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The Electromagnetic Spectrum and the Model of the Atom Part I

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Title: The Electromagnetic Spectrum and the Model of the Atom Part I


1
The Electromagnetic Spectrum and the Model of the
Atom Part I
  • Chemistry Mrs. Cameron

2
The Purpose of Science
  • The purpose of science to make models that
    explain natural phenomena.

3
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.

4
Radioactivity and Light have been tools to
discover the structure of atoms.
  • Do you Remember?
  • Thompsons Cathode Ray Experiment
  • Rutherfords Gold foil Experiment

5
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6
Thomsons Plum-Pudding Model
7
Ernest Rutherfords Gold Foil Experiment
8
  • 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
9
  • 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

10
New Evidence Continuous and Line Spectra
  • Lets Talk about LIGHT

White light
is actually made up of many colors
11
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.

12
When white light is passed through a prism, the
light refracts, or bends to display all of its
component colors
13
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
14
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.
15
The Wave Nature of Light
16
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.

17
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18
Wave Interference patterns
19
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)

20
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.

21
The Wave Nature of Light
22
? Ionizing Radiation ?
? Nonionizing Radiation ?
23
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24
X-rays
25
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)

26
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

27
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

28
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?

29
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

30
The Photoelectric Effect
31
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.

32
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
33
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.

34
Continuous spectrum of white light
Line spectra of elements
35
Line Spectra and the Quantization of Energy
36
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)

37
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.

38
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39
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
40
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

41
  • 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

42
Emission of Light During Movement of Electrons
http//iws.collin.edu/biopage/faculty/mcculloch/14
06/outlines/chapter2010/Ma10-8b.JPG
43
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44
  • 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

45
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

46
  • 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

47
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
48
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.
49
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

50
Electron Diffraction Patterns
http//www.microscopy.ethz.ch/TEM_ED_examples.htm
51
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!
52
? 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
53
To be Continued
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