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The ATOM
A History of the Atom
as brief as possible
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Democritus – matter
can be divided –
coined ‘atom’
Aristotle – matter is
continuous
Jump to 1700’s
Mendeleev –
organized periodic
table
Dalton – measured
pure substances
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Rutherford –
discovered nucleus
Bohr – found pattern
with atomic number
Schrodinger –
movement of
electrons
Planck and
Heisenberg – nuclear
chemistry
What the Heck?
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Atom comes from the Greek word
meaning indivisible.
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Today we know that what we once knew
as the atom, the smallest particle of life, is
in fact divisible and contains smaller
particles
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The Electron, Proton and Neutron
In an Atom:
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Protons and
Neutrons are in the
Nucleus
Electrons are in the
Electron cloud or
energy levels.
In the Atom:
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Nucleus is positively
charged in the center
of the atom
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Electron cloud is
negatively charged
area surrounding the
nucleus
To Illustrate
Dalton’s Laws
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Still followed to this day:
Law of conservation of mass: matter cannot
be created or destroyed in ordinary chemical
or physical changes
 Law of definite composition: a chemical
compound contains the same elements in
exactly the same proportions regardless of
the size of the sample or the source
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Modern Atomic Theory
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1) All Matter is made up of very small particles
called atoms
2) Atoms of the same element have the same
chemical properties
3) While individual atoms of a given element
may not all have the same mass any sample of
the element will have a definite average mass
that is characteristic.
Modern Atomic Theory Cont.
4) Compounds are formed whenever two
or more elements unite, with each atom
loosing its characteristic properties as a
result of the combination
 5) Atoms are not subdivided in physical or
chemical reactions
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Protons, Neutrons and Electrons
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Atomic Number = # of protons = # of Electrons
Atomic weight = # of protons + # of neutrons
So if you had an atom of Lead
Atomic Number of lead = 82
 Protons = 82
 Electrons = 82
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Atomic Mass of Silicon = 207.2
 Neutrons = Mass-Protons = 207.2 - 82 =
125
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Use Your Periodic Table
Name
Symbol
Atomic
Mass
Proton
Electron
Si
87.62
22
Iron
Bi
195.08
Zinc
77
71
35
Neutron
Masses of Subatomic Particles
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Proton - 1.67265 x 10-24 g
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Neutron - 1.67495 x 10-24 g
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Electron - 9.10953 x 10-28 g
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Electrons are 10,000 times smaller than
Protons and Neutrons!
Isotope
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Isotope - atoms with the same number of
protons but different numbers of neutrons.
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Number of protons make up the identity of
the atom. I.E. anything with 6 protons is
Carbon.
And why is the mass number a
decimal?
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Carbon: (2 natural isotopes)
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Carbon - 12 • (total weight = 12)
• (98.89 percent of Carbon atoms)
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Carbon - 13
• (total weight = 13)
• (1.11 percent of Carbon atoms)
(12 x .9889)+(13 x .0111)
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= 12.01 the atomic mass of Carbon
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Where the heck did that
equation come from?
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(12 x .9889) + (13 x .0111) = 12.01
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(12 x .9889) the weight of Carbon - 12
times the percent of Carbon - 12
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(13 x .0111) the weight of Carbon - 13
times the percent of Carbon - 13
Try it yourself:
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Neon - 22
(Total weight = 22)
 (10% of Neon)
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Neon - 20
(Total weight = 20)
 (90% of Neon)
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Once More:
Hf - 176 = 5%
 Hf - 177 = 19%
 Hf - 178 = 27%
 Hf - 179 = 14%
 Hf - 180 = 35%
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