Chapt4 2, 4-3
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Chapt4 2, 4-3

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Chapt4 2, 4-3 Chapt4 2, 4-3 Presentation Transcript

  • Chapter 4 Atoms and their structure
  • Idea of the atom
    • Original idea Ancient Greece (400 B.C.)
    • Democritus, Leucippus and Aristotle- Greek philosophers.
  • History of the atom
    • Late 1700’s - John Dalton- England.
    • Teacher- summarized results of his experiments and those of others.
    • Elements substances that can’t be broken down
    • In Dalton’s Atomic Theory - combined idea of elements with that of atoms.
  • Dalton’s Atomic Theory
    • All matter is made of tiny indivisible particles called atoms.
    • Atoms of the same element are identical, those of different atoms are different.
    • Atoms of different elements combine in whole number ratios to form compounds.
    • Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.
  • Parts of Atoms
    • J. J. Thomson - English physicist. 1897
    • Made a piece of equipment called a cathode ray tube.
    • It is a vacuum tube - all the air has been pumped out.
    • A limited amount of other gases are put in
  • Thomson’s Experiment + - Metal Disks Voltage source
    • Passing an electric current makes a beam appear to move from the negative to the positive end
    Thomson’s Experiment + - Voltage source
  • Thomson’s Experiment
    • By adding an electric field he found that the moving pieces were negative
    + - Voltage source
  • Thomson’s Experiment
    • Used many different metals and gases
    • Beam was always the same
    • By the amount it bent he could find the ratio of charge to mass
    • Was the same with every material
    • Same type of piece in every kind of atom
  • Thomsom’s Model
    • Found the electron.
    • Couldn’t find positive (for a while).
    • Said the atom was like plum pudding.
    • A bunch of positive stuff, with the electrons able to be removed.
  • Rutherford’s experiment
    • When the alpha particles hit a florescent screen, it glows.
    • Here’s what it looked like (pg 72)
  • Lead block Uranium Gold Foil Flourescent Screen
  • He Expected
    • The alpha particles to pass through without changing direction very much.
    • Because…
    • The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles.
  • What he expected
  • Because, he thought the mass was evenly distributed in the atom
  • Because, he thought the mass was evenly distributed in the atom
  • What he got
    • Atom is mostly empty.
    • Small dense, positive piece at center.
    • Alpha particles are deflected by it if they get close enough.
    How he explained it +
  • +
  • Modern View
    • The atom is mostly empty space.
    • Has two regions.
      • Nucleus -
      • Electron cloud-
  • Subatomic particles Electron Proton Neutron Name Symbol Charge Relative mass Actual mass (g) e - p + n 0 -1 +1 0 1/1840 1 1 9.11 x 10 -28 1.67 x 10 -24 1.67 x 10 -24
  • Structure of the Atom
    • There are two regions.
      • 1. The nucleus- contains protons and neutrons.
        • Positive charge.
        • Almost all the mass.
      • 2. Electron cloud- most of the volume of an atom (empty space).
        • The region where the electron can be found .
  • Counting the Particles Atomic Number Atomic Mass
    • Atomic Number = number of protons (# p)
      • # of protons determines kind of atom.
      • Helps find the atom in the periodic table
      • In neutral atom, atomic number = # of e
    • Atomic mass - avg mass of all isotopes of the atom
    • Mass Number = (rounded atomic mass) # p + # n
      • All the things with mass.
      • NOT on the periodic table
  • Symbols
    • if an element has an atomic number of 34 and a mass number of 79 what is the
      • number of protons
      • number of neutrons
      • number of electrons
      • Name
  • Isotopes
    • Dalton was wrong.
    • Atoms of the same element not identical, may have different masses.
    • called isotopes – different number of neutrons .
    • Do isotopes of an element have the same number of protons? Electrons?
  • Nuclear Symbols
    • Contain the symbol of the element, the mass number and the atomic number.
    X Mass number Atomic number
  • Naming Isotopes
    • Can also be represented as follows:
      • Put the mass number after the name of the element.
    • carbon- 12
    • carbon -14
    • uranium-235
  • Symbols
    • Find the
      • number of protons
      • number of neutrons
      • number of electrons
      • Atomic number
      • Mass Number
      • Name
    Na 24 11
  • Symbols
    • Find the
      • number of protons
      • number of neutrons
      • number of electrons
      • Atomic number
      • Mass Number
      • Name
    Br 80 35
  • Symbols
    • if an element has 91 protons and 140 neutrons what is the
      • Atomic number
      • Mass number
      • number of electrons
      • Complete symbol
      • Name
  • Symbols
    • if an element has 78 electrons and 117 neutrons what is the
      • Atomic number
      • Mass number
      • number of protons
      • Complete symbol
      • Name
  • Measuring Atomic Mass
    • Atomic mass – average mass of isotopes in an element
    • Unit is the Atomic Mass Unit (amu)
    • 1 amu = One twelfth the mass of a carbon-12 atom.
      • 1 amu ~ mass of p +
      • and n 0
      • Mass of isotopes not
      • Exact ( silicon-30 29.974 amu)
  • Calculating Atomic Mass
    • Atomic mass:
    • Atomic mass = (m 1 * % 1 ) + (m 2 * % 2 )+ (m n * % n )
    • m 1 = mass of Isotope 1
    • % 1 = abundance or percent of isotope 1 MUST BE IN DECIMAL
    • Use for as many isotopes in an element.
  • Atomic Mass
    • Calculate the atomic mass of copper if copper has two isotopes. 69.1% has a mass of 62.93 amu and the rest has a mass of 64.93 amu.
  • Atomic Mass
    • Magnesium has three isotopes . 78.99% magnesium 24 with a mass of 23.9850 amu, 10.00% magnesium 25 with a mass of 24.9858 amu , and the rest magnesium 25 with a mass of 25.9826 amu . What is the atomic mass of magnesium?
  • Atomic Mass
    • Is not a whole number because it is an average.
    • are the decimal numbers on the periodic table.