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Unit 5 Atoms and the Periodic Table

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Unit 5 Atoms and the Periodic Table Powered By Docstoc
					Unit 5: Atoms and the
Periodic Table
Chapter 4
Part 1: Atomic Structure
Objectives
 Explain Dalton’s atomic theory, and
  describe why it was more successful
  than Democritus’s theory.
 State the charge, mass, and location
  of each part of an atom according to
  the modern model of the atom.
 Compare and contrast Bohr’s model
  with the modern model of the atom.
Atomic Models
 Atom comes from the Greek word that
  means “unable to be divided.”
 Democritus (4 b.c.) - came up with the
  first theory of atomic structure; said
  that the universe was made of invisible
  units called atoms, but was unable to
  provide evidence.

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   John Dalton (1808) - used scientific
    research to claim that atoms could not
    be divided, that all atoms of an
    element were exactly alike, and that
    atoms of different elements could join
    to form compounds.


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   Neils Bohr (1913) - said that the
    electrons in an atom orbit around the
    nucleus like planets around the sun.
    The path of an electron is determined
    by how much energy it has, which puts
    it in a specific energy level.

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   Modern wave model (1925) - says that
    electrons behave more like waves on a
    vibrating string, and move back and
    forth between energy levels. Thus, an
    electron’s exact location at any given
    moment cannot be determined.


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  An Atom’s Contents
     Atoms contain smaller pieces called
      subatomic particles.


PARTICLE     CHARGE     LOCATION     MASS

proton       positive   in nucleus   1 amu

electron     negative   around       almost 0 amu
                        nucleus
neutron      neutral    in nucleus   1 amu
 Unreacted atoms (atoms that are not
  part of a chemical compound) have no
  overall charge.
 That means the number of positive
  charges (protons) must equal the
  number of negative charges
  (electrons).
 Neutrons do not affect the overall
  charge.
Electrons and Energy Levels
 If an electron does not have much
  energy, it will be closer to the nucleus
  in the first or second energy level.
 The first energy level will hold only 2
  electrons.
 If an atom has more than 2 electrons,
  the first two will fill the first level, and
  the rest will begin filling the second
  energy level.
 The second level will hold 8 electrons.
 If the first and second levels are full
  and there are still leftover electrons,
  they will go to the third level, which
  holds 18 electrons.
 Once the first three levels are full
  (that’s 28 electrons!), electrons must
  go to the fourth level, which holds 32
  electrons.
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Orbitals
 There are four different types of orbitals that
  can be found within the energy levels. They
  are s, p, d, and f.
 An s orbital has the lowest amount of
  energy and can hold 2 electrons.
 A p orbital has more energy than s orbitals,
  and there are 3 of them. Each one can hold
  2 electrons, for a total of 6.
 There are 5 d orbitals and 7 f orbitals. F
  orbitals have the most energy, and each
  one holds 2 electrons.
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Valence Electrons
 Electrons found in the outermost
  energy level of an atom.
 Each atom contains between 1 and 8
  valence electrons.
 These electrons are the ones that are
  used to form chemical bonds with
  other atoms to form molecules or
  compounds.
Part 2: The Periodic Table
Objectives
 Relate the organization of the periodic table
  to the arrangement of electrons within an
  atom.
 Explain why some atoms gain or lost
  electrons to form ions.
 Determine how many protons, neutrons,
  and electrons at atoms has given its symbol,
  atomic number, and mass number.
 Describe how the abundance of isotopes
  affects an element’s average atomic mass.
 Locate alkali metals, alkaline-earth
  metals, and transition metals in the
  periodic table.
 Locate semiconductors, halogens, and
  noble gases in the periodic table.
 Relate an element’s chemical
  properties to the electron arrangement
  of its atoms.
The Periodic Table
 The periodic law says that if the elements
  are arranged in a specific order, similarities
  in their properties will occur in a regular
  pattern.
 Period - horizontal row; the number of
  protons (and therefore, the number of
  electrons) increases by one as you move
  from left to right.
 Group (or family) - vertical column; elements
  in the same group have the same number of
  valence electrons, so they have similar
  characteristics.
The Periodic Table




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Groups
   Group 1 - Alkali metals (1 v.e.)
   Group 2 - Alkaline-earth metals (2 v.e.)
   Groups 3 -12 - Transition metals (number of
    v.e. varies)
   Group 13 - Boron family (3 v.e.)
   Group 14 - Carbon family (4 v.e.)
   Group 15 - Nitrogen family (5 v.e.)
   Group 16 - Oxygen family (6 v.e.)
   Group 17 - Halogens (7 v.e.)
   Group 18 - Noble (or inert) gases (8 v.e.)
 If an atom has only 1 v.e., it will be
  very reactive because it will want to
  stabilize itself by giving away its v.e.
  The goal of an atom is to become
  stable by having a totally full or totally
  empty outer energy level. This will
  cause it to be an ion with a +1 charge.
 Likewise, an atom with 7 v.e. will be
  very reactive because it only needs 1
  v.e. to be stable. Where can it find 1
  v.e.?
Calculating P, N and E
 Atomic number - the number of protons in
  an atom (and thus, the number of electrons)
 Mass number - the number of protons plus
  neutrons in an atom
 Isotope - atoms of an element that have the
  same number of protons, but different
  number of neutrons. This will not affect the
  number of electrons, but will affect the
  mass.
 Average atomic mass - the average mass of
  all the isotopes of an element
   To calculate the number of neutrons in
    an atom:
    Mass number - atomic number = neutrons

     Mass is measured in atomic mass units
      (amu).
     1 amu = 1/12 the mass of a standard C-
      12 atom
Part 3: Using Moles to Count
Atoms Objectives
 Explain the relationship between a
  mole of a substance and Avogadro’s
  constant.
 Find the molar mass of an element by
  using the periodic table.
 Solve problems converting the amount
  of an element in moles to its mass in
  grams, and vice versa.
Using Moles to Count Atoms

 We use moles to count atoms because
  they are so small and so numerous. If
  we know the mass of the atom, we can
  estimate how many atoms are in a
  sample of a substance by counting
  them in groups.
 The mole has a value of 6.022 x 1023
  particles in exactly 1 mole of substance.
  This value is called Avogadro’s
  constant.
Molar Mass
   The mass in grams of 1 mole of a
    substance is its molar mass, which is
    the same as its average atomic mass
    on the periodic table.
Using Conversion Factors
 A fraction that equals 1.
 For example: 12 in/1 ft = 1;
      5280 ft/ 1 mi = 1; 365.25 d/1 yr = 1

   Example: If you have 5.50 mol of Fe,
    and Fe has a molar mass of 55.85
    g/mol, what is its mass in g?
    5.50 mol Fe x 55.85 g Fe= 307 g Fe
                  1 mol Fe
Examples
   If you have 2.50 mol of S, and S has a
    molar mass of 32.07 g/mol, what is its
    mass in g?

    2.50 mol S x 32.07 g S = 80.18 g S
                 1 mol S
 A) 1.80 mol Ca
 B) 0.50 mol C
 C) 3.20 mol Cu
 A) 72.14 g Ca
 B) 6.01 g C
 C) 203.36 g Cu
   How many moles are in 620 g of Hg?

    620 g Hg x 1 mol Hg = 3.09 mol Hg
             200.59 g Hg
 A) 352 g Fe
 B) 11 g Si
 C) 205 g He
 A) 6.30 mol He
 B) 0.39 mol Si
 C) 51.25 mol He

				
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