 years

  38. K-40 is part of your body • K-40 is the largest radiation source in you • Human body has ≈160 grams of Potassium • 0.0117% K-40*160 grams  0.0187g K-40 • This produces 4,400 disintegrations/sec • Emission mostly β (beta ray or electrons) • We will look at a potassium sample today

  39. Radioisotopes used internally for medical applications will have short half-lives so that they decay rapidly and do not remain in the body for prolonged periods (Tc-99, ½=6hrs)

  40. Ionizing Radiation • A large dose of ionizing radiation can destroy living cells, causing death. • A small dose of ionizing radiation may not cause visible symptoms but might lead to a genetic mutation or cancer.

  41. What’s an “MeV”? • Alternative energy unit used in physics • Equals 1 electron passing through 1 volt field • Same as 1.6*10-19Joule • Convenience is simple numbers, fewer zeros • Energy of visible light in eV in diagram below • Nuclear particles & rays >106 times light energy • 1 MeV = 10^6 eV

  42. Light energy in electron volts

  43. Health professionals who work with X rays or other kinds of ionizing radiation protect themselves by surrounding the source with a thick layer of lead or other dense material. Protection is also afforded by controlling the distance between the worker and the radiation source because radiation intensity (I) decreases with the square of the distance from the source. We will demonstrate inverse square law this week in lab. The intensities (I) of radiation at two different distances (d) are given by the equation: I1d12 = I2d22 Radiation Intensity

  44. Inverse Square Law

  45. Inverse square law • Intensity of radiation falls with distance • A non-linear relationship • Twice as far  ¼ the intensity • Works with all forms of point source radiation • Why planet Mercury is so hot, Venus so cold

  46. Inverse Square Example • Radiation intensity is 2250 at 2 meters, what is intensity at 3 meters? • I1 * d12 = I2 * d22 • 2250 * 22 = I2 * 32 • I2 = 2250*4/9 = 1000 • Moving away by 1 meter cuts radiation by more than half

  47. Mass into Energy Enormous ratio between mass & Energy c = 3*10^8 meters/sec c2 = 9*10^16 meters2/sec2 How much energy is that ? 1 gram U235 converts to 3.4*10^8 kcal Hiroshima bomb converted only a few grams

  48. Equivalence of Mass & EnergyAlbert Einstein’s famous equation E=mc2

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