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Ch. 19: Radioactivity and Nuclear Chemistry

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  1. Ch. 19: Radioactivity and Nuclear Chemistry Dr. Namphol Sinkaset Chem 201: General Chemistry II

  2. I. Chapter Outline • Introduction • Types of Radioactivity • The Valley of Stability • Radiometric Dating • Nuclear Fission • Nuclear Fusion • Radiation and Life

  3. I. Introduction • Antoine-Henri Becquerel discovered radioactivity when he placed some rock crystals on a photographic plate. • He called the rays that were emitted uranic rays because they came from uranium in the crystals. • Marie Curie changed the name to radioactivity when she discovered polonium and radium.

  4. II. Types of Radioactivity • Ernest Rutherford and others worked on figuring out what radioactivity was. • Discovered that radioactive emissions were produced from unstable nuclei. • Several types of radioactivity • alpha (α) decay • beta (β) decay • gamma (g) ray emission • positron emission • electron capture

  5. II. Review of Atomic Symbols

  6. 0 p n e 1 1 -1 1 0 II. Subatomic Particles • The term nuclide is used to refer to a particular isotope of an element. • Each nuclide is composed of subatomic particles. • Each subatomic particle has its own representation in nuclear chemistry.

  7. II. Shedding Helium

  8. II. Nuclear Equations • In a nuclear reaction, elements change their identity. • Nuclear equations are balanced by ensuring the sum of mass numbers and the sum of atomic numbers on both sides are equal.

  9. II. α Partcles – Dangerous? • Alpha particles are the most massive particles emitted by nuclei. • They have the potential to interact with and damage other molecules. • Alpha radiation has the highest ionizing power, but it has the lowest penetrating power.

  10. II. Emitting an Electron

  11. II. Dangers of Beta Particles • Beta particles are less massive than alpha particles, so they have less ionizing power. • However, they have greater penetrating power. Sheet of metal or thick block of wood needed to stop them.

  12. II. Gamma Ray Emission • This type of radiation involves emission of high-energy photons, not particles. • Gamma rays have no mass and no charge as they are a type of EM radiation. • Gamma rays can be emitted along with other types of radiation. • Gamma rays have low ionizing power, but very high penetrating power.

  13. II. Antiparticles of Electrons!!

  14. II. Electron Capture • Instead of emitting particles, a nucleus can pull in an e- from an inner orbital. • When an e- combines with a proton in the nucleus, a neutron is formed. • proton + electron  neutron

  15. II. Radioactive Decay Summary

  16. II. Sample Problems • Write a nuclear equation for the positron emission of sodium-22. • Write a nuclear equation for electron capture in krypton-76. • Potassium-40 decays into argon-40. Identify the type of radioactive decay.

  17. III. Why Is There Radioactivity? • When a nuclide undergoes radioactive decay, it becomes more stable. • The strong force binds protons and neutrons together, but it only works at very short distances. • Stability of nucleus is a balance between +/+ repulsions and the strong force attraction.

  18. III. Importance of Neutrons • Neutrons are key to nuclei stability because they increase strong force attractions, but lack charge repulsion. • However, since neutrons occupy energy levels like e-, cannot just stuff nucleus with neutrons. • Nuclear stability is indicated by the ratio of neutrons to protons (N/Z).

  19. III. The Valley of Stability • For lighter elements, N/Z for stable isotopes is about 1. • For Z > 20, stability requires higher N/Z. • No stable isotopes above Z = 83. • Thus, nuclides decay to get back to the valley of stability.

  20. III. Magic Numbers • Nucleons occupy energy levels in the nucleus, so certain numbers of nucleons are stable. • N or Z = 2, 8, 20, 28, 50, 82, and N = 126 are uniquely stable and are called magic numbers.

  21. III. Journey to Valley of Stability • Atoms w/ Z > 83 undergo decay in one or more steps to become stable. • The successive decays to become stable are known as a decay series.

  22. IV. Radioactivity is Everywhere • Everything around us contains at least some nuclides which are radioactive. • Radioactivity is found in the ground, in our food, in our air. • Radioactivity is in our environment because of some long decay times, and the constant production of radioactive nuclides through various decay series.

  23. IV. Radioactivity is 1st Order • All radioactive nuclides follow 1st order kinetics. • Thus, ln Nt/N0 = -kt. • Since decay is 1st order, half lives are independent of initial concentration.

  24. IV. Sample Problem • How long would it take for a 1.35-mg sample of Pu-236 to decay to 0.100 mg?

  25. IV. Rate of Decay and Amount are Interchangeable

  26. IV. Radiocarbon Dating • Radioactive C-14 is continuously taken up by living organisms, so the amount is in equilibrium with the amount in the atmosphere. • When the organism dies, it no longer takes in C-14. The C-14 continuously decays in the remains. • Age can be determined by comparing rate of decay in remains to rate of decay in atmosphere.

  27. IV. Sample Problem • An ancient scroll is claimed to have originated from Greek scholars in about 500 B.C. A measure of its C-14 decay rate gives a value that is 89% of that found in living organisms. How old is the scroll? Could it be authentic?

  28. V. Making New Elements • Enrico Fermi attempted to synthesize a new element by bombarding U-238 with neutrons. • He detected beta particles, but never confirmed the chemical products.

  29. V. Nuclear Fission • Meitner, Strassmann, and Hahn repeated Fermi’s experiment. • They discovered that elements lighter than uranium were produced w/ a lot of energy.

  30. V. Nuclear Chain Reaction

  31. V. Source of Energy in Fission • U-235 + n  Ba-140 + Kr-93 + 3n • If we look at exact masses, we find that mass of products is 235.86769 amu and mass of reactants is 236.05258 amu. • Mass is not conserved!! • In nuclear reactions, mass can be converted into energy via E = mc2.

  32. V. The Mass Defect • All stable nuclei have masses less than their components which is known as the mass defect. • When the mass defect is used in E = mc2, the nuclear binding energy is calculated. • The nuclear binding energy is the energy needed to break up a nucleus into its component nucleons.

  33. V. Calculating Binding Energies • A useful conversion between mass and energy is 1 amu = 931.5 MeV. Note that 1 MeV = 1.602 x 10-13 J. • The mass defect of a helium nucleus is 0.03038 amu, so its binding energy is 28.30 MeV.

  34. V. Comparing Nuclei Stability • In order to see which nuclei are more stable than others, the binding energy per nucleon is calculated. • This is simply the binding energy divided by the number of nucleons in the nuclide.

  35. VI. Nuclear Fusion • Smaller nuclides can combine into more stable nuclides in a process called fusion. • Fusion is the energy source of the sun and used in hydrogen bombs. • High temps are needed to overcome the +/+ repulsions.

  36. VII. Radiation Risks • There are 3 categories of radiation effects. • Acute radiation damage: large amounts of radiation in short time. Immune and intestinal cells most susceptible. • Increased cancer risk: low dose over time. Damage occurs to DNA. • Genetic defects: high radiation exposure to reproductive cells causing problems in offspring. Not seen in humans, even Hiroshima survivors.

  37. VII. Measuring Exposure • There are several ways to measure exposure to radiation. • curie (Ci): exposure to 3.7 x 1010 decay events per second. • gray (Gy): exposure to 1 J/kg body tissue. Also have the rad (radiation absorbed dose) which is 0.01 J/kg body tissue. • rem (roentgen equiv. man): multiplication of rads by the biological effectiveness factor, which depends on the type of radiation.

  38. VII. Sources of Radiation

  39. VII. Results of Radiation Exposure

  40. VII. Applications of Radioactivity • Medicine • Use of radiotracers to track movement of compound or mixture in body. I-131 for thyroid, labeled antibodies to locate infection, P-32 for cancer. • Gamma rays to kill cancer cells. • Kill microorganisms • Sterilize medical devices. • Kill bacteria and parasites in food. • Sterilize harmful insects