The Electromagnetic Spectrum and the Model of the Atom Part I

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The Electromagnetic Spectrum and the Model of the
AtomPart I

• Chemistry Mrs. Cameron

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The Purpose of Science

• The purpose of science to make models that
  explain natural phenomena.

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Scientific Models

• A model is the best possible explanation which
  accounts for all observed phenomenon and has
  predictability.
• 1) An unanswered question means change the
  model.
• 2) Predictability is the test of a good or
  true model.

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Radioactivity and Light have been tools to
discover the structure of atoms.

• Do you Remember?
• Thompsons Cathode Ray Experiment
• Rutherfords Gold foil Experiment

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Thomsons Plum-Pudding Model
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Ernest Rutherfords Gold Foil Experiment
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• Results of the Rutherford experiment

(a) The results that the metal foil experiment
would have yielded if the plum pudding model had
been correct
(b) Actual results
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• Rutherfords Nuclear Model of the atom
• Small, dense, positively charged nucleus
• Contains protons (1 charge)
• Contains neutrons (no charge)
• Remainder of the atom is mostly empty space
• Contains electrons (-1 charge) in the empty
  space

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New Evidence Continuous and Line Spectra

• Lets Talk about LIGHT

White light
is actually made up of many colors
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White Light

• Given off by objects heated to a very high
  temperature.
• When an object is heated, it first gives off a
  red glow then, as more energy is added it begins
  to glow white hot.

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When white light is passed through a prism, the
light refracts, or bends to display all of its
component colors
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The Wave Nature of Light

• All waves exhibit similar characteristics and
  properties
• Crest the top of a wave
• Trough the bottom of the wave
• (on next slide)
• Origin the center line through which the wave
  oscillates

Wave crest or peak
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The Wave Nature of Light
Crest ?
Origin ?
?trough
Amplitude distance from the origin to the crest
or the origin to the trough of a wave. (Indicates
intensity) Frequency (?) the number of waves
that pass a given point in a given amount of
time Wavelength (?) - distance from crest to
crest or trough to trough Speed (Velocity) the
amount of distance covered in a specified amount
of time.
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The Wave Nature of Light
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Wave Behaviors

• Reflection a wave strikes an object and bounces
  off
• Refraction the bending of waves
• Diffraction the bending of waves around an
  opening or around the edge of an object
• Interference the ability of two or more waves
  to add together forming regions of large or small
  amplitude.

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Wave Interference patterns
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Light Exhibits Interference

• Constructive interference waves in-phase
  create waves of greater amplitude ( they add)
• Destructive interference waves out-of-phase
  create waves of lower amplitude (they cancel out)

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The Electromagnetic Spectrum

• Visible light is one type of electromagnetic
  radiation.
• The arrangement of electromagnetic waves by their
  wavelength is called the electromagnetic spectrum
• The spectrum ranges from high energy, shorter
  wavelengths of radiation to long wavelength,
  lower energy radiation.
• The visible portion of the spectrum is very
  small.

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The Wave Nature of Light
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? Ionizing Radiation ?
? Nonionizing Radiation ?
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X-rays
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Units for Wave Characteristics

• Wavelength (? lambda) meters with a Greek
  prefix (nanometers , nm)
• Frequency (? nu) cycles per second or Hertz
• (1/s, sec-1, Hz)
• Speed meters per second
• (m/s)

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Wave Calculations

• All electromagnetic radiation travels at the same
  speed through the vacuum of space
• 3.00 x 108 m/s (c)
• Wavelength and frequency are inversely
  proportional ? c/?
• Energy and frequency are directly proportional
• E h?
• E energy in joules,
• h Planks constant 6.626 ? 10-34 Js

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The Dual Nature of Light

• Light behaves as both a wave and as a particle.
• Evidence
• 1) The glowing of heated metals
• (first infra red, then visible as Tº?)
• 2) The photoelectric effect

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The Work of Max Planck(1858-1947)

• Able to predict the wavelengths of light changes
  with temperature
• Energy must be emitted in a Quantized way, or
  restricted to certain quantities.
• Quantum (singular) or Quanta (Plural)
• A quantum is a packet of energy
• VERY SMALL
• Related through Plancks Constant
• h 6.626 ? 10-34 Js, E h?

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The Photoelectric Effect

• When light is shined on a piece of metal,
  electrons are ejected from the metal
• Only light containing enough energy (of a certain
  wavelength and frequency) works

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The Photoelectric Effect
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Einstein and the Photoelectric effect

• Albert Einstein (1879-1955) proposes that light
  has a particle nature too.
• He relates Plancks idea of quantized energy to
  light.
• Light energy quanta photons
• Photons transfer energy to electrons when light
  strikes the metal.
• Photons must be of sufficient energy for this to
  occur.

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What do you see?
Depending on how you look at this it can be an
old lady or a young lady turning her head. The
picture has a dual nature
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Line Spectra of Elements

• A line spectrum is a spectrum that contains only
  certain colors, or wavelengths of light.
• The rainbow is a continuous spectrum
• Elements emit line spectra when they are
  vaporized in an intense flame or with electricity.

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Continuous spectrum of white light
Line spectra of elements
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Line Spectra and the Quantization of Energy
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Neils Bohr (1885-1962)Planetary Model of the Atom

• Combines Rutherfords nuclear model and Plancks
  quanta of energy.
• Electrons must have certain energy levels in
  which they travel. (Energy Level Postulate)
• Energy of Electrons must be quantized to explain
  line spectra.
• (Transitions Between Energy Levels)

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Bohr Model

• Ground state energy level closest to the
  nucleus
• Excited state energy levels farther away from
  the nucleus
• Energy levels represented by quantum numbers ex.
  n 1, n 2, n 3 etc.

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Emission Spectrum of Hydrogen

• In general, the line spectrum of an element is
• rather complicated
• The line spectrum of hydrogen, with a single
  electron,
• is the simplest

http//upload.wikimedia.org/wikipedia/commons/4/4c
/Emission_spectrum-H.png
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The Evidence of Line Sprectra

• Atomic line spectra tell us that when an excited
  atom loses energy, not just any arbitrary amount
  can be lost
• This is possible if the electron is restricted to
  certain energy levels
• The energy of the electron is said to be
  quantized

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• Radiation absorbed electron moves from the
  ground state to an excited state.
• Cant maintain this higher energy level
• Electron falls back down to the ground state
• Radiation is emitted as it returns to the
  ground state
• The energy emitted or absorbed ? energy levels

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Emission of Light During Movement of Electrons
http//iws.collin.edu/biopage/faculty/mcculloch/14
06/outlines/chapter2010/Ma10-8b.JPG
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• J.J. Balmer equation for visible spectrum of
  Hydrogen
• v 1/? 1.097 x 107/m (1/n12 - 1/n22)
• Which led to..
• The Rydberg equation can be used to calculate all
  the spectral lines of hydrogen

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Rydberg Equation
Rydberg constant 109,678 cm-1

• Used to calculate the wavelengths of all the
  spectral lines of hydrogen.
• Atomic spectra indicate that when an excited atom
  loses energy, the energy is in discrete amounts
  - or quantized.
• n1 and n2 are positive integers

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• Bohr proposed that the electrons moved around the
  nucleus is fixed paths or orbits much like the
  planets move around the sun
• The orbits, labeled with the integer n, have
  energy
• This equation allows the calculation of the
  energy of any orbit

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Bohr was able to use this model to calculate the
energies of the light given off by the hydrogen.
E h?
Unfortunately, the model became too
mathematically complicated for any element with
more electrons than hydrogen. Also, electrons not
completely explained as particles
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Limitations of the Bohr Model
It cannot explain the spectra of atoms other
than hydrogen. Electrons do not move about the
nucleus in circular orbits. However - the
model introduces two important ideas The
energy of an electron is quantized electrons
exist only in certain energy levels described by
quantum numbers. Energy gain or loss is
involved in moving an electron from one energy
level to another.
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Matter Waves

• Louis DeBroglie (1892-1897)
• Electrons (matter) have a dual nature just like
  light.
• Proved it mathematically
• Evidence diffraction patterns produced by beams
  of electrons

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Electron Diffraction Patterns
http//www.microscopy.ethz.ch/TEM_ED_examples.htm
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Matter Wave Equations
E mc2 and E h? ? c/? (solve
for ?) ? c/ ? (Replace ? with c/ ?) E
h c/ ? (Both sides of the equation equal E so
you can equate one to the other) hc/ ?
mc2 (solve for ?) ? hc h mc2
mc (replace the speed of light with the speed of
the particle) ? h ms You can calculate
the wavelength of any object!
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? h ms Why do we not see these
waves? Because for relatively large objects
their mass is too large and their speed too slow
for the wavelengths to be observed. Remember, h
6.626 x 10-34 J s Electrons - small mass
(9.109 x 10-28 g) - travel REALLY fast so we
can observe the wave behaviors
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To be Continued