Atomic Physics

1
Atomic Physics

• Chapter 28

2
Atomic Models
3
Introduction

• How do neon signs work?

4

• Our main focus will be on the hydrogen atom.
• It is the simplestatomic system.

5

• Why is it important to study the hydrogen atom?
• Studying the quantum numbers for the allowed
  states of hydrogen will help us to describe the
  allowed states of more complex atoms.
• The hydrogen atom is an ideal system for relating
  theory to experimentation.
• Much that we learn about hydrogen can be extended
  to single electron ions like He and Li.

6
Early Models Of The Atom

• The Greek model
• Tiny, hard, indestructible sphere
• 3

7

• The J. J. Thomson model
• A volume of positive charge is embedded with
  negative charges called electrons

8

• The Rutherford model
• A positive nucleus orbited by electrons.
• The nucleus contains 99.9 of the atoms mass

9

• The Rutherford model
• Which force holds the electrons in orbit?
• The Coulomb force

10
Problems with the Rutherford Model

• There were two basic difficulties with the
  Rutherford model.
• It could not explain why atoms radiate discrete
  frequencies.
• Accelerating electrons should radiate
  electromagnetic waves.

11
Electron Transitions

• Using a high voltage to move electrons through a
  gas causes the gas electrons to become excited
  and to jump from lower energy levels to higher
  energy levels.
• Photons of various wavelengths are produced when
  electrons fall from higher energy levels to lower
  energy levels.

12
Emission Spectra

• The emission spectrum of hydrogen
• Can be produced by applying a high voltage across
  an evacuated glass tube filled with hydrogen
• The observed wavelengths are characteristic only
  of hydrogen
• 279, 57

13
The Balmer Series

• In the Balmer Series
• nf 2
• There are four prominent wavelengths
• 656.3 nm (red)
• 486.1 nm (green)
• 434.1 nm (purple)
• 410.2 nm (deep violet)
• 278, 28.7

14
Balmer Wavelengths
15
The Balmer Series Wavelength Equation

• RH is the Rydberg constant
• RH 1.0973732 x 107 m-1

16
Two Other Important Series

• Lyman series (UV)
• nf 1
• Paschen series (IR)
• nf 3
• 70

17
Spectral Lines

• How many different spectral lines could be
  produced by an electron in the n 3 state?
• Three

18

• How many different spectral lines could be
  produced by an electron in the n 4 state?
• Six

19
Photon Energy

• The equation for determining the energy of the
  emitted photon in any series

20
The Absorption Spectrum

• An element can absorb the same wavelengths that
  it emits.
• The spectrum consists of a series of dark lines.

21
Identifying Elements

• The absorption spectrum was used to identify
  elements in the solar atmosphere were identified
  in this way.
• Helium was discovered.

22
Thermal vs. Atomic Spectra

• How could you tell if the light from a candle
  flame is thermal or atomic in origin?

23

• If the spectrum is continuous, the source must be
  thermal.

24
Auroras

• What is the origin of the colors in the aurora
  borealis?

25

• High speed particles from space interact with the
  earths magnetic field.

26
The Bohr Theory Of Hydrogen

• At the beginning of the 20th century, scientists
  wondered why atoms only radiated certain
  wavelengths.
• Bohr provided an explanation.

27
Four Assumptions of The Bohr Theory

• 1) The electron orbits the proton due to the
• Coulomb force which produces centripetal
• acceleration.

28

• 2) Only certain electron orbits are stable
• and do not radiate energy.

29

• 3) Radiation is only emitted when an
• electron drops from a more energetic
• state to a lower state.

30

• 4) The radius of the electrons orbit is
• determined by the electrons orbital
• angular momentum.
• 28.6

31
Total Energy of the Hydrogen Atom

• The total energy of the hydrogen atom can be
  determined by using this equation.

32
The Bohr Radius

• An electron can exist only in certain allowed
  orbits determined by the integer n.
• When n 1, we have what is known as the Bohr
  radius (ao).
• ao 0.0529 nm

33
Orbital Radii

• A general equation for finding the radius of any
  orbit

34
Energy States

• The energy for various energy states can be found
  by using
• n 1 is the ground state

35
Ionization Energy

• The minimum energy required to ionize the atom is
  called the ionization energy.
• An electron is completely removed from the atom.

36
The Hydrogen Spectrum

• The general expression for determining
  wavelengths of the various series in the hydrogen
  spectrum

37
Bohrs Correspondence Principle

• Quantum mechanics is in agreement with classical
  physics when the energy differences between
  quantized levels are very small.

38
Successes of the Bohr Theory

• It accounted for the Balmer series and other
  series.

39

• It predicted a value for the Rydberg constant
  that agreed strongly with the experimental value.

40

• It gave an expression for the radius of the
  hydrogen atom.

41

• It predicted the energy levels of hydrogen.

42

• It also works with hydrogen-like (one electron)
  atoms.
• Singly ionized helium

43

• It also works with hydrogen-like (one electron)
  atoms.
• Doubly ionized lithium

44

• It also works with hydrogen-like (one electron)
  atoms.
• Triply ionized beryllium

45
Four Quantum Numbers

• The state of an electron is specified by four
  quantum numbers.
• These numbers describe all possible electron
  states.
• The total number of electrons in a particular
  energy level is given by

46
Principle Quantum Number

• The principal quantum number (n) where n 1, 2,
  3,
• Determines the energy of the allowed states of
  hydrogen
• States with the same principal quantum number are
  said to form a shell
• K, L, M, (n 1, 2, 3, )

47
Orbital Quantum Number

• The orbital quantum number (l) where l ranges
  from 0 to (n 1) in integral steps
• Allows multiple orbits within the same energy
  level
• Determines the shape of the orbits
• States with given values of n and l are called
  subshells
• s (l 0), p (l 1), d (l 2), f (l 3), etc

48
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49
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50
Electron Subshells
51

• Generally, the electrons in the s subshell are at
  the lowest energy level and those in the f
  subshell in the highest shell occupy the highest
  energy level.

52

• As the shell number (n) increases the energy
  difference between the shells diminishes, as
  shown by the decreasing distance between each
  successive shell.

53
Electron Subshells
54
Magnetic Quantum Number

• The magnetic quantum number (ml) where ml ranges
  from - l to l in integral steps
• Explains why strong magnetic fields can cause
  single spectral lines to split into several
  closely spaced lines
• Called the Zeeman effect

55
Spin Magnetic Quantum Number

• The spin magnetic quantum number (ms) where ms
  can only be 0.5 or 0.5
• Accounts for the fine structure of single
  spectral lines in the absence of a magnetic field

56
Hydrogen Like Atoms

• Two important equations for hydrogen-like atoms
• Orbital energy
• Orbital radius

57
Angular Momentum

• Physicists agreed that angular momentum was
  quantized but no one was able to explain why.
• 28.10

58
Electron Standing Waves

• de Broglie stated that an electron orbit would be
  stable if it contained an integral number of
  electron wavelengths.
• Analogous to standing waves in a string

59
Wave Properties

• It became generally agreed upon that wave
  properties were involved in the behavior of
  atomic systems.

60
Quantum Mechanics And The Hydrogen Atom

• A review of the various quantum number ranges
  which are used to determine allowable states
• n can range from 1 to infinity in integral steps
• l can range from 0 to (n - 1) in integral steps
• ml can range from l to l in integral steps
• ms can only be ½ or ½

61
The Spin Magnetic Quantum Number

• The spin magnetic quantum number explains the
  splitting of each energy level into two (the
  Zeeman Effect).
• It explains how two very closely spaced lines may
  be formed in the spectra of certain gases.
• Electron spin (spin-up and spin-down)

62
Questions

• 2, 8, 12
• Pg. 910

63
Electron Clouds

• The electron may be found at various distances
  from the nucleus but the probability of finding
  it at a distance corresponding to the first Bohr
  orbit is a maximum.
• It can be found in a spherical region known as
  the electron cloud.
• 281, 282

64
The State of an Electron

• The state of an electron is specified by four
  quantum numbers.
• These numbers describe all possible electron
  states.
• The total number of electrons in a particular
  energy level is given by

65
The Pauli Exclusion Principle

• Two electrons in an atom can never have the same
  set of quantum numbers.
• Because of this, the elements all have different
  chemical properties.
• The n 1 energy level is filled with electrons
  first.

66
The Pauli Exclusion Principle And The Periodic
Table

• Mendeleev arranged the elements in a periodic
  table according to their atomic masses and
  chemical similarities.
• He left gaps which were filled in within the next
  20 years.
• Vertical columns have similar chemical
  properties.
• 15

67
The Periodic Table
68
Special Groups Within the Periodic Table

• Noble gases
• The outer shell is filled.
• Alkali metals
• The outer shell has only one electron.
• Halogens
• The outer shell needs one electron.

69
The Dow Corning Periodic Table
70
X-Rays

• X-rays are emitted when a metal target is
  bombarded with high-energy electrons to produce
• A broad continuous band
• Bremsstrahlung
• Characteristic x-rays
• K??and K?
• 284, 285

71
X-Ray Photons

• What can the incoming electron from an electron
  gun do to a K-shell electron in a tungsten target
  atom?
• It can knock a K-shell electron out of its energy
  level. Then, an electron from a higher energy
  level can fall into the K-shell (n 1).
• The energy lost by the falling electron shows up
  as an emitted x-ray photon.

72
Characteristic X-Rays

• K-shell emission produces higher-intensity x-rays
  than Bremsstrahlung.
• The x-ray photon comes out at a single
  (characteristic) wavelength.
• K??or K?

73
Ka X-Rays

• When an incoming electron forces an electron out
  of the K shell an electron can drop down from the
  n 2 level and a Ka x-ray photon is emitted.

74
Kb X-Rays

• When an incoming electron forces an electron out
  of the K shell an electron can drop down from the
  n 3 level and aKb x-ray photon is emitted.

75

• Which x-ray photon has the highest energy?

76
Ka X-Ray Wavelengths

• The wavelength of the emitted Ka x-ray
  photon is given by

77
Electron Shielding

• One electron in the K shell partially shields the
  other from the charge of the nucleus.
• Because of this, we use Zeff (Z - 1) in the Ka
  equation.

78
K? X-Ray Wavelengths

• The wavelength of the emitted K? x-ray
  photon is given by

79
Electron Shielding

• One electron in the K shell and eight electrons
  in the L shell partially shield the M-shell
  electrons from the charge of the nucleus.
• Because of this, we use Zeff (Z - 9) in the K?
  equation.

80
Atomic Transitions

• Atoms will only emit or absorb EM radiation at
  certain frequencies corresponding to transitions
  involving the various energy states.

81
Stimulated Absorption

• In the stimulated absorption process, light may
  be used to stimulate electrons to higher excited
  states.
• Only certain frequencies will do this.
• 28.17

82
Spontaneous Emission

• When the electrons randomly fall back to their
  original orbits we call this spontaneous
  emission.
• 286

83
Spontaneous Emission
84
Stimulated Emission

• In stimulated emission, all of the electrons can
  be made to fall back at the same time and thus
  produce bright, coherent light.
• This is the basis for the operation of LASERS.

85
Stimulated Emission
86
Lasers

• LASER - Light Amplification by Stimulated
  Emission of Radiation

87
Population Inversion

• In a laser, electrons are stimulated so that
  there are more electrons in the excited state
  than in the ground state.
• This is called a population inversion.
• 287

88
Laser Requirements

• There are three conditions for laser action to
  occur.
• A population inversion
• The excited state must be a metastable (long
  lifetime) state.
• The photons must be confined long enough to
  stimulate further emissions.

89
He-Ne Lasers

• The operation of a He-Ne laser
• An oscillator is used to sweep electrons through
  a thin glass tube containing a He-Ne mixture.
• The neon atoms are raised to a metastable state
  by collisions with excited helium atoms.
• Electrons simultaneously returning to a lower
  energy state emit coherent photons of a
  particular wavelength. (632.8 nm)
• 28.22a, 71, 288

90
Laser Frequencies

• Frequency ranges of lasers
• Infrared (CO2)
• Visible (red, green, blue)
• Ultraviolet

91
Laser Applications

• Medical
• Welding detached retinas
• Laser surgery
• Laser vision correction (Lasik)

92
Lasik Surgery

• An ultra-thin flap is created on the eye's
  surface during LASIK corrective eye surgery.
  After laser energy is applied to reshape the eye,
  the flap is replaced to serve as a type of
  natural bandage.

93

• Surveying and distance measurement

94

• Cutting and drilling metals in industry

95

• Fiber optic communications

96
Holography

• Used in the production of three-dimensional
  images
• Interference patterns are placed on film.
• Used to protect credit cards
• 283, 284

97
Making Holograms
98
CDs and DVDs

• Information is stored in binary form.
• Pits and land areas (ones and zeros)
• The laser beam follows a spiral path.
• A diffraction grating is used to provide
  tracking.
• 40 second memory for music CDs

99
Infrared Remote Control

• A different infrared wavelength is assigned to
  each number or function.
• TV and stereo remote controls use IR.
• Some computers and calculators use IR.
• My MAC PowerPoint remote uses RF.
• Dont confuse IR with RF controls.
• MAC Photo Booth Demo

100
Semiconductor Devices

• Doping
• Donor atoms
• N-type semiconductor
• Acceptor atoms
• P-type semiconductors

101
Semiconductor Devices

• P-N junctions
• Diodes
• Forward bias
• Reverse bias
• Half-wave rectifiers
• Full-wave rectifiers
• Transistors

102
Transistors

• Junction transistors
• Types
• npn
• pnp
• Parts of a transistor
• Emitter
• Base
• Collector
• 227

103
Semiconductor Devices

• Integrated circuits
• What are they?
• Where are they used?
• What are the advantages of integrated circuits?

104
Computer Memory
105
Questions

• 7, 9 - 11, 15
• Pg. 910