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Chapter%204.1%20Properties%20of%20Light

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Chapter 4.1 Properties of Light Wave-Particle Nature of Light Electrons and light have a dual wave-particle nature. Electromagnetic Radiation (EMR) – PowerPoint PPT presentation

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Title: Chapter%204.1%20Properties%20of%20Light


1
Chapter 4.1 Properties of Light
  • Wave-Particle Nature of Light
  • Electrons and light have a dual wave-particle
    nature.
  • Electromagnetic Radiation (EMR)
  • Form of energy that exhibits wavelike behavior
    and travels at the speed of light.
  • Speed of Light (C) 3 x 10 8 m/s

2
Components of a Wave
  • Wavelength (?) lambda
  • Units any unit of length (m)
  • Distance between corresponding points of a wave.
  • Crest to Crest or Trough to Trough

3
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4
Components of a Wave
  • Frequency (?) nu
  • Units Hertz (Hz) or 1/s
  • How often a wavelength passes a given point in
    time.

5
Components of a Wave
  • Amplitude
  • Height of the wavelength.
  • Measured from the origin to crest or origin to
    trough.
  • Brightness of light.

6
Wavelength vs. Frequency
  • C ? ?
  • Inversely proportional.
  • As wavelength increases, frequency decreases.

7
  • Spectrums
  • Range of wavelengths for a series of waves.
  • Electromagnetic Spectrum
  • Consist of all electromagnetic radiation.
  • Continuous Spectrum
  • Spectrum where all wavelengths within a given
    range are together.
  • Examples Visible Light, X-Rays, U.V. Light, etc

8
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9
EMR Spectrum
  • 7 Parts
  • Longest wavelength to Shortest
  • Radio
  • Microwaves
  • Infrared
  • Visible Light
  • U.V. Light
  • X-Rays
  • Gamma-Rays

10
  • Problems
  • What is the wavelength of EMR that has a
    frequency of 7.50 x 10 12Hz?

11
  • Problems
  • 1. Determine the frequency of light with a
    wavelength of 4.257 x 10-7 cm.
  • 2. What is the wavelength of U.V light that has a
    frequency of 4.50 x 10 16 Hz?
  • 3. What is the wavelength and color of light,
    that has a frequency of 6.00 x 10 14 KHz?

12
Chapter 4.2 Quantum Theory
  • Photoelectric Effect
  • Emission of electrons by certain metals when
    sufficient light shines on them.
  • Photoelectric effect

13
Chapter 4.2 Quantum Theory
  • Photoelectric Effect

14
Chapter 4.2 Quantum Theory
  • Photoelectric Effect

15
Chapter 4.2 Quantum Theory
  • Quantum
  • Finite quantity of energy that can be gained or
    lost by an atom.
  • Plancks Equation
  • h 6.63 x 10 34 Js
  • E quantum of energy
  • Photon
  • An individual quantum of light, caused by
    electrons losing quanta of energy.

E h ?
16
Chapter 4.2 Quantum Theory
  • Visible Light Emissions
  • As electrons gain quanta of energy they release
    it in the form of photons.
  • Energy States of an Atom
  • Ground State- an atoms lowest energy level.
  • Excited State- an atoms highest energy level.
  • , is produced when
    electrons drop from the excited to the ground
    states.

Line Spectrum
17
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18
Chapter 4.2 Quantum Theory
19
Chapter 4.2 Quantum Theory
  • Problems
  • What is the energy of U.V. light with a frequency
    of 4.50 x 10 16 Hz?
  • E h ?, h 6.63 x 10 34 Js

20
Chapter 4.2 Quantum Theory
  • Problems
  • Determine the energy of light that has a
    wavelength of 450nm.

21
  • Equations

22
  • Problems
  • 1. What is the energy of a photon of green light
    with a frequency of 5.80 x 1014 1/s?
  • 2. What is the energy, in joules, of a quantum of
    radiant energy whose wavelength is 6.82 x 10 6
    cm?
  • 3. Determine the wavelength of a photon that has
    3.11 x 10 19 J of energy.
  • 4. Determine the frequency, in MHz, of a photon
    that has wavelength of 1.36 x 10 10 nm.

23
Summary
  • restricted the amount of energy
    that an object emits or absorbs as a quantum. (E
    h? ?)
  • used Plancks theory and
    explained the photoelectric effect.
  • light travels as tiny particles,
    photons.

Planck
Einstein
Compton
24
Bohrs Model
  • The Line Spectra demonstrates that the energy
    levels of an electron in an atom are quantized
  • Similar to the rungs of a ladder, nothing exist
    in between.
  • (For Hydrogen (1 p 1 e- )
  • 1st Energy Level n 1
  • 2nd so on n 2,3,4,5,6, 8
  • Only electrons dropping from a Higher Level to a
    Lower one emit EMR
  • A Number of Possibilities for electron drops

25
Hydrogens Line Spectrum
  • Several Series of lines are observed
  • Electron Drops to the n 1 Level
  • Lyman Series (U.V. Range)
  • Electron Drops to the n 2 Level
  • Balmer Series (Visible Range)
  • Electron Drops to the n 3 Level
  • Paschen Series (Infrared Range)

26
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27
  • The Lines become more closely spaced as the
    levels increase
  • Theres an upper limit to how high the electron
    can jump.
  • The Bohr model explained spectral lines but not
    how atoms bonded.
  • Ultimately Displaced

28
1924 Louis de Brogile
  • French Graduate Student (asked an important
    question)
  • If light behaves as waves particles, can
    particles of matter behave as waves?
  • Derived an Equation
  • Predicts that all matter exhibits wavelike
    motions.

h Planks Con. m mass v - velocity
29
  • Large Objects
  • Small Wavelengths
  • 200 g Baseball _at_ 30 m/s - ? 10-32 cm
  • Undetectable
  • Small Objects Large Wavelengths
  • 9.11 x 10-28 g _at_ 30 m/s - ? 10-3 cm
  • Very Detectable w/ proper instruments
  • New Ballgame Classical Mechanics vs. Quantum
    Mechanics
  • New method for describing the motions of
    subatomic particles

30
Heisenbergs Uncertainty Principle
  • It is impossible to know exactly both the
    velocity the position of a particle at the same
    time.
  • Accuracy of V ? then Position ?

31
Classical Vs Quantum
  • Classical adequately describes the motions of
    bodies much larger than the atoms of which they
    are composed. It appears that such a body loses
    energy in any amount
  • Quantum describes the motions of subatomic
    particles and atoms as waves. These particles
    gain or lose energy in packages called quanta.

32
Quantum Mechanical Model
  • Modern description of the electrons derived from
    the mathematical solution to the Schrodinger
    equation.
  • Erwin Schrodinger - used wave mechanics to show
    the electrons about the nucleus emit vibration
    frequencies that were constant.
  • Quantum Numbers - specify the properties of
    atomic orbitals and their electrons.
  • distance from the nucleus.
  • ..\..\Videos\phy03_vid_quantum_300.mov

33
Chapter 4.4 Quantum Numbers
  • Principal Quantum Number (n)
  • Main energy level surrounding the nucleus.
  • Size of each orbital.
  • Primary distance from the nucleus.
  • Has values of n 1 to 7, 1 is the closest 7 is
    the farthest from the nucleus.

34
Chapter 4.4 Quantum Numbers
  • Orbital Quantum Number (l)
  • Shape of the orbitals.
  • Referred to as subshells.

35
Chapter 4.4 Quantum Numbers
p orbital
s orbital
d orbital
f orbital
36
Chapter 4.4 Quantum Numbers
  • Magnetic Quantum Number (m)
  • Orientation of an orbital about the nucleus.
  • l s
  • m 0
  • l p
  • m -1, 0, 1
  • l d
  • m -2, -1, 0, 1, 2
  • l f
  • m -3, -2, -1, 0, 1, 2, 3

37
Chapter 4.4 Quantum Numbers
  • s orbital, 1 orientation.

38
Chapter 4.4 Quantum Numbers
  • p orbital, 3 orientations.

px orbital
py orbital
pz orbital
pxyz orbital
39
Chapter 4.4 Quantum Numbers
  • d orbital, 5 orientations.

40
Chapter 4.4 Quantum Numbers
  • f orbital, 7 orientations.

41
Chapter 4.4 Quantum Numbers
  • Spin Quantum Number(1/2 , -1/2)
  • Indicates two possible states on an electron in
    an orbital.

42
Chapter 4.4 Quantum Numbers
  • Magnetism
  • Caused by the motion of electrons about the
    nuclei of atoms.
  • Diamagnetism substance is weakly repelled by a
    magnetic force.
  • Paramagnetism substance is weakly attracted by
    a magnetic force.
  • Ferromagnetism Strong attraction by a magnetic
    force.

43
Chapter 4.4 Quantum Numbers
Principal Energy Level Sublevels Orbitals
N1 1s 1s
N2 2s, 2p 2s(one) 2p(three)
N3 3s, 3p, 3d 3s(one) 3p(three) 3d (five)
N4 4s, 4p, 4d, 4f 4s(one) 4p(three) 4d (five) 4f(seven)
44
Chapter 4.4 Quantum Numbers
Principal Q.N. Orbitals per Main level (n2) Electrons per main level (2n2)
1 1 2
2 4 8
3 9 18
4 16 32
45
Chapter 4.4 Quantum Numbers
Orbital Max electrons
s 2
p 6
d 10
f 14
46
Rules Governing Electron Configurations
  • Electron Configuration arrangement of electrons
    in the atom
  • Rules
  • Aufbau Rule electron occupies the lowest energy
    level that will receive it.
  • Hunds Rule orbitals of equal energy each
    receive one electron (equal spin) before any
    receive two.
  • Paulis Exclusion Principle no two electrons
    can have the same set of 4 quantum numbers
    (n,l,m,s)

47
Orbital Notation
  • Orbital Notation
  • Orbital represented by a line ____ or
  • Electron is represented by an ½ Arrow ???
  • ½ (?)
  • - ½ (?)
  • Examples
  • H (1 e-1) He (2 e-1) ?

48
Order of Energy Levels
  • 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s
    5f 6d
  • Number - principal quantum number, the main
    energy level
  • Letter orbital quantum number, the shape
  • Useful Diagram

1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d 7d
4f 5f 6f 7f
49
Orbital Notation
  • Write the orbital notation for the following
    elements
  • Al
  • Zn
  • P
  • Cl Atomic Orbitals

1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d 7d
4f 5f 6f 7f
50
Electron-Configuration Notation
  • Eliminates the lines arrows
  • Superscripts are used to illustrate the number of
    electrons in the sublevel
  • Same order of sublevels
  • Examples
  • H (1 e-1) - 1s1
  • He (2 e-1) - 1s2
  • Li (3 e-1) - 1s22s1

51
Electron-Configuration Notation
  • Write the short-hand electron-configuration for
    the following
  • Br
  • Pb
  • Cs
  • Kr

1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
3d 4d 5d 6d 7d
4f 5f 6f 7f
52
Exceptions to Aufbau
  • All elements prefer a more stable configuration
    of electrons.
  • Fully filled and ½ filled orbitals are more
    stable than others.
  • Elements that are 1 shy of a full or ½ filled d
    orbital configuration will have electrons
    transfer from the s to the d to reach this stable
    state.
  • Example if you have a 4s2 and 3d4 the actual
    configuration should be 4s1 and 3d5.

53
Practice
  • Ba
  • Mg
  • W
  • Ag
  • Sb

54
Noble Gas Configuration
  • Shortest method of writing the electron
    configurations.
  • Use the last noble gas to occur prior to the
    element that is being configured.
  • Start at the nS where n equals the period in
    which the element being configured can be found.
  • Example Zr
  • Noble gas would be Kr and start configuration at
    5s.
  • Kr 5s24d2

55
Writing in noble gas configuration
  • Al
  • Write the noble gas configuration for the
    following
  • V
  • Rb
  • I
  • Hg
  • U
  • W

56
Identifying Electrons
  • Paired electrons when 2 electrons are within
    the same orbital.
  • Unpaired electrons when a single electron is
    within an orbital.

57
Identifying Electrons
  • How many unpaired electrons does the following
    elements have?
  • Na
  • O
  • B

58
Electron Dot Notation
  • Describes the number of electrons in an atoms
    outer electron cloud, or highest energy level, as
    dots around the symbol.
  • Valence electron - refers to the highest energy
    level electron within an atom.

59
Electron Dot Notation
  • Draw the electron-configuration and electron-dot
    notation for Nitrogen.

60
Electron Dot Notation
  • Write the noble gas-configuration and electron
    dot notation for the following.
  • Al
  • As
  • I
  • Sr
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