Unit 3 Atomic structure


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Unformatted text preview: ATOMIC STRUCTURE ATOMIC STRUCTURE AND THEORY & NUCLEAR CHEMISTRY CHAPTERS 4 & 25 Democritus (450 BC): Matter is composed of tiny indivisible particles. Dalton’s Atomic Theory (1803): Based on four statements. Each element is composed of atoms. Atoms of the same element are alike and differ from any other element. Atoms can not be created nor destroyed by chemical reactions. Atoms combine in simple whole number ratios. Early Models of the Atom Early Models of the Atom Atom: The smallest particle of an Atom: The smallest particle of an element that retains the chemical identity of that element. Approximately 114 elements. The alphabet of chemistry. The chemist’s goal is to understand the atoms that make up the world. Magnetism: The theory for the discovery of the electron and proton. Electrostatics: Opposite charges attract and likes repel. Benjamin Franklin: An object could have one of two charges (positive or negative). Michael Faraday (1839): Drew the connection between the electrical current and the structure of the atom. Discovering Atomic Discovering Atomic Structure Equipment to Isolate Sub­ Equipment to Isolate Sub­ Atomic Particles Eugene Goldstein: Designed the cathode ray tube and the canal ray tube. Vacuum inside, current flows between cathode and the anode (blue), source of electrons or x­rays. Discovery of the Electron Discovery of the Electron J.J. Thompson Defines the charge of the electron. Designs the mass spectrometer. Begins to define the mass of the electron. Places bar magnets around the tube and creates a magnetic field in which the beam of blue light bends toward the positive pole of the magnet. Defining the Electron Defining the Electron Robert Milikan Defines a more accurate mass of the electron. “Oil Drop Experiment” Timed the rate at which the charged oil drop would rise and fall depending upon the strength of the electromagnetic field vs. the strength of gravity. Discovery of the Proton Discovery of the Proton Eugene Goldstein: Defined the charge of the proton. J.J. Thompson: Defined the mass of the proton to be 1837 times the mass of the electron. The Nuclear Atom The Nuclear Atom Rutherford Gold Foil Experiment Concluded a nucleus existed. Contained the positive charge as well as the majority of the mass. Atomic Structure Atomic Structure Atoms are composed of protons, neutrons, and electrons. Create chart on the board of sub­ atomic particles for comparison. Atomic number = Z = # of protons. Atomic mass = M = # of protons and neutrons. An individual atom is electrically neutral. Protons = Electrons. p = Z, Charge = p ­ e, and n = M – Z. Atomic Number and Mass Atomic Number and Mass Created when an atom gains or loses one or more electrons. It acquires a net electrical charge. Charge of ion = the number of protons – the number of electrons. IONS IONS Dalton’s postulate that all atoms of a given element are identical is not exactly true. An isotope is an atom that has the same number of protons but a different number of neutrons. Isotopes Isotopes Why? Mendeleev Moseley Periodic Table Periodic Table Development Based on increasing atomic mass. Predictions. Based on increasing atomic number. Why did Mendeleev’s work? Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Metals: Properties include luster or shine, Metals: Properties include luster or shine, good conductors of heat/electricity, malleable, and ductile. Nonmetals: Large variation in properties. Metalloids or Semimetals: The in between. Family Names: Family Names: Representative (1A – 8A) Transition (1B – 8B) Inner Transition Alkali Metals Alkaline Earth Metals Boron family Carbon family Nitrogen family Oxygen family Halogens Noble gases “Strong Nuclear Force” between extremely close subatomic particles. Atomic #’s 1 thru 20: protons = neutrons. Past 20: Need more neutrons than protons to be stable. Past 83: All are radioactive. Too many or too few neutrons = radioactivity. Nuclear Chemistry Nuclear Chemistry The Changing Nucleus Radioactivity Radioactivity Alpha, Beta, and Gamma Distinguished by charge, mass, and penetrating power. Radioactive decay: when an atom emits one of these kinds of radiation. Decay = nucleus decays to form a new nucleus releasing radiation in the process. Nuclear reactions change the composition of the atom’s nucleus. Types of Radiation Name Identity Charge Penetrating Ability Alpha (α) Beta (β) Helium­4 nuclei 4 2He Electrons 0 ­1e 2+ Low, stopped by paper Medium, stopped by heavy clothing High, stopped by lead 1­ Gamma High energy (γ ) none Nuclear Equations Nuclear Equations Keep track of the reaction’s components. The sum of the mass numbers and atomic numbers are the same before and after. Sample problems. Harmful to living things: Effects on living tissue: Effects of Radiation Effects of Radiation Effects depend on the amount and type. Ionizing radiation: disrupts living cells. Somatic damage: direct. Genetic damage: reproduction. Uses of Nuclear Chemistry Uses of Nuclear Chemistry Radioactive dating. Smoke detectors. Imaging. Radiotracers. Cancer therapy. Food preservation. Energy. Measuring Radiation Measuring Radiation Equipment: Dosimeter. Geiger counter. A large nucleus is split into two smaller nuclei of approximately the same mass. The “missing mass” = energy. Nuclear chain reactions. Nuclear reactors. Three Mile Island (loss of coolant). Chernobyl (failure of the moderator). Waste disposal (fuel rods and burial). Nuclear Energy Nuclear Energy Nuclear Fission Two small nuclei join to form a large nucleus. Difficult to produce and control. Benefits: Electron cloud repulsion. Nucleus repulsion. Nuclear Energy Nuclear Energy Nuclear Fusion Problem: Uses hydrogen – abundant. No radioactive waste. High temps required. The half­life of a radioactive isotope is the time it takes for one half of a sample of that isotope to decay. Uranium­238: 4.5 billion years (alpha decay). Carbon­14: 5,730 years (beta decay). Sample problems. Half­life Half­life ...
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