Ch. 19: Radioactivity and Nuclear Chemistry

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

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

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.

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

II. Review of Atomic Symbols

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.

II. Shedding Helium

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.

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.

II. Emitting an Electron

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.

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.

II. Antiparticles of Electrons!!

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

II. Radioactive Decay Summary

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.

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.

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).

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.

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.

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.

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.

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.

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

IV. Rate of Decay and Amount are Interchangeable

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.

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?

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.

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.

V. Nuclear Chain Reaction

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.

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.

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.

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.

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.

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.

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.

VII. Sources of Radiation

VII. Results of Radiation Exposure

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

Ch. 19: Radioactivity and Nuclear Chemistry