Atomic Theory

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AtomicTheory

• A Brief History

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Atoms are made up of subatomic particles called
protons, neutrons and electrons

• How do we know that?

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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.

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• 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

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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)
• 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

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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

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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.

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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.
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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

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John Dalton(early 1803)

• Matter consists of tiny particles called atoms
  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

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Why were Daltons views accepted?

• The scientific method is now the proper way to
  do science.
• Daltons theory was based on experimental
  observations the law of Conservation of Mass and
  the law of Definite Proportions.
• Daltons theory correctly predicted the outcome
  of future experiments. These predictions became
  the law of Multiple Proportions.

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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.

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Amedo Avogadro

• Avogadro, among other achievements, was able to
  explain the existence of diatomic molecules.
• Avogadros Law Equal volumes of any gas at the
  same temperature and pressure, have the same
  number of particles.
• 1 mole 22.4 Liters

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J.J. Thomson set up a crookes tube with a anodic
and cathodic ends
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• 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?

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• 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

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Demonstration

• Molecular Expressions Electricity and Magnetism
  - Interactive Java Tutorials Crookes Tube

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• He concluded that the particles in the tube were
  negatively charged and had mass
• mass 9.109 x 10-31kg

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• 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!

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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

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J.J. Thomsons Plum Pudding Model
Positively charged pudding
Negatively charged particles later named electrons
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Thomson movie
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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

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• 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?

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Rutherfords Hypothesis England, 1911

• Rutherford hypothesized that the particles were
  travelling through a void and occasionally
  bouncing off a concentrated positive charge.

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Conclusion

• There must be a dense region with positive
  charges surrounded by the electrons
• 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.

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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

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Were 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

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Coulombs Law

• Since like charges repel, how can the nucleus be
  stable with protons () and neutrons (0)?
• Coulombs 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!

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Problems with Rutherfords 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.

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Elemental Quandary

• The Rutherford model was unable to explain the
  difference in the visible spectrum for each
  element.

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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 (?)

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Electromagnetic Spectrum

• All of the frequencies or wavelengths of
  electromagnetic radiation.

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Wavelength

• The wavelength is the distance between repeating
  units of a wave pattern (?) and measured in nm

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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.

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Hydrogen
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Carbon
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Oxygen
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Xenon
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Compare these spectrum

• Hydrogen, Carbon, Oxygen and Xenon

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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.

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Neils Bohr(early 1900s)

• 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.

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The Bohr Atom
Light
Excited States
-

Ground State
Nucleus
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The Bohr Planetary Atomic Model
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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

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The Modern Atom

• The modern atom is further defined by the works
  of these scientists
• de Broglie
• Max Plank
• Albert Einstein
• Heisenberg
• Erwin Schrodinger

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Problems with the Planetary Model

• This model only works for Hydrogen

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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
• Planks Constant
• hv
• Founded quantum mechanics theory
• He was also an accomplished musician!

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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.

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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

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Erwin Schrödinger Austria, 1920s

• 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.

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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

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The Modern Atomic View
The Wave-Mechanical Model
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Another View
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The Theory

• No two electrons can have the same quantum number
  (Pauli Exclusion Principle)
• 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

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Lets Review
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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.

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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

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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

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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

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Max Plank

• Father of Quantum Physics
• Electrons absorb and emit energy in discrete
  packets called quanta

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Erwin Schrödinger

• Electrons exist in specific orbitals and are
  assigned separate quantum numbers

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Summary

• The model of the atom changed over time. How?
  What? When? Where? Why?
• Get into your study groups and each student
  answer a different question.
• Write your responses on the bottom of your notes
  page.