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4.1

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  • 1. Introduction to Atoms CHAPTER 4 SECTION 1
  • 2. History of Atom All atoms share the same basic structure During past 200 years, scientists have proposed different models. An atom is the smallest particle of an element. Atomic theory grew as a series of models that developed from experimental evidence. As more evidence was collected, the theory and models were revised.
  • 3. Dalton’s Model Based on experiments, Dalton developed a theory of structure of matter 4 main concepts:  All matter is composed of tiny, indivisible particles called atoms  Atoms of each element are exactly alike  Atoms of different elements have different masses  Atoms of different elements can join to form compounds
  • 4. Dalton’s Model
  • 5. Thomson’s Model End of 1800s Thomson discovered that atoms were not simple, solid spheres Atoms contained subatomic particles  Very small, negatively charged  Called them electrons
  • 6. Discovery of the ElectronIn 1897, J.J. Thomson used a cathode raytube to deduce the presence of a negativelycharged particle: the electron
  • 7. Modern Cathode Ray Tubes Television Computer MonitorCathode ray tubes pass electricitythrough a gas that is contained at avery low pressure.
  • 8. Thomson’s Model Also knew that atoms were electrically neutral  Must contain enough positive charge to balance negative charge of electrons Developed model where electrons were stuck into a positively charged sphere  Like chocolate chips in cookie dough
  • 9. Thomson’s Model
  • 10. Rutherford’s Model By early 1900s, scientists knew that positive charge of atom comes from subatomic particles called protons A proton is a positive charged particle in the nucleus of an atom. 1911—Rutherford begins to test theory His experiments led him to believe that protons are concentrated in a small area at center of atom  Called this area the nucleus
  • 11. Rutherford’s Model Rutherford’s model describes an atom as mostly empty space, with a center nucleus that contains nearly all the mass  Like the pit in a peach
  • 12. Bohr’s Model Modified Rutherford’s model in 1913 Proposed that each electron has a certain amount of energy  Helped electron move around nucleus Electrons move around nucleus in region called energy levels The energy level is the specific amount of energy an electron has. Energy levels surround nucleus in rings, like layers of onion
  • 13. Bohr’s Model Has been called planetary model  Energy levels occupied by electrons are like orbits of planets at different distances from the sun (nucleus)
  • 14. Electron Cloud Model Model accepted today Electrons dart around in an energy level Rapid, random motion creates a “cloud” of negative charge around nucleus Electron cloud gives atom its size and shape
  • 15. Electron Cloud Model
  • 16. Findings Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron) 1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton
  • 17. Atomic Number Atoms are composed of identical protons, neutrons, and electrons  How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS The “atomic number” of an element is the number of protons in the nucleus # protons in an atom = # electrons
  • 18. Atomic NumberAtomic number (Z) of an element isthe number of protons in the nucleusof each atom of that element. Element # of protons Atomic # (Z) Carbon 6 6 Phosphorus 15 15 Gold 79 79
  • 19. Mass NumberMass number is the number ofprotons and neutrons in the nucleusof an isotope: Mass # = p+ + n0Nuclide p+ n0 e- Mass #Oxygen - 18 8 10 8 18Arsenic - 75 33 42 33 75Phosphorus - 31 15 16 15 31
  • 20. The Complete Set-UP  Contain the symbol of the element, the mass number and the atomic number.Superscript → Mass numberSubscript → Atomic number X
  • 21. IsotopesDalton was wrong about all elements of the same type being identicalAtoms of the same element can have different numbers of neutrons.Thus, different mass numbers.These are called isotopes.
  • 22. Isotopes Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.
  • 23. Naming IsotopesWe can also put the massnumber after the name of theelement:carbon-12carbon-14uranium-235
  • 24. Isotopes are atoms of the same element havingdifferent masses, due to varying numbers ofneutrons. Isotope Protons Electrons Neutrons NucleusHydrogen–1 (protium) 1 1 0Hydrogen-2(deuterium) 1 1 1Hydrogen-3 1 1 2 (tritium)
  • 25. IsotopesElementsoccur innature asmixtures ofisotopes.Isotopes areatoms of thesame elementthat differ inthe number ofneutrons.

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