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Atomic
Theory
A Brief History
Atoms are made up of subatomic
particles called protons, neutrons
and electrons
 How do we know that?
Vocabulary
 Atom: The smallest unit of an element,
having all the characteristics of that element
and consisting of a dense, central, positively
charged nucleus surrounded by a system of
electrons.
 Molecule: The smallest particle of a
substance that retains the chemical and
physical properties of the substance and is
composed of two or more atoms.
 Compound: A compound is a substance
made up of atoms representing more than
one element bonded together and exhibiting
distinct physical and chemical characteristics

Example: H2O, H2SO4
Background
 Law of Conservation of Mass (Lavoisier, 1789)

During a chemical reaction, the total mass of the
reactants is equal to the total mass of the
products.
 Law of Definite Proportions (Proust, 1799)
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When atoms combine to form compounds, they
always combine in the same simple, small whole
number proportions.
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Example: Water is always H2O
Example: Sulfuric Acid is always H2SO4
Aristotle
(circa. 400 B.C.)
 Matter is not made of particles, but rather is
continuous.
 The continuous matter is called “hyle.”
 There were only four elements

Earth, Air, Fire, Water
Democritus
(circa. 400 B.C.)
 Matter is made of empty space and
tiny particles called “atoms.”
 Atoms are indivisible.
 There are different types of atoms for
each material in the world.
Why was Democritus Ignored?
Because the early Greek philosophers
did not experiment and because
Aristotle was an established teacher and
because the church was opposed to
“soul atoms”, the views of Democritus
were not accepted until the 19th century.
Pre-Atomic Theory Postulates
 Law of Conservation of Mass

During a chemical reaction, the total mass of the
reactants is equal to the total mass of the
products.
 Law of Definite Proportions

When atoms combine to form compounds, they
always combine in the same simple, small whole
number proportions.


Example: Water is always H2O
Example: Sulfuric Acid is always H2SO4
John Dalton
(early 1803)
 Matter consists of tiny particles called atoms
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which are indivisible and indestructible.
All atoms of a particular element are identical.
Atoms of different elements differ in mass and
properties.
Atoms combine in whole number ratios to form
compound atoms.
In chemical reactions, atoms are combined,
separated, or rearranged but are never created,
destroyed, or changed
Why were Dalton’s views accepted?
 The scientific method is now the proper way to
“do science.”
 Dalton’s theory was based on experimental
observations: the law of Conservation of Mass
and the law of Definite Proportions.
 Dalton’s theory correctly predicted the outcome
of future experiments. These predictions
became the law of Multiple Proportions.
The Dalton Atom
 John Dalton examined the empirical
proportions of elements that made up
chemical compounds.
 At this stage, the atom was still seen as an
indivisible object, with no internal structure.
Amedo Avogadro
 Avogadro, among other achievements, was
able to explain the existence of diatomic
molecules.

Avogadro’s Law: Equal volumes of any gas at
the same temperature and pressure, have the
same number of particles.

1 mole = 22.4 Liters
J.J. Thomson set up a crookes
tube with a anodic and cathodic
ends
 When electricity was applied to the tube, a
beam was emitted from the cathodic (-) plate
 Thomson then assumed the particles
emitted were negative
 To test this theory, he applied a magnetic
field to the tube and “bent” the beam
 What happens with like charges?
 He tested the tube further by applying an
electrical field to the tube using paddles
 The tube turned around
 Thomson determined that the tube turned as
tiny particles hit the paddles
Demonstration
 Molecular Expressions: Electricity and
Magnetism - Interactive Java Tutorials:
Crookes Tube
 He concluded that the particles in the tube
were negatively charged and had mass

mass = 9.109 x 10-31kg
 Since these particles are negatively charged,
but the atoms are neutral, there must be
other particles in an atom
 Problem: This requires too many electrons!
Thomson Model
 The discovery of the electron by J. J.
Thomson showed that atoms did have some
kind of internal structure.
 The Thomson model of the atom described
the atom as a "pudding" of positive charge,
with negatively charged electrons embedded
J.J. Thomson’s Plum Pudding
Model
Positively
charged
“pudding”
Negatively
charged particles
later named
electrons
Thom
son
movie
Milliken and the Oil Droplet
 In 1909, Robert Milliken performed an
experiment using droplets of oil to determine
the charge of an electron.

electrons, e, e-, -1.602 x 10-19C
 Ernest Rutherford conducted experiments to
test the Thomson model
 He directed alpha particles through a thin
gold foil and measured them with a film
 Most particles went through the foil
•But, some were deflected, Why?
Rutherford’s Hypothesis
England, 1911
 Rutherford hypothesized that the particles
were travelling through a void and
occasionally bouncing off a concentrated
positive charge.
a
Conclusion
 There must be a dense region with positive
charges surrounded by the electrons
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An atom is mostly empty space with a dense
region in the middle.
This dense region is called the “nucleus”
He measured the number of particles deflected
and the angles and calculated that the radius of
the nucleus was 1/10,000 of the whole atom
Problem: Electrons should spiral into the
nucleus.
Let there be protons!
 The discovery was made and protons were
recognized
 The mass of a proton is 2000x the mass of an
electron
 1.673 x 10-27 kg
We’re not done yet ...
 30 years later, Irene Curie, the daughter of
the great Madame Curie, produced a beam of
particles that could go through almost
anything
 And James Chadwick determined this beam
was not affected by a magnetic field (no
charge!)
 Neutrons were given credit
Coulomb’s Law
 Since like charges repel, how can the
nucleus be stable with protons (+) and
neutrons (0)?
 Coulomb’s Law: the closer two charges are,
the greater the force between them
 As the distance between like charges
decreases, the force between them
increases.
 Try it!
Problems with Rutherford’s model
 According to classical physics, an electron in
orbit around an atomic nucleus should emit
photons continuously as they are accelerating
in a curved path.
 The loss of energy should cause the electron
to collide with the nucleus and collapse the
atom.
Elemental Quandary
 The Rutherford model was unable to explain
the difference in the visible spectrum for each
element.
Visible-line Spectrum
 When an elemental gas is excited by
electricity, it emits a distinct visible light
pattern.
 The color of each spectral line is identified by
the wavelength ()
Electromagnetic Spectrum
 All of the frequencies or wavelengths of
electromagnetic radiation.
Wavelength
 The wavelength is the distance between
repeating units of a wave pattern (λ) and
measured in nm
Frequency
 Frequency is the measurement of the
number of times that a repeated event occurs
per unit of time (Hz)
 The blue wave has the greatest frequency.
Hydrogen
Carbon
Oxygen
Xenon
Compare these spectrum
•Hydrogen, Carbon, Oxygen and Xenon
In comes Niels Bohr
Denmark, 1913
 In 1913, Bohr proposed that electrons were
restricted to certain fixed circular orbits.
 Orbits are energy levels
 Electrons can jump from ground state to an
excited state by absorbing energy or a photon
with the precise wavelength.
Neils Bohr
(early 1900’s)
Electrons travel around the nucleus in
specific energy levels.
 Electrons have a ground state and an
excited state
 Electrons do not radiate energy in their
normal energy level called the ground
state.
 Electrons absorb energy and move to
energy levels further from the nucleus
called excited states.
 Electrons lose energy (light) as they
return to lower energy levels.

The Bohr Atom
Excited States
+
Nucleus
Ground
State
The Bohr Planetary Atomic Model
The Bohr Atom
In the Bohr Model the neutrons
and protons occupy a dense
central region called the
nucleus, and the electrons
orbit the nucleus much like
planets orbiting the Sun
The Modern Atom
 The modern atom is further defined by the
works of these scientists:
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de Broglie
Max Plank
Albert Einstein
Heisenberg
Erwin Schrodinger
Problems with the Planetary Model
 This model only works for Hydrogen
Max Plank
Germany, 1918
 Energy is gained or lost in discrete
“packets” called quanta
 Calculated the amount of energy
and determined that it is a constant
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Plank’s Constant
hv
 Founded quantum mechanics
theory
 He was also an accomplished
musician!
de Broglie, 1924
 Electrons move like waves and so have
properties of waves.
Albert Einstein
 Einstein was simultaneously working on the
photoelectric effect, the theory of relativity
and the energy-mass relationship.
Heisenberg, 1925
 Heisenberg proposed that it is not possible to
know the position and momentum of an
electron at the same time.

Heisenberg Uncertainty Principle
Erwin Schrödinger
Austria, 1920’s
 Electrons have characteristics associated
with waves and particles; wave-particle
duality.
 Electrons are located around the nucleus
in “orbitals”

An orbital is a probability that an electron
will be there
 4 quantum numbers indicate the
probable location of the electron wave.
Schrödinger Wave Equation
2/x2 + 2/y2 + 2/z2 + 82m/h2(E-V)=0
(E-V) = 2 2me4/h2n2
 The equation predicts the orbital
The Modern Atomic View
The Wave-Mechanical Model
Another View
The Theory
 No two electrons can have the same
quantum number (Pauli Exclusion Principle)
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No two electrons can occupy the same space
at the same time
A quantum number is an address of the
electron
 Electrons exist in orbitals around the nucleus
Let’s Review
The Dalton Atom
 John Dalton examined the empirical
proportions of elements that made up
chemical compounds.
 At this stage, the atom was still seen as an
indivisible object, with no internal structure.
Thomson Model
 The discovery of the electron by J. J.
Thomson showed that atoms did have some
kind of internal structure.
 The Thomson model of the atom described
the atom as a "pudding" of positive charge,
with negatively charged electrons embedded
Rutherford Model
 The Rutherford model described the atom
made up of a dense nucleus of
approximately containing positively charged
particles, surrounded by an electron cloud of
approximately.
 “Nuclear Model”
Niels Bohr
 The Bohr Model is probably familar as the
"planetary model" of the atom, the figure is
used as a symbol for atomic energy
 The neutrons and protons occupy a dense
central region called the nucleus, and the
electrons orbit the nucleus much like
planets orbiting the Sun
Max Plank
•Father of Quantum Physics
•Electrons absorb and emit
energy in discrete “packets”
called quanta
Erwin Schrödinger
 Electrons exist in specific orbitals and are
assigned separate quantum numbers
Summary
 The model of the atom changed over time.
How? What? When? Where? Why?
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Get into your study groups and each student
answer a different question.
Write your responses on the bottom of your
notes page.