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Atomic Structure
Atoms and their structure
Mr. Bruder
Dalton’s Atomic Theory

John Dalton (1766-1844) had four theories
1.
All elements are composed of submicroscopic indivisible particles
called atoms
Atoms of the same element are identical. The atoms of anyone
element are different from those of any other element
Atoms of different elements can physically mix together or can
chemically combine w/ one another in simple whole-number ratios
to form compounds
Chemical reactions occur when atoms are separated, joined, or
rearranged. However, atoms of one element are never changed into
atoms of another elements as a result of a chemical reaction
2.
3.
4.
Atoms & Subatomic Particles

Atom- smallest particle of an element that
retains the properties of that element
A Helpful Observation
Gay-Lussac- under the same conditions of
temperature and pressure, compounds
always react in whole number ratios by
volume.
 Avagadro- interpreted that to mean
 at the same temperature and pressure, equal
volumes of gas contain the same number of
particles
 (called Avagadro’s hypothesis)

Electron
J.J Thomson (1856-1940) – discovered the
electron in 1897
 Electron is the negative charged subatomic
particle
 An electron carries exactly one unit of
negative charge & its mass is 1/1840 the
mass of a hydrogen atom
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Cathode Ray
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The Cathode Ray tubes pass electricity through
a gas that is contained at a very low pressure
Thomson’s Experiment
Voltage source
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Thomson’s Experiment
Voltage source
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Thomson’s Experiment
Voltage source
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Thomson’s Experiment
Voltage source
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Passing an electric current makes a beam
appear to move from the negative to the
positive end
Thomson’s Experiment
Voltage source
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+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Thomson’s Experiment
Voltage source
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+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Thomson’s Experiment
Voltage source
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+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Thomson’s Experiment
Voltage source
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By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field
Thomson’s Experiment
Voltage source
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 By adding an electric field he found that the
moving pieces were negative
Thomson’s Atomic Model
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Thomson’s Atomic Model
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Thomson though electrons were like plums
embedded in a positively charged “pudding”, so
his model was called the “plum pudding” model
Thomsom’s Model
Found the electron
 Couldn’t find
positive (for a while)
 Said the atom was
like plum pudding
 A bunch of positive
stuff, with the
electrons able to be
removed

Mass of Electron
In 1909 Robert Millikan determined the mass of an
electron with his Oil Drop Experiment
 He determined the mass to be 9.109 x 10-31 kg
 The oil drop apparatus

Millikan’s Experiment
Atomizer
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Oil
Microscope
Millikan’s Experiment
Atomizer
Oil droplets
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Oil
Microscope
Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking off
electrons
Millikan’s Experiment
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Millikan’s Experiment
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They put an electric charge on the plates
Millikan’s Experiment
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Some drops would hover
Millikan’s Experiment
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Millikan’s Experiment
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Measure the drop and find volume from 4/3πr3
Find mass from M = D x V
Millikan’s Experiment
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From the mass of the drop and the charge on
the plates, he calculated the charge on an electron
Proton
In 1886 Goldstein discovered the Proton
 Proton is a positively charged subatomic
particle found in the nucleus of a atom

Radioactivity
Discovered by accident
 Bequerel
 Three types
– alpha- helium nucleus (+2 charge, large
mass)
– beta- high speed electron
– gamma- high energy light

Ernest Rutherford
Rutherford (1871-1937) proposed that all mass
and all positive charges are in a small
concentrated region at the center of the atom
 He used the Gold-Foil Experiment to prove his
theory
 In 1911 he discovered the Nucleus
 Nucleus- central core of an atom, composed of
protons and neutrons
 The nucleus is a positively charged region and it
is surrounded by electrons which occupy most of
the volume of the atom

Rutherford’s Experiment
Used uranium to produce alpha particles
 Aimed alpha particles at gold foil by
drilling hole in lead block
 Since the mass is evenly distributed in
gold atoms alpha particles should go
straight through.
 Used gold foil because it could be made
atoms thin
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Lead
block
Uranium
Florescent
Screen
Gold Foil
What he expected
Because
Because, he thought the mass was
evenly distributed in the atom
What he got
How he explained it
Atom is mostly empty
 Small dense,
positive piece
at center
 Alpha particles
are deflected by
it if they get close
enough
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Rutherford’s Model
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Discovered dense
positive piece at the
center of the atom
Nucleus
Electrons moved
around
Mostly empty space
Neutron
James Chadwick (1891-1974) – discovered
the neutron in 1932
 Neutron is a subatomic particle with no
charge but their mass is nearly equal to that
of a proton

Bohr Model
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1.
2.
3.
4.
Bohr changed the Rutherford model and
explained how the electrons travel.
Bohr explained the following in his model:
Electrons travel in definite orbits around the
nucleus
Electrons are arranged in concentric circular
paths or orbitals around the nucleus
Electrons don’t fall into the nucleus because
electrons in particular path have fixed energy
and don’t lose energy
His model was patterned after the motion of the
planets around the sun. It is often called the
Planetary model.
46
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Quantum Theory


1.
2.
3.
Bohr explained how electrons were moving via
Quantum Theory
Key Terms:
Energy Levels- Regions around the nucleus
where the electron is likely moving
Quantum- Amount of energy required to move
an electron from one energy level to the next
Quantum Leap- Abrupt Change
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Bohr’s Model
Increasing energy
Fifth

Fourth
Third
Second
First
Nucleus
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Further away from
the nucleus means
more energy.
There is no “in
between” energy
Energy Levels
Bohr’s Model cont.
Energy levels are not equally spaced.
 Energy levels more closely spaced further
from the nucleus
 Higher energy level occupied by an
electron, the more energetic that electron is.
 Amount of energy gained or lost by an
electron is not always the same amount.
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50
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The Quantum Mechanical Model
Energy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move
from one energy level to another.
Since the energy of an atom is never “in between”
there must be a quantum leap in energy.
Schrodinger derived an equation that described the
energy and position of the electrons in an atom
Schrödinger’s Equation
The wave function is a F(x, y, z)
 Actually F(r,θ,φ)
 Solutions to the equation are called orbitals.
 These are not Bohr orbits.
 Each solution is tied to a certain energy
 These are the energy levels
Animation
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The Quantum Mechanical Model
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The atom is found inside
a blurry “electron cloud”
A area where there is a
chance of finding an
electron.
Draw a line at 90 %
Modern View
The atom is mostly
empty space
 Two regions
 Nucleus- protons and
neutrons
 Electron cloud- region
where you have a
chance of finding an
electron

Quark
Protons & Neutrons can still be broken down into
a smaller particle called the Quark
 The Quark is held together by Gluons

Density and the Atom
Since most of the particles went through, it
was mostly empty.
 Because the pieces turned so much, the
positive pieces were heavy.
 Small volume, big mass, big density
 This small dense positive area is the nucleus

Atomic Particles
Particle
Charge
Mass (kg)
Location
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
Neutron
0
1.675 x 10-27
Nucleus
Subatomic particles
Relative Actual
mass (g)
Name Symbol Charge mass
Electron
e-
-1
1/1840 9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
Symbols

Contain the symbol of the element, the mass
number and the atomic number
Mass
number
Atomic
number
X
Sub-atomic Particles
Z - atomic number = number of protons
determines type of atom
 A - mass number = number of protons +
neutrons
 Number of protons = number of electrons if
neutral

Symbols
A
X
Z
23
Na
11
Atomic Structure Symbols
Proton = p+
 Electron = e Neutron = n0

Atomic # - Subscript
 Mass # - Superscript

235
92
U
Rules for Atomic Structure
1.
2.
3.
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Atomic # = # of Protons
# of Protons = # of Electrons
Mass # = # of Protons + # of Neutrons
# of Neutrons = Mass # - # of Protons
If you know the Mass # & Atomic # you
know the composition of the element
Symbols
 Find
the
–number of protons
–number of neutrons
–number of electrons
–Atomic number
–Mass Number
80
35
Br
Symbols
 if
an element has an atomic
number of 34 and a mass number
of 78 what is the
–number of protons
–number of neutrons
–number of electrons
–Complete symbol
Symbols
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of protons
–Complete symbol
Example
Element Atomic Mass # Protons Electro Neutro
K
#
19
ns
19
11
16
5
17
46
35
ns
19
23
35
Isotopes
Isotope- atoms that have the same number
of protons but different number of neutrons
 Since isotopes have a different number of
neutrons the isotope has a different mass
number.
 Isotopes are still chemically alike because
they have the same number of protons and
electrons

Examples of Isotopes
Isotopes


Isotopes are atoms of the same element with different
masses.
Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
© 2009, Prentice-Hall, Inc.
Naming Isotopes
Put the mass number after the name of the
element
 carbon- 12
 carbon -14
 uranium-235

Electrical Charges
Electrical charges are carried by particles of
matter
 Atoms have no net electrical charges
 Given the number of negative charges
combines with the number of positive
charges = Electrically Neutral
 All elements are Electrically Neutral

Atomic Mass vs. Atomic Weight
Atomic Mass is for a single element
 Most elements are Isotopes
 How do we find their mass?
 We use Atomic Weight
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Measuring Atomic Mass
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Unit is the Atomic Mass Unit (amu)
One twelfth the mass of a carbon-12 atom
Each isotope has its own atomic mass. We need
the average from the percent abundance
Each isotope of an element has fixed mass and a
natural % abundance
You need both of these values to find the Atomic
Weight
Calculating Atomic Weight
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1.
2.
3.
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Cl-35 34.969amu and 75.77% abundance
Cl-37 36.966amu and 24.23% abundance
To solve for Cl-35
AMU x Abundance
34.969 x .7577
= 26.496
You solve for Cl-37
Atomic Weight Cont.
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1.
2.
3.
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Cl-37
AMU x Abundance
36.966 x .2423
= 8.957
Now you combine your two answers
26.496 + 8.957=
35.453
Look at Cl on the table. What is the Atomic
Weight?
Example
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Calculate the atomic weight of copper.
Copper has two isotopes. One has 69.1%
and has a mass of 62.93 amu. The other has
a mass of 64.93 amu. What is the atomic
weight???
Atomic Weight & Decimals
Atomic Weight- of an element is a
weighted average mass of the atoms in a
naturally occurring sample of an element
 Atomic Weights use decimal points because
it is an average of an element
