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Atoms and the Periodic Table Atoms and the by zhangyun


									Atoms and the Periodic
       Mr. Holmes
 Russian chemist, Dimitri Mendeleev,
  searched for a way to organize the
  elements in the 1800’s.
 He arranged the elements in order of
  increasing mass, which revealed a pattern
  in their chemical properties.
 Elements are arranged by increasing
  atomic number and chemical properties.
 In 1800, there were only 31 elements to
 Elements can appear in three various
    – Solids
    – Liquids (aqueous)
    – Gases (gaseous)
   The periodic table is divided into three different
    regions (families)
    – Metal
    – Nonmetals
    – Metalloids
   There are two groups within the nonmetals
    – Noble gases
    – Halogens (are a series of nonmetal elements from
      Group 17)
 The horizontal rows of the periodic table
  are called periods.
 Elements increase by one proton and one
  electron as you go from left to right on the
  periodic table.
 Each vertical column in the periodic table
  is a group, which make up the families.
 Each group is given a roman numeral.
             Early Greek Theories
              400 B.C. - Democritus thought matter
               could not be divided indefinitely.
             • This led to the idea of atoms in a void.
                                  earth         air
             • 350 B.C - Aristotle modified an earlier
               theory that matter was made of four
               “elements”: earth, fire, water, air.
             • Aristotle was wrong. However, his
 Aristotle     theory persisted for 2000 years.
                John Dalton
 1800 -Dalton proposed a modern atomic model
  based on experimentation not on pure reason.
          •  All matter is made of atoms.
          •  Atoms of an element are identical.
          •  Each element has different atoms.
          •  Atoms of different elements combine
             in constant ratios to form compounds.
           • Atoms are rearranged in reactions.
• His ideas account for the law of conservation of
  mass (atoms are neither created nor destroyed)
  and the law of constant composition (elements
  combine in fixed ratios).
   Adding Electrons to the Model
Materials, when rubbed, can develop a charge
difference. This electricity is called “cathode rays”
when passed through an evacuated tube (demos).
These rays have a small mass and are negative.
Thompson noted that these negative subatomic
particles were a fundamental part of all atoms.
     1) Dalton’s “Billiard ball” model (1800-1900)
        Atoms are solid and indivisible.
     2) Thompson “Plum pudding” model (1900)
        Negative electrons in a positive framework.
     3) The Rutherford model (around 1910)
        Atoms are mostly empty space.
        Negative electrons orbit a positive nucleus.
     Ernest Rutherford (movie: 10 min.)
 Rutherford   shot alpha () particles at gold foil.
 Zinc sulfide screen     Thin gold foil
 Lead block
  substance path of invisible
Most particles passed through.
 So, atoms are mostly empty.
Some positive -particles
 deflected or bounced back!
Thus, a “nucleus” is positive &
 holds most of an atom’s mass.
 Atoms are the basic unit of matter.
 Atoms are made of small subatomic
    – Protons
    – Neutrons
    – Electrons
   Protons and neutrons are made up of 3
    smaller particles called quarks.
 Electrons have a negative charge
 Protons have a positive charge
 Neutrons have no charge
 Protons and neutrons make up the
 Electrons surround the nucleus
               The Atom
 Protons and
  neutrons are
  the same size
  and mass .
 Electrons are
  much smaller
  and lighter than
  protons and
  Atomic numbers, Mass numbers
 There  are 3 types of subatomic particles. We
  already know about electrons (e–) & protons (p+).
  Neutrons (n0) were also shown to exist (1930s).
 They have: no charge, a mass similar to protons
 Elements are often symbolized with their mass
  number and atomic number                      16
                                 E.g. Oxygen: 8 O
 These values are given on the periodic table.
 For now, round the mass # to a whole number.
 These numbers tell you a lot about atoms.
   # of protons = # of electrons = atomic number
   # of neutrons = mass number – atomic number
 Calculate # of e–, n0, p+ for Ca, Ar, and Br.
     Atomic   Mass   p+   n0   e–
Ca    20       40    20   20   20
Ar    18       40    18   22   18
Br    35       80    35   45   35
      Bohr - Rutherford diagrams
 Puttingall this together, we get B-R diagrams
 To draw them you must know the # of protons,
  neutrons, and electrons (2,8,8,2 filling order)
 Draw protons (p+), (n0) in circle (i.e. “nucleus”)
 Draw electrons around in shells
     He         Li
                                Li shorthand

    2 p+              3 p+             3 p+   2e– 1e–
    2 n0              4 n0             4 n0

Draw Be, B, Al and shorthand diagrams for O, Na
Be               B
     4 p+               5 p+            13 p+
     5 n°               6 n°            14 n°

     O                  Na

         8 p+ 2e– 6e–      11 p+ 2e– 8e– 1e–
         8 n°              12 n°
 The groups of the periodic table also tell
  you the number of valence electrons of
  each atom.
 The periods tell the valence shell of the
1A                                 8A

     2A   3A   4A   5A   6A   7A
 Isotopes have the same number of
  protons but differ in their number of
 Boron, for example, naturally appears as
  B-11. An isotope of B-11 would be B-10
  and B-12
    Which of the following is an isotope?
    a.   Li-8
    b.   Mg-24
    c.   N-14
    d.   C-13
 A substance formed by the chemical
  combination of two or more elements in
  definite proportions.
 Compounds are formed by forming
  chemical bonds
 Bond formation involves the electrons that
  surround each atomic nucleus.
       Types of Chemical Bonds
   Covalent Bonds
    – Sharing of electrons between molecules
   Ionic Bonds
    – Transfer of electrons between molecules
   Van der Waals Forces
    – Intermolecular forces of attraction between

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