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					Atomic Structure

The History of the Atomic Model
John Dalton (1766 – 1844)
    Dalton proposed the Atomic Theory in
     1805 which stated that
1.   All matter was composed of small
     indivisible particles termed atoms
2.   Atoms of a given element possess unique
     characteristics and weight
3.   Three types of atoms exist: simple
     (elements), compound (simple molecules),
     and complex (complex molecules)
    Dalton's theory identified chemical
     elements as a specific type of atom.
John Dalton (1766 – 1844)

   Dalton inferred proportions of elements
    in compounds by taking ratios of the
    weights of reactants, setting the atomic
    weight of hydrogen to be identically one.
   Dalton’s view of the atom is that of a
    solid sphere, similar to a billiard ball.
J. J. Thomson (1856 – 1940)

   Thomson discovered the electron, and
    showed that it was part of the structure
    of the atom.
   He did this using Gas Discharge tubes.
   The gas discharge tube is a glass tube
    that has most of the air pumped out, and
    two electrodes at either end.
Gas Discharge Tube
J. J. Thomson (1856 – 1940)

 In 1887, Thomson determined the
  rays were due to negatively charged
  particles that have mass.
 He called them ‘corpuscles’, later
  called electrons.
 Thomson received the Nobel Prize in
  1906 for his discovery.
J. J. Thomson (1856 – 1940)
   He theorized that the electrons were
    much like raisins in pudding, where
    the pudding is the positive particle.
Ernest Rutherford (1871-1937)

 Henri Becquerel discovered
  radioactivity in 1896.
 Ernest Rutherford discovered that
  radiation could be split into 3 types of
  beams with magnetic fields.
     Alpha particles (positive particle)
     Beta particles (negative particle)
     Gamma rays (neutral wave)
Ernest Rutherford (1871-1937)

 Rutherford became interested in the
  atomic structure.
 He theorized that using Thomson’s
  model, alpha particles should pass
  through atoms unaffected.

                  Alpha particles would pass directly
                  through the gold atoms according
                  to the Thomson model of the atom
                  since it is “a sea of positive charge
                  embedded with negative charges.”
Ernest Rutherford (1871-1937)

 Rutherford devised an experiment to
  test his theory.
 Alpha particles can be detected as a
  flash of light when they strike zinc
  sulfide.
 A sheet of gold is hammered into a
  very thin sheet of foil.
 Flashes of light should form only
  opposite the emitter.
Ernest Rutherford (1871-1937)

 Rutherford was shocked that the a
  particles bounced backwards.
 He concluded that the atom contained
  a tiny, dense core that was positively
  charged.
                     The atom is mostly empty
                     space since the majority of
                     alpha particles pass through
                     unaffected, but there is a tiny
                     positively charged core that
                     deflected the positively
                     charged alpha particle.
Anatomy of Rutherford’s
Model
The Electron Problem

 The nucleus contains positively
  charged particles called protons.
 Electrons are 1/2000 the mass of
  protons.
 Opposites attract!

 Why don’t the electrons crash into the
  nucleus?
James Maxwell (1831-1879)

 “[The work of Maxwell] ... the most
  profound and the most fruitful that
  physics has experienced since the
  time of Newton.” – Albert Einstein
 At the beginning of the 19th century,
  there were two theories regarding the
  nature of light.
James Maxwell (1831-1879)
 Newton proposed that light was a
  particle.
 Huygen proposed the light was a wave.
 James Maxwell proposed that light is an
  electromagnetic wave.
 The frequency of electromagnetic
  waves is continuous, and visible light is
  just one part of the range of frequencies
  possible.
 The range of frequencies is known as
  the Electromagnetic Spectrum.
James Maxwell (1831-1879)
Max Planck (1858-1947)

   Max Planck became interested in studying
    black bodies that emit electromagnetic
    radiation.
   When a solid is heated, it begins to glow.
   The hotter the object, the higher the
    frequency of electromagnetic radiation
    emitted.
   The problem was that if light is a continuous
    wave, then the math predicts a curve that is
    not experimentally observed.
Max Planck (1858-1947)
                Actual curve




                      Predicted curve
Max Planck (1858-1947)

 Planck was able to find a
  mathematical formula that fit the
  empirical data IF he treated light as if
  it was discrete, NOT continuous.
 He treated the energy from light in
  discrete amounts, or packets he
  called a quantum of energy.
 This was the first indication of light
  acting as a particle instead of a wave.
Max Planck (1858-1947)

   “But even if the radiation formula should
    prove to be absolutely accurate it would
    after all be only an interpolation formula
    found by happy guesswork, and would thus
    leave one rather unsatisfied. I was,
    therefore, from the day of its origination,
    occupied with the task of giving it a real
    physical meaning …”
       (Max Planck, 1919 Nobel Prize address,
        'The Origin and Development of the
        Quantum Theory')
Albert Einstein (1879-1955)

 In 1887, Heinrich Hertz discovered the
  Photoelectric effect accidentally.
 This occurs when electromagnetic
  radiation (light) is shone on a metal
  plate, which results in flowing
  electrons.
Photoelectric Effect
Albert Einstein (1879-1955)

 The classical theory of light (as a
  wave) predicted that the brighter the
  light, the greater the number of
  liberated electrons.
 However, experimental evidence
  showed that brightness of light does
  not increase electron flow.
 However, a higher frequency of light
  does!
Albert Einstein (1879-1955)

   In the first of 4 of Einstein’s major papers,
    he proposed an explanation for the
    Photoelectric effect.
   Einstein used Planck’s quantum theory to
    say that light consisted of streams of
    Planck’s quanta, and called them photons.
   It is the collision of photons with electrons
    that breaks them free from the atom.
   The higher the energy/frequency, the
    greater the chance the collision will break
    an electron free.
Albert Einstein (1879-1955)
Um, what about those
electrons?
 Now that we understand the quantum
  nature of light, we can now return to
  atomic theory, and consider how Neils
  Bohr used the idea of light as discrete
  energy packets to explain the nature
  of electrons.
 Questions – Page 173 #1, 3, 5-7
   Answers to homework
1. The first experimental observation was that the
electromagnetic radiation emitted by black bodies did
not agree with the classical theory of light as a wave.
Planck deduced that the amount of energy released
could be predicted if light were treated as packets of
energy, or quanta.

The second experimental observation was that the
amount of electrons liberated by the photoelectric
effect only increased in response to increased
frequency, not brightness. Einstein used Planck’s
quantum theory to explain the photoelectric effect as
due to photons colliding with electrons.
   Answers to homework
3. The photoelectric effect occurs when
electromagnetic radiation (e.g. visible light, ultraviolet,
etc) is directed at a metal. This results in the release
of electrons from the metal (i.e. current) due to
collisions between photons and electrons.
     Answers to homework
6.
E = hf = 6.6x10-34 J/Hz x 5.5 x1014 Hz = 3.63x10-19 J

b.

7 a. UV = 9.9x10-19 J, IR = 2.178x10-19 J

b. UV:IR = 4.5:1

c. Visible light photons are between UV and IR in
energy, and X-rays have more energy than all of
them.

				
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