Transcript cc 6 atomic theory

History of Atomic Theory
Scientists B.C.
Democritus
400 BC
Coined the term “atom”
Aristotle
384-322 BC
Believed matter is continuous
Dalton’s Atomic Theory
-early 1800’s-
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All matter is composed of tiny, indivisible
particles called atoms.
Atoms of the same element have the same
properties (mass, size, etc.).
In a chemical reaction, matter cannot be created
or destroyed. (Law of Conservation of Mass)
Compounds always contain elements in the
same ratio by mass (Law of Definite Proportions)
Atomic size
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A penny contains 2.4 x 1022 atoms
Radius of one atom is around 2 x 10-10m
or .2 nm
Scanning tunneling microscope can
generate images of individual atoms.
Thomson’s Cathode Ray Tube
-late 1800’s-
Showed that electrons are negatively charged particles.
Image from Addison Wesley Chemistry
Thomson’s “plum pudding” model
Rutherford’s gold foil exp.
-early 1900’s-
Conclusion:
Most of an atom’s volume is empty space.
Rutherford’s “planetary” model
5 Models of the Atom
(a) Dalton's model (1803)
(c) Rutherford's model (1909)
(e) Electron-cloud model (present)
(b) Thomson's model (1897)
(d) Bohr's model (1913)
© Prentice-Hall, Inc.
Subatomic charge location mass Other feature
particle
proton
+
Nucleus 1 amu Defines the
element
-atomic no.
neutron
0
Nucleus 1 amu Change no. to
form isotopes
electron
-
Electron ~0
cloud
atom’s volume
-dictates
reactivity
Nuclear Forces
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Short-range forces that hold the nuclear
particles together.
Isotopes
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Atoms of the same
element that differ in
mass
Atomic no.=# protons
#protons=#electrons
Mass no.=#protons +
# neutrons (nucleons)
Num
f neutrons
Isotopes of Hydrogen
Nuclide
Protons
Neutrons
Protium
1
0
Mass
Number
1
Deuterium
1
1
2
tritium
1
2
3
Isotopes can be written two ways
108
47
Ag
207
82
Pb
80
35
Br
or bromine-80
Electrons
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Found in an electron cloud outside of the
nucleus (but not in paths like the planets)
1st energy level holds 2 electrons
2nd energy level holds up to 8
3rd energy level holds up to 18
Periodic Table
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Arranged by increasing atomic number
Rows are called periods
Columns are called groups
Average Atomic Mass
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An element’s atomic mass is the weighted
average of its naturally occurring isotopes.
Average Atomic Mass
Multiply the mass of each isotope by its
abundance to get the weighted average.
(% x mass)+ (% x mass) + . . .
100
Ex.: Boron is 80.20% boron-11 (atomic
mass 11.01 amu) and 19.80% boron-10
(atomic mass 10.01 amu).
What is the average atomic mass of
boron?
(11.01amu)(80.20) + (10.01amu)(19.80) =
100
=10.81 amu
Sample Problem
ex.: Neon has 2 isotopes.
Neon-20 has a mass of 19.992amu and
neon-22 has a mass of 21.991amu.
In an average sample of neon atoms, 90% will be
neon-20 and 10% will be neon-22.
Calculate the average atomic mass.
(90 x 19.992amu)+(10 x 21.991amu)
100
= 20.192 amu