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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 • Accidents • 3-mile Island (USA), Chernobyl (USSR), Japan • Potassium Iodide protection

  2. 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 ignored 1/1836 =0.054% (considered 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

  3. 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

  4. Protons & Neutrons • Protons at core of Element, with positive charge • Positive charge balances negative electron charge • Element Atomic Number = number of Protons in that element • Total number protons = nuclear charge = “Z” • Electron has least mass • One electron = 0.910938291*10-27 grams • Proton has mass of 1836 times that of electron • One proton = 1.6726218*19-24 grams • Neutron slightly heavier than Proton … but no charge • One neutron = 1.6749286*10-24 grams • Isotopes • Same Element and atomic number, different number of neutrons • Approximately the same number of neutrons as protons • Large variations between elements • Atomic Weight (most often used) • Weighted average of isotope masses • What’s on the periodic chart

  5. Isotopes • Same element can have different mass • Same # of protons • Protons determine # of electrons for bonding • Neutral element has equal # protons and electrons • variable # of neutrons • No charge but mass about same as proton • No influence in chemical reactions • “goes along for the ride” • Show slide of isotopes • Thousands of isotopes for ≈110 elements

  6. 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.

  7. 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

  8. Conservation of Energy • Chemical change • Reactions cause change, but not total amount of energy • Total energy the same before & after reaction • Different forms of energy can be inter-converted • Potential energy into kinetic (roller coaster) • Chemical Bonds broken generate heat • Burning wood, gasoline • Reaction simply rearranges the relationships • Conversion may happen between solid, liquids, gases • Quantity of element atoms will be the same • Energy insignificant compared to mass • Generally ignored in chemical formulas

  9. Conservation of Mass (among other things) • Conservation rules: • Sum of Products = Sum of Ingredients • Numerous conservation rule examples … • Conservation of Mass, • Conservation of Energy, • Conservation of Momentum, • Conservation of Angular Momentum • Chemical change cannot create or destroy mass • Reactions cause change in form but not total amount • Total mass the same before and after reaction • Mass is invariant to chemical reactions

  10. Conservation of Mass (and other things) Einstein would disagree … • Mass is convertible to energy and vice versa • Applies particularly to nuclear reactions • Escaping radiation equivalent to some mass loss • Conversion ratio is enormous, E=mC2 • LOT of energy equivalent to small amount of mass • Practical impact in non-nuclear chemistry • Ignore radiation for mass conservation • Ignore heat for mass conservation

  11. 11.1 Nuclear Reactions • 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:

  12. Writing Nuclear Reactions • “stacked” numbers difficult to write 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

  13. Cathode Rays, Crookes 1895Early 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.

  14. 60 years of breakthroughs • 1884 - Chemistry innovations, 56 theses, Arrhenius • 1888 - Proton Discovery, Goldstein • 1895 – X-Ray discovery, Roentgen • 1896 - Radioactivity Discovery, Baqerel • 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

  15. 11.2 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.

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

  17. E=mc2, so light energy has mass Although small, mass of light is an important consequence

  18. Space distortion due to massGeneral Relativity predicted light deflection by mass in 1916, proven by experiment in 1919

  19. Gravity Well or “Black Hole”space is so distorted that light cannot pass nearby, a consequence of general relativity

  20. 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

  21. 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.

  22. Alpha rays move at ~0.1c (”c” is speed of light), 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.

  23. 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

  24. 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

  25. 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

  26. 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

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

  28. 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”

  29. 11.3 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.

  30. 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.

  31. Radioactivity • Emissions from Pitchblende (uranium ore) • Found to expose photographic film • 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 • Gama (γ), radiation similar to X-Ray • Very penetrating, used for imaging

  32. 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

  33. 11.4 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.

  34. 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

  35. Unstable nucleus emits a helium nuclei Alpha Decay a common event, new element with lower atomic number (by two) is formed

  36. 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.

  37. Proton Emissionloss of positive charge reduces atomic number by 1

  38. Neutron Emission

  39. Beta ParticleUnstable isotope emits electron having negative charge, nucleus of parent changes charge in the other direction, creating a new element. Example is Carbon-14 used for dating historical objects. Note another particle with zero mass and charge, the “Neutrino” … more about that later

  40. Electron capture, symbolized E.C., is a process in which the nucleus captures an inner-shell electron from the surrounding electron cloud, thereby converting a proton into a neutron. • The conversion of mercury- 197 into gold-197 is an example of electron capture.

  41. Electron CaptureInner orbit electron can be “snagged” by nucleus

  42. 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.

  43. What about anti-matter? • 1928 prediction by Paul Dirac that wave equations allow for negative matter • Anti-particles are the same as conventional 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 • Just like they say in Star Trek !

  44. Anti-Particles, Anti-Matter • 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

  45. 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.

  46. Positron EmissionUnstable nuclei can also emit a positive electron, a form of “anti-matter” which turns into energy upon meeting it’s opposite. Mass and charge must balance

  47. Emission of g rays causes no change in mass or atomic number. • g emission usually accompanies emission of other rays but it is often omitted from nuclear equations. • Their penetrating power makes them both dangerous to humans and useful in medical applications.

  48. 11.5 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.

  49. 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.

  50. Carbon-14 Dating • Carbon14 discovered in 1940 at UC • 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