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Atomic Theory
Democritus

The Greek philosopher Democritus (460 B.C. – 370
B.C.) was among the first to suggest the existence
of atoms (from the Greek word “atomos”)
 His atomic model was called the “solid sphere”
model
He believed that atoms were indivisible and
indestructible
 He named the smallest piece of matter “atomos,”
meaning “not to be cut.”

Democritus
Based on his theory:
 Atoms were infinite in number, always moving
and capable of joining together
 Atoms were small, hard particles that were all
made of the same material but were different
shapes and sizes.
 Problems with the theory:
did not explain chemical behavior
 was not based on the scientific method – but just
philosophy

John Dalton

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1766 - 1844
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.
Was called the “billiard ball” or “solid
sphere” model
John Dalton

Improved on Democritus’s theory
Dalton’s theory was based on scientific
experimentation
 His ideas account for the law of conservation
of mass - atoms are neither created nor
destroyed
 Agree with the law of constant composition elements combine in fixed ratios

John Dalton

Problems with Dalton’s Theory:
 Discoveries were made that showed
that atoms are divisible into smaller
subatomic particles:
 Electrons, protons, and neutrons
 There are many other particles
J.J. Thomson

Called the “plum pudding” model

Provided the first hint that an atom is
made of even smaller particles.
Thomson Model

He proposed that atoms were made from
a positively charged substance with
negatively charged electrons scattered
about, like raisins in a pudding.
Thomson Model

Thomson studied the passage of an electric
current through a gas.

As the current passed through the gas, it
gave off rays of negatively charged
particles.
J. J. Thomson
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Improvements:
 Discovered particles smaller than the
atom exist, therefore, the atom was
divisible
 He also concluded that there must be
positively and negatively charged
particles within the atom.
 Discovered the electron
Problems:
 Could not discover the proton
Ernest Rutherford’s
Gold Foil Experiment - 1911
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•
Positively charged alpha particle were fired at a thin sheet
of gold foil.
Some of the positively charged particles bounced away
from the gold sheet as if they had hit something solid. He
knew that positive charges repel positive charges.
Rutherford’s Findings
Conclusions/Improvements:
The nucleus is small
The nucleus is dense
The nucleus is positively charged
The electrons move around the
nucleus like planets orbit the sun.
e) Called the nuclear model or the
planetary model
a)
b)
c)
d)
Problems: Based on classical physics,
later abandoned because of the
discoveries in quantum physics.
Bohr Model

In 1913, the Danish scientist Niels Bohr
proposed an improvement to Rutherford’s
model. His model is sometimes referred
to as the Rutherford-Bohr model .
Bohr’s Model
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Electrons orbit the nucleus in orbits that have a
set size and energy.
The energy of the orbit is related to its size. The
lowest energy is found in the smallest orbit.
Radiation is absorbed or emitted when an
electron moves from one orbit to another.
Problems:
 We cannot predict the exact location and orbit
of electrons.
 Predictions fail to work on larger atoms.
James Chadwick
 1932 – confirmed the existence of the neutron – a
particle with no charge, but a mass nearly equal to
a proton
 He was trying to identify the extra mass of an
atomic nucleus by firing alpha particles into a
beryllium target and allowing the resulting radiation
to interact with paraffin wax.
 The interactions between the radiation and the
hydrogen in the wax led to the discovery of the
neutron.
James Chadwick
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Improvements:
With the discovery of the neutron, the
atomic model seemed more complete than
ever.
The overall charges remained the same
Now there no longer seemed to be a
discrepancy between the atomic mass and
the atomic number.
Electron Cloud
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A space in which
electrons are likely to
be found.
Electrons whirl about
the nucleus billions of
times in one second
They are not moving
around in random
patterns.
Location of electrons
depends upon how
much energy the
electron has.
Electron Cloud

Depending on their energy they are locked
into a certain area in the cloud.

Electrons with the lowest energy are found
in the energy level closest to the nucleus

Electrons with the highest energy are
found in the outermost energy levels,
farther from the nucleus.
Atomic Structure
The Atom
An atom consists of:
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nucleus – contains protons and neutrons
electrons in space around the nucleus.
The atom is mostly empty space
Electron cloud
Nucleus
ATOMIC COMPOSITION
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Protons (p+)
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Electrons (e-)
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+ electrical charge
mass = 1.672623 x 10-24 g
relative mass = 1.007 atomic mass units (amu)
but we can round to 1
negative electrical charge
relative mass = 0.0005 amu
but we can round to 0
Neutrons (no)
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no electrical charge
mass = 1.009 amu but we can round to 1
Atomic Number, Z
All atoms of the same element have
the same number of protons in the
nucleus, Z
13
Al
26.981
Atomic number
Atom symbol
AVERAGE Atomic Mass
Atomic Number
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Atoms are composed of identical protons,
neutrons, and electrons
 How then are atoms of one element different
from another element?
Elements are different because they contain
different numbers of PROTONS
The “atomic number” of an element is the number
of protons in the nucleus
# protons in an atom = # electrons
Atomic Number
Element
# of protons
Atomic # (Z)
Carbon
6
6
Phosphorus
15
15
Gold
79
79
Mass Number, A
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A carbon atom with 6 protons and 6
neutrons is the mass number = 12
atomic mass units
Mass Number (A)
= # protons + # neutrons
NOT on the periodic table…(Round the
AVERAGE atomic mass on the table)
A boron atom can have
A = 5 p + 5 n = 10 amu
A
10
Z
5
B
Mass Number
Mass number is the number of
protons and neutrons in the nucleus
of an isotope: Mass # = p+ + n0
p+
n0
e- Mass #
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Nuclide
Oxygen - 18
Complete Symbols/ Nuclear Notation
Contain the symbol of the element,
the mass number and the atomic
number.
Mass
Superscript →
number

Subscript →
Atomic
number
X
Symbols

Find each of these:
a) number of protons
b) number of
neutrons
c) number of
electrons
d) Atomic number
e) Mass Number
80
35
Br
Symbols

If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons
b) number of neutrons
c) number of electrons
d) complete symbol
Symbols
 If an element has 91
protons and 140 neutrons
what is the
a) Atomic number
b) Mass number
c) number of electrons
d) complete symbol
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
a) Atomic number
b) Mass number
c) number of protons
d) complete symbol
Isotopes
Isotopes
 Dalton
was wrong about all
elements of the same type being
identical
 Atoms of the same element can
have different numbers of
neutrons.
 Thus, different mass numbers.
 These are called isotopes.
Isotopes
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Frederick Soddy (1877-1956) proposed the
idea of isotopes in 1912
Isotopes are atoms of the same element
having different masses, due to varying
numbers of neutrons.
Soddy won the Nobel Prize in Chemistry in
1921 for his work with isotopes and
radioactive materials.
Naming Isotopes
 We
can also put the mass
number after the name of the
element:
carbon-12
carbon-14
uranium-235
Isotopes are atoms of the same element having
different masses, due to varying numbers of
neutrons.
Isotope
Protons Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Isotopes
Elements
occur in
nature as
mixtures of
isotopes.
Isotopes are
atoms of the
same element
that differ in
the number of
neutrons.
Isotopes &
Their Uses
Bone scans with
radioactive technetium99.
Isotopes & Their Uses
The tritium content of ground water is used to
discover the source of the water, for example, in
municipal water or the source of the steam from a
volcano.
Atomic Mass
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How heavy is an atom of oxygen?
 It depends, because there are different
kinds of oxygen atoms.
We are more concerned with the average
atomic mass.
This is based on the abundance
(percentage) of each variety of that
element in nature.

We don’t use grams for this mass because
the numbers would be too small.
Measuring Atomic Mass
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Instead of grams, the unit we use is the Atomic
Mass Unit (amu)
It is defined as one-twelfth the mass of a carbon-12
atom.
 Carbon-12 chosen because of its isotope purity.
Each isotope has its own atomic mass, thus we
determine the average from percent abundance.
To calculate the average:
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Multiply the atomic mass of each isotope by
it’s abundance (expressed as a decimal),
then add the results.
If not told otherwise, the mass of the
isotope is expressed in atomic mass units
(amu)
Atomic Masses
Atomic mass is the average of all the
naturally occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
Carbon = 12.011
% in nature
98.89%
1.11%
<0.01%
Average Atomic Mass
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weighted average of all isotopes
on the Periodic Table
round to 2 decimal places
Avg.
Atomic
Mass
(mass)(% )  (mass )(% )

100
D. Average Atomic Mass

EX: Calculate the avg. atomic mass of oxygen if
its abundance in nature is 99.76% 16O, 0.04%
17O, and 0.20% 18O.
Avg.
(16)(99.76 )  (17)(0.04)  (18)(0.20)
 16.00
Atomic 
100
amu
Mass
D. Average Atomic Mass

EX: Find chlorine’s average atomic mass if
approximately 8 of every 10 atoms are chlorine35 and 2 are chlorine-37.
Avg.
Atomic
Mass
(35)(8)  (37)(2)

 35.40 amu
10