Atoms_ Molecules and Ions History by ewghwehws

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									    Atoms, Molecules and Ions
            History
• Greeks
• Democritus 460-370 BC “atomos”
• Aristotle- elements.
• Alchemy
• 1660 - Robert Boyle- experimental definition
  of element.
• Lavoisier (1734-1794)- Father of modern
  chemistry. He wrote the book -1789.
                  Laws
• Conservation of Mass
• Law of Definite Proportion- compounds
  have a constant composition.
• Multiple Proportions- When two elements
  form more than one compound, the ratios of
  the masses of the second element that
  combine with one gram of the first can be
  reduced to small whole numbers.
          Multiple What???
• Water has 8 g of oxygen per g of hydrogen.
• Hydrogen peroxide has 16 g of oxygen per
  g of hydrogen.
• 16/8 = 2/1
• Small whole number ratios.
      Dalton’s Atomic Theory
            1803-1807
1) Elements are made up of atoms
2) Atoms of each element are identical. Atoms
  of different elements are different.
3) Compounds are formed when atoms
  combine. Each compound has a specific
  number and kinds of atom.
4) Chemical reactions are rearrangement of
 atoms. Atoms are not created or destroyed.
       A Helpful Observation
• Gay-Lussac- under the same conditions of
  temperature and pressure, compounds
  always react in whole number ratios by
  volume.
• Avogadro- interpreted that to mean nat the
  same temperature and pressure, equal
  volumes of gas contain the same number of
  particles.
• (called Avogadro’s Hypothesis)
 Thomson’s Experiment (1897)
• Used the Cathode Ray Tube(CRT) to
  discover the electron.
      Thomson’s Experiment




Passing an electric current makes a beam
appear to move from the negative (cathode)
to the positive end (anode).
   Proving Electrons had Mass
• Thomson devised a cathode ray tube
  with a paddle wheel built inside. When
  the high voltage electricity was turned
  on the paddle wheel began to rotate and
  move away from the cathode and
  towards the anode.
      Determining Charge
• Thomson concluded from this evidence
  and from his previous experiments that
  tiny particles were being emitted from the
  atoms of the cathode. These tiny particles
  were negatively charged. He called these
  particles "electrons."
Thomson’s Model
of the Atom
• Given these experimental results,
  Thomson proposed in 1897 what has
  since been referred to as the "plum
  pudding" model of the atom. This model
  depicts the atom as a diffuse cloud of
  positive charge with the negative
  electrons embedded randomly in it, like
  plums in pudding.
   Millikan (1909) - Mass of the
 Electron “(Oil Drop Experiment)”
• Using this information he calculated the mass
  of the electron as 9.11 X 10-28 grams.
               Radiation
• In 1896, the French scientist, Henri
  Becquerel, found that a piece of a mineral
  containing uranium could produce its
  image on a photographic plate in the
  absence of light. He attributed this
  phenomenon to a spontaneous emission
  of something which he called "radiation"
  which originated from the uranium.
                   Radioactivity
• Discovered by accident
• Bequerel
• Marie and Pierre Curie isolated the
  radioactive components at Bequerel’s
  suggestion
• Three types
• –alpha- helium nucleus (+2 charge, large
  mass)
•   __beta-   high speed electron
• –gamma- high energy light
         Rutherford (1911)
• A fluorescent screen would detect
  radiation by flashing whenever it was
  struck by the radiation.

• Rutherford observed that the radiation
  was diffracted into three beams by the
  charged plates.
Rutherford “Gold Leaf Experiment”
• Rutherford set up the apparatus which
  would bombard thin gold foil with alpha
  particles. These particles would then be
  detected by the fluorescent screen.

 He anticipated that all of the alpha
 particle detection would occur on the
 screen directly behind the foil.
 Rutherford got surprisingly different results
• "It was about as credible as if you had
  fired a 15-inch shell at a piece of tissue
  paper and it came back and hit you!"
      Atomic Model Revised
• Rutherford suggested that the atom was
  mostly empty space with a highly charged
  center. Most of the particles pass through
  the atom undisturbed, but a few get too
  close to the center and are deflected.
   "Planetary Model."



• To account for these results Rutherford
  proposed a new model of the atom in
  1913. This model had the following
  characteristics:

 The atom is mostly empty space with the
 majority of its mass concentrated in the
 center of the atom which he called the
 "nucleus."
    Mass Number and Atomic Mass
   • By the early 1930's the major subatomic
     particles had been discovered and their
     physical properties had been described.
   • James Chadwick (1932) - discovers
     Neutrons.
Particles    Mass (grams) Relative Mass (amu)      Relative
Charge
Proton      1.67262 X 10-27       1.007            +1
Electron 9.10939 X 10-31      5.486 X 10-4 (~ 0)   -1
Neutron     1.67493 X 10-27       1.009             0
          Atomic Mass Unit
• amu - the unit that we use to measure
  atoms.
• 1 amu ~ mass of one mole of the following:
  ~hydrogen ~ 1 proton ~ 1 neutron
• Actual masses:
• Particle       Charge           Mass (amu)
• Proton         positive(+1)     1.0073
• Neutron        none(neutral) 1.0087
• Electron       negative (-1)    5.486 x 10-4
• Hydrogen       none (neutral) 1.0079
          Measuring Atoms
              o
  Angstrom (A) - a convenient non-SI unit of
  length used to express atomic dimensions.

• 1 Angstrom = 1 x 10-10 meters

• most atoms have diameters between 1 x 10-10 m
                                    o
  and 5 x 10 -10 m, or between 1 - 5A.
                Isotopes
• All atoms have the same number of protons.
• The number of neutrons may vary for a
  given element.
• Isotope - atoms of a given element that
  differ in the number of neutrons and
  consequently the mass.
• May be written as the symbol 12C or simply
                                 6
  carbon-12 as opposed to the isotope 14C or
                                       6
  carbon-14.
 Atomic Numbers and Mass Numbers
• Atomic Number - is the number of protons, which is
  shown as the subscript. (it is also the number of
  electrons in a neutral atom.)
• Mass Number - is the total number of protons plus
  neutrons in the atom. (which represent essentially
  all the mass of the atom.)
• ex.       12C
             6
• has 6 Protons, 6 Electrons and 6 Neutrons
• Number of Neutrons = Mass # - Atomic #
                   6     =    12 -         6
                Symbols
• Contain the symbol of the element, the mass
  number and the atomic number.

           Mass
           number


           Atomic
           number
                    X
                 Nuclides
• Nuclide - an atom of a specific isotope
• ex.   14C or carbon-14
          6


• 6 protons, 6 electrons and 8 neutrons
• # of neutrons = atomic mass - atomic number
•      8         =          14       -   6
                   Symbols
• Find the
  –   number of protons
  –   number of neutrons
  –   number of electrons    24
  –
  –
      Atomic number
      Mass Number
                             11   Na
  –   Name
            Symbols
 Findthe
  –number of protons
  –number of neutrons
                         80
  –number of electrons
                         35   Br
  –Atomic number
  –Mass Number
  – Name
            Symbols
   an element has 91 protons and
 if
  140 neutrons what is the
  –Atomic number
  –Mass number
  –number of electrons
  –Complete symbol
   – Name
             Symbols
 ifan element has 78 electrons and
  117 neutrons what is the
   –Atomic number
   –Mass number
   –number of protons
   –Complete symbol
   – Name
                 Atomic Mass
•   How heavy is an atom of oxygen?
•   There are different kinds of oxygen atoms.
•   More concerned with average atomic mass.
•   Based on abundance of each element in nature.
•   Don’t use grams because the numbers would be
    too small.
        Measuring Atomic Mass
•   Unit is the Atomic Mass Unit (amu)
•   One twelfth the mass of a carbon-12 atom.
•   6 p+ and 6 n0
•   Each isotope has its own atomic mass
•   we get the average using percent
    abundance.
        Calculating averages


                    Average =
(% as decimal x mass) + (% as decimal x mass)
+ (% as decimal x mass )
                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?
• If not told otherwise, the mass of the isotope is
  the mass number in amu
                 Periodic Table
• The rows on the periodic
  chart are periods.
• Columns are groups.
• Elements in the same
  group have similar
  chemical properties.
                 Groups




These five groups are known by their names.
Periodic Table


         Nonmetals are on the
         right side of the
         periodic table (with
         the exception of H).
Periodic Table


         Metalloids border the
         stair-step line (with
         the exception of Al
         and Po).
Periodic Table


         Metals are on the left
         side of the chart.
      Diatomic Molecules




These seven elements occur naturally as molecules
containing two atoms.

								
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