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Parts of the Atom and Isotopes

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					Parts of the Atom and
       Isotopes



      Ms. Kakaliou
 A Modern Model Emerging
    1919 Moseley discovered the proton
    1932 British physicist James Chadwick discovered the
     neutron
    Review of the particles making up atom

Particle   Symbol    Relative     Relative   Actual Mass (g)
                     Electrical   Mass
                     Charge       (amu)
Electron   e-        -1           1/1837     9.11 x 10-28
Proton     p+        +1           1          1.673 x 10-24
Neutron    n0        0            1          1.674 x 10-24
Subatomic Particles - Roles
   Electrons
       Surround the nucleus
       Responsible for the volume of the atom

   Protons and Neutrons
       Located in the nucleus – makeup the small,
        dense, positively charged nucleus
       Responsible for the mass of the atom
Subatomic Particles-Definitions
  Atomic number: number of protons
      Identifies the element
  Mass Number: total number of protons and neutrons
  Notations
 Mass Number (superscript)
                              X OR Element’s Name – Mass Number
   Atomic Number (subscript)
  Examples
   197
      Au                        OR gold – 197
  79
  Number of Neutrons: (Mass Number – Atomic Number)
  Gold, Au : has 79 protons, 79 electrons, (197-79)=118 neutrons
Subatomic Particles-Definitions
   Number of protons = number of electrons, atoms are
    electrically neutral

   Number of protons ≠ number of electrons, atoms
    exist as charged species called ions

   Cation: Positively charged ion – forms when electrons
    leave the atomic structure (Na+: 11p+, 10e-)

   Anion: Negatively charged ion – forms when
    electrons are entering the atomic structure
    (F -: 9p+, 10e-)
Subatomic particles-Definitions
         Examples
    197
         Au3+
     79
   3 electrons have left the structure gold atom
   79 protons, 76 electrons, 118 neutrons

    31
         Si 2-
     14

   2 electrons have entered the structure of Si atom
   14 protons, 16 electrons, 17 neutrons
                        Isotopes
   Most elements exist as mixtures of isotopes
   Each isotope has a fixed mass and a natural percent abundance


        Isotope            Mass (amu)       Natural Abundance
                                                    %
          1H                       1.0078              99.985
          2H                       2.0140               0.015
          10B                     10.0129                20.0
          11B                     11.0093                80.0
          12C                      12.000               98.89
          13C                     13.0033                1.11
         Isotopes - Definition
   Isotopes are atoms that have the same
    number of protons, but different number of
    neutrons

   Isotopes have different mass numbers

   Isotopes of an element are chemically the
    same, because the electrons, not the
    neutrons, determine the atom’s chemical
    properties
                Relative Masses
   Atomic Mass Unit (amu or u) is defined as one-twelfth the
    mass of carbon-12 atom

   1 amu = (1/12) mass of C-12 atom

   The masses of all the atoms of the elements of the
    periodic table are expressed in relation to the mass of the
    C-12 atom

   Example : magnesium-24
   24 amu = 24 (1/12) mass of C-12 atom = 2 mass of C-12
    atom

   Magnesium-24 has 2 times more mass than the C-12
    atom
      Atomic Mass – Weighted
             Average
   The atomic mass of an element – the one
    reported at the P.T.- is not a whole number!

   The atomic mass of an element is a weighted
    average mass of the atoms in a naturally
    occurring sample of the element (all isotopes
    contribute)
        Atomic Mass – Weighted
               Average
   A weighted average reflects both the mass and the
    relative abundance of the isotopes as they occur in
    nature

   In order to calculate the atomic mass of an element
    you need to know:

       The number of stable isotopes of the element
       The mass of each isotope
       The natural percent abundance of each isotope
        Atomic Mass Calculation
      Isotopes          Mass (amu)          Abundance %
         35Cl              34.969                  75
         37Cl              36.966                  25
    Cl-35 contribution :34.969 amu × 0.75 = 26.227 amu
    Cl-37 contribution :36.966 amu × 0.25 = 9.231 amu
    Atomic mass for chlorine:              = 35.458 amu

NOTE: the atomic mass is closest to the most abundant isotope
  (Cl-35 mass= 34.969 amu)

				
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posted:10/12/2010
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