1. Atomic Theory and Spectroscopy

2. Electromagnetic Radiation
Energy emitted by electrons can be detected at any part of the electromagnetic spectrum

3. Electromagnetic Radiation
So, just what is EMR?
- an oscillating electric and magnetic field
which travels through space OR
- a discrete series of “particles” that possess a
specific energy but have no mass
BOTH!

4. Waves

5. Measuring Waves Two properties can be measured: Wavelength (l)
The distance from the same point on successive waves (measured in meters)

6. Measuring Waves Frequency ()
The number of times a wave travels up and down per second
Measured in cycles per second or hertz (Hz)

7. The frequency () is the number of wave crests per second which pass a reference point.

8. The amplitude (A) is the height of the wave

9. Measuring radiation

10. Measuring Radiation
All radiation constantly travels through space at the same velocity (speed) =
3.0 x 108 m/s
(299,792,458 meters per second)
The speed of electromagnetic radiation ---“The speed of light” = c

11. c=  l
c = 3.0 x108
If either the frequency or wavelength is known, the other can be calculated

12. A red light has a wavelength of 728 *10-9 m. 
What is the speed of the wave in m/s?   
What is the frequency of the light?

13. A certain blue light has a frequency of 6. 91 x 1014 Hz
A certain blue light has a frequency of 6.91 x 1014 Hz. What is the wavelength of the light?

14. Determine the wave length of light with a speed of 50*106 Hz (/sec)

15. If I have a wavelength of 780 *10-9 m and what is the frequency?

16. Microwave ovens often employ radiation with a frequency of 2
Microwave ovens often employ radiation with a frequency of 2.45 x 109 /s. What is the wavelength (in cm) of this radiation?

17. A purple light has a frequency of 7. 42 x 1014 Hz
A purple light has a frequency of 7.42 x 1014 Hz.      What is its wavelength?     What is the energy of one quanta of light

18. m/s
Not just a good idea,
it’s the law!

19. Properties of Light
Form of energy, detectable with the eye, which can be transmitted from one place to another at a finite velocity

20. Theory of Light •Two complimentary theories to explain
how light behaves and the form by
which it travels:
Particle Theory
Wave Theory

21. Particle Theory
• Release of a photon

22. Electromagnetic spectrum and Energy
longest

23. Planck’s Quantum Theory
In 1900, German Physicist Max Planck proposed: “Radiant energy may only be absorbed or emitted in discrete amounts: quanta.”

24. Electromagnetic spectrum and Energy
Planck discovered the energy of a wave or photon of light is constant
h = (6.63 x 10-34J/Hz)
Planck’s constant (h)

25. Energy and Radiation
If the frequency of a wave is known, then the energy of the wave can be calculated as well in the same manner
E=h  Or
E= h (c/)

26. Energy and Radiation  = 1.035 x 108 Hz E=hv
Sample: What is the energy of a wave which has a frequency of x 108 Hz?
6.63 x10-34 J
1.035 x 108 Hz
E =
6.86 x 10-26J = Hz
 = x 108 Hz
6.63 x10-34 J
h = Hz

27. Photons
Photons are a particle of radiation or an individual quantum of light
Quantum is a finite quantity of energy that can be gained or lost by an atom

28. The Nature of Light and Radiation
If e- are exposed to energy which matches those levels (a photon), they leap to unstable higher energy levels.
As they fall back they emit that energy ( a photon) at a wavelength and frequency which can be detected.

29. Photoelectric Effect
Albert Einstein proposed that light not only behaves as a wave, but as a particle too
Albert Einstein did not get the Nobel Prize for Relativity. He received it for work that he did between 1905 and 1911 on the Photoelectric Effect.

30. Photoelectric Effect In 1905, Einstein applied this quantum theory to
explain the photoelectric effect:

32. Photoelectric Effect
-if EMR was absorbed as a wave, then the number of electrons ejected and the energy of the electrons ejected should vary only with the intensity of the light

33. Einstein Photoelectric Effecthv
Ehv - Ee- = Ionization Energy

34. Photoelectric Effect NUMBER of e-: does vary with EMR intensity
ENERGY of e-: vary only with MR frequency AND: no effect if freq is below a threshold value!

35. Photoelectric effect
The reason that certain types of light cause this effect but others do not has to do with threshold energy
Remember that electrons must gain energy in certain amounts (E = h)
Only certain types of light with “just the right” frequency can cause electrons to become excited

37. The Nature of Light and Radiation

38. Photoelectric Effect View EMR as a collection of particles(called
photons), with each photon having the following energy:
E = hν

39. Photoelectric Effect Each photon will cause an electron to be ejected
IF the energy of the photon is above a minimum (threshold) value. Any energy of the photon above that needed to eject the electron would be transferred to the electron as kinetic energy

40. Photoelectric Effect
Increased EMR intensity translates to an increase in the number of photons (increasing the number of electrons ejected)

41. Einstein
Einstein did not receive his Nobel Prize of the theory of relativity but for his work on the photoelectric effect

42. Typed of Spectra

43. Hydrogen Line Spectrum
Because hydrogen atoms emit only specific frequencies of light indicate that the energy differenced between the atom states are fixed

44. The Bohr Atom
The electron in a hydrogen atom can exist only in discrete orbits
The orbits are circular paths about the nucleus at varying radii

45. Light leads to Electrons
Electrons exist in specific, discrete levels or quanta of energy
The lowest energy of each electron is known as its ground state
If the right amount (quantum) of energy is applied, then an electron will “leap” to another energy level

46. The Bohr Atom Each orbit corresponds to a particular energy
Orbit energies increase with increasing radii

47. The Bohr Atom The lowest energy orbit is called the ground state
After absorbing energy, the e- jumps to a higher energy orbit (an excited state)

48. The Bohr Atom
When the e- drops down to a lower energy orbit, the energy lost can be given off as a quantum of light
The energy of the photon emitted is equal to the difference in energies of the two orbits involved

49. Bohr’s Model of H Atom Linked an atom’s e- to photon emission
Said that e’ can circle the nucleus only is specific paths or orbits.
Orbits are separated from the nucleus by large empty spaces (nodes) where the e- cannot exist. (Rungs on a ladder)
Remember, E increases as e- are further from the nucleus.

50. Bohr’s Model of H Atom
While an e- is in orbit, it cannot gain nor lose E
It can, move to a higher E level IF it gains the amount of E = to the difference between the higher E orbit and the lower energy orbit.
When falls, it emits a PHOTON = to the difference.
Absorption gains E/Emission gives off E

51. Take the good with the bad (Bohr’s Model)
Led scientists to believe a similar model could be applied to all atoms.
Bad:
Yikes!!! Did not explain the spectra of atoms with more than one e’ or chemical behavior of atoms!

52. The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much like planets orbiting the sun.
But I was wrong! They’re more like bees around a hive.
WRONG!!!
Neils Bohr

53. The Nature of Electrons
If light acts as both particles and waves, then so do electrons.
If electrons act like waves, then how can we locate them?

54. Atoms are like onions and ogres they have lots of layers
plantanswers.tamu.edu

55. Orbits
We like to draw orbits as circular
but they really aren’t

56. Orbits
Turns out, there's no reason to assume that electron orbits are circular.

57. Orbits
There are lots of paths an electron can take in order to get around the nucleus.
(or through the nucleus)

58. Orbits
In fact it's very rare for an atom's electron to be in a circular orbit.

59. Electron Facts
The electron moves at different speeds. Fast near the nucleus and slow when it's far from the nucleus.

60. Well we really don’t Because atoms are so small, no one can see them.
For this reason, we just can't say exactly where the electron is, as it moves about the nucleus. In those orbit pictures, we know the electron will be somewhere on the white orbit lines.

61. This makes the models involving orbits, whether circles, or ellipses, wrong, because the orbits are pretty specific about the electron's location.

62. So how do we locate electrons in the modern atom?
Quantum Numbers

63. Quantum Numbers
Specify the properties of atomic orbitals and the properties of electrons in orbital

64. Quantum Numbers
Don’t tell us where the electron is, just where it's most likely to be.
The probable location of electrons are described using an address

65. The four quantum numbers
their symbols are n, l, m and s.
EVERY electron in an atom has a specific,
unique set of these four quantum numbers.

66. Quantum Numbers There are 4 parts to each address
Principle quantum number (n)
Angular quantum number (l)
Magnetic quantum number (m)
Spin quantum number (s)

67. The Principal Quantum Number
The main energy level of an electron

68. Principal Quantum Number (n)
Describes the size of the orbital.
Since the distance from of an electron from the nucleus is directly proportional to the energy of the electron
It must be a whole number n=1, n=2 …

69. Angular quantum number (l) Azimunthal
Describes the shape
The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels).
s p d f g h...

70. Angular (Azimuthal) Momentum Quantum Number
Letter
Shape s
Sphere p
Dumbell d
Cloverleaf f
“funky”

71. Each suborbital can hold a maximum of 2 electrons per orbital

72. Angular (Azimuthal) Momentum Quantum Number
Letter
Max number of suborbitals
Max. # of e- s 1 2 p 3 6 d 5 10 f 7 14

73. Quantum Numbers
The more crowded the dots are in a particular region, the better chance you have to finding your electron there.

74.                     
s Orbital sphere

75. Sizes of s orbitals Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.

76.                     
p Orbital dumbbell

77.                     
two p Orbitals

78.                     
All three p Orbitals

79. orbital
                       
d Orbitals Cloverleaf

80. orbital
                       
f Orbitals “funky”
                                                                                                      

82. Orbital filling table

83. The Magnetic Quantum Number
describes the orientation in space of a particular orbital.
x, y, z
It is called the magnetic quantum number because the effect of different orientations of orbitals was first observed in the presence of a magnetic field.)

84. Magnetic quantum number (m)
The orientation in space
What axis the orbit is located on

85. The Spin Quantum Number
allows two electrons of opposite spin (or symmetry) into each orbital.
+1/2 or -1/2
 or 

86. Spin quantum number (s)
+1/2
-1/2
       
               

87. Aufbau (filling up) Principle
the number of electrons in an atom is equal to the atomic number
each added electron will enter the orbitals in the order of increasing energy
an orbital can only hold 2 electrons.

88. Hund’s Rule
For orbitals of equal energy, one electron goes into each orbital until all orbitals are half-full.
2p4
2s2
1s2

89. Aufbau Principle
Lower-energy orbitals of an atom are filled with electrons first.
1s 2s 2p 3s 3p 4s 3d 4p …

90. Pauli Exclusion Principle
No two electrons in an atom can have an identical set of four quantum numbers. (electrons sharing an orbital have different spins)
2s2
1s2

91. Key Terms
Ground State Electron Configurations (1s22s2…/Orbital Notations __ __ __ __
Aufbau Principle (lowest first/periodic guide)
Hund’s Rule: Single e’ before pairing begins
Pauli Exclusion (up/down—no 2 e’ same set of quant. #’s)
Highest Occupied Level
Inner-Shell e’
Noble Gas Notation

92. Spin Diagram for Hydrogen
Electron configuration
for Hydrogen
1s1

93. Spin Diagram for Helium
Electron configuration
for Helium
1s2

94. Spin Diagram for lithium
Electron configuration
for lithium
1s22s1

95. Spin Diagram for beryllium
Electron configuration
for berylium
1s22s2

96. Electron configuration
Spin Diagram for boron
Electron configuration
for boron
1s22s22p1

97. Spin Diagram for carbon
Electron configuration
for carbon
1s22s22p2

98. Spin Diagram for nitrogen
Electron configuration
for nitrogen
1s22s22p3

99. Spin Diagram for oxygen
Electron configuration
for oxygen
1s22s22p4

100. Spin Diagram for fluorine
Electron configuration
for fluorine
1s22s22p5

101. Electron configuration
Spin Diagram for neon
Electron configuration
for neon
1s22s22p6

102. Standard Notation or Complete electron configuration of Fluorine
Number of electrons
in the sub level 2,2,5
1s2 2s2 2p5
Main Energy Level Numbers 1, 2, 2
Sublevels

103. Shorthand Notation or Noble Gas Configuraiton
Use the last noble gas that is located in the periodic table right before the element.
Write the symbol of the noble gas in brackets.
Write the remaining configuration after the brackets.
Ex: Fluorine: [He] 2s2 2p5

104. Blocks in the Periodic Table

107. You broke your big toe!  The x ray they take of  toe uses waves that have a length 2.19 x 1010m.      What is the speed of the wave in m/s?      What is the wavelength in nm?      What is the frequency of the x ray?