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Chapter Three
ATOMIC THEORY NOTES
Important Concepts in a
Nutshell
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First person to theorize that matter was made up of tiny
particles was a Greek philosopher (algebra teacher),
DEMOCRITUS around 400 B.C.
Later Aristotle theorized that matter was continuous and
couldn’t be broken down into smaller parts. (He was
Wrong! He set atomic theory back thousands of years)
Neither Democritus nor Aristotle had any experimental
evidence to support their theories.
It would be another 2000 years before our modern
atomic theory would evolve.
Serious Atomic Research Begins
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In the 1790’s, because of the development of accurate
measuring instruments, scientists began to collect
data that would help develop atomic theory.
In 1806 John Dalton, an English school teacher,
developed a theory that is almost completely accepted
today.
Dalton’s Atomic Theory
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All matter is composed of extremely small
particles called atoms
Atoms of a given element are identical in size,
mass and other properties. Atoms of different
elements are different in size, mass and
properties.
Atoms cannot be subdivided, created or
destroyed.
Atoms of different elements combine in simple
whole-number ratios to form compounds.
In chemical reactions, atoms combine, are
separated or rearrange to form products.
Other laws developed in the 1790’s include:
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LAW OF CONSERVATION OF MASS: Matter can
neither be created nor destroyed. (In other words; the
mass of the reactants = the mass of the products)
LAW OF DEFINITE PROPORTIONS: Compounds
contain the same elements in the same proportions by
mass no matter the size of the sample. EXAMPLE---In 18 grams of water, there are 2 g of hydrogen and 16 g
of oxygen. A 1: 8 ratio.
In 50 grams of water there are 5.56 g of hydrogen and
44.44 g of oxygen which reduces down to a 1:8 ratio. No
matter how large or small the sample of water, the ratio
of the mass of hydrogen to the mass of oxygen is
ALWAYS 1:8.
Other “Laws” continued
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LAW OF MULTIPLE PROPORTIONS: If two or more
compounds are composed of the same two
elements, then the amount of the second element in
each separate compound which combines with the
same amount of the first element, form small, wholenumber ratios.
EXAMPLE: In compound
A, 1 g nitrogen combines with 1.14 g of oxygen. In
compound B,1 g nitrogen combines with 2.28 g
oxygen. The ratio of oxygen in compound A to the
oxygen in compound B is ……..
1.14 g : 2.28 g
which reduces to a 1:2 ratio
(a simple whole number ratio)
Cathode Ray Tubes
AKA Crookes Tube
In the late 1880’s, many experiments were
performed in which electric current was passed
through various gases at low pressures. They
were carried out in glass tubes known as cathoderay tubes.
Inferences Made From Cathode
Rays
 Cathode rays were deflected by a
negative magnetic field. Conclusion:
Cathode rays are negatively charged
 Paddle wheel moved when placed in
front of cathode ray. Conclusion:
Cathode rays are made up of particles
with mass
 Cathode ray tubes were made with
different metals and gases and the
results were always the same.
Conclusion: Electrons are present in all
matter
Other Inferences Made
 Since atoms have negative electrons but
are electrically neutral themselves, they
must contain some positive charge to
balance the negative electrons.
 Since electrons have so much less mass
than atoms, atoms must have some other
particles that account for most of their
mass.
Rutherford’s Gold Foil Experiment
Rutherford’s Conclusions based on
Gold Foil Experiment
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Atoms are made up of mostly empty space
There must be a very small, positive, dense bundle of
matter at the center of the atom (nucleus)
The electrons most likely orbit the nucleus like planets
orbit the sun
Other Facts about atomic structure
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Atoms are neutral because they contain equal
numbers of protons and electrons
Mass of a neutron is 1.675 x 10-27 kg
Mass of a proton is 1.673 x 10-27 kg
Mass of an electron is 9.109 x 10-31 kg
Short-range forces called “nuclear forces” hold
the protons and neutrons together
The density of a nucleus is extremely high: 2 x
108 metric tons/ mL
Facts about atomic structure cont.d
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The symbol “Z” represents atomic number (number of
protons) Z=9 means atomic number 9
The number of protons in the nucleus is what gives an
atom its identity. For example: an atom with 9 protons
is fluorine. The number of neutrons can vary but the
number of protons is always 9. If it doesn’t have 9
protons, THEN IT’S NOT FLUORINE!
ISOTOPES are atoms of the same element (same # of
protons) with different numbers of neutrons, therefore
different masses and thus different mass numbers!
Ex. Copper-63 (29 protons and 34 neutrons)
Copper-65 (29 protons and 36 neutrons)
Calculating Average Atomic
Mass
 Copper is a mix of Cu-63 and Cu-65. 69.17% of the
atoms are Cu-63 with a mass of 62.929 599 and the
other 30.83% are Cu-65 with a mass of 64.927 793.
 Let MAMA help you calculate average atomic mass
(it’s really MxA + MxA)
 “M” stands for mass of isotope and “A” stands for
abundance (percentage)
 Multiply the mass times the abundance and add to the
other mass x abundance. Don’t forget to change the
percent to a decimal
 Ex. (62.929 599 x .6917 )+ (64.927 793 x .3083 )=
Average Atomic Mass
Definition of Mole
 A mole is the amount of a substance that contains
as many particles as there are atoms in exactly 12 grams
of carbon-12. What does that mean?
 There are 6.022 136 7 x 1023 atoms in 12g of carbon. That
means….if you have that many atoms or molecules, you have
a mole!
 This number is called “Avogadro’s Number”
 One mole contains 6.022 136 7 x 1023 atoms or 6.022 136 7 x
1023 molecules or 6.022 136 7 x 1023 ions or 6.022 136 7 x
1023 particles etc.
Not as weird as you think!!!
 It’s not any different than the idea of a “dozen”. A dozen
is 12 eggs or 12 doughnuts or 12 chicken wings or 12
pencils…….
 A gross is 144 eggs or 144 pencils or 144 sodas…..
 A ream is 500 sheets of paper.
 So “MOLE” is the easy way to say ………
6.022 136 7 x 1023 atoms or
602 213 670 000 000 000 000 000 atoms
MOLAR MASS
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The mass of one mole of a pure
substance is called the molar
mass of that substance. The
molar mass of an element is numerically equal to the
average atomic mass of the element contained in the
periodic table EXCEPT THE UNITS ARE GRAMS NOT
AMU’s.
For example:
1 atom of Mg has a mass of 24.3 atomic mass units
1 mole (6.02 x 1023 atoms) of Mg has a mass of 24.3 grams
1 atom of Au has a mass of 196.967 atomic mass units
1 mole (6.02 x 1023 atoms) of Au has a mass of 196.967 grams
Gram Mole Conversions
Conversions from moles to grams and grams to moles
are simple if you use dimensional analysis.
Example: How many moles are contained in 57 grams
of sodium?
Your conversion ratio is the molar mass of sodium
which is found on the periodic table. 23g Na
1 mole Na
Begin by writing 57 g Na then multiply it by the
reciprocal of the molar mass so that the grams units
cancel.
57 g Na [ 1 mole Na]
Since the 23 is in the
[23 g Na ]
denominator, you should
divide 57 by 23 to get your ans.