1. Waves, light, and energy: Where chemistry and physics collide
EMR info
Waves, light, and energy: Where chemistry and physics collide

2. List as many interactions of light and matter as you can.
Before we get started….
What is light?
Is it matter?
What forms of light exist?
List as many interactions of light and matter as you can.
think how light changes matter, and how matter changes light
What are some uses of light?

3. First things first: Waves
First things first: Waves
a and b represent wavelength (λ)- the distance of a wave from crest to successive crest; measured in meters

4. Waves: amplitude
The height of a wave from crest to midline or trough to midline; measured in meters

5. Terms you need to know: Wavelength (λ) Amplitude
Wavelength (λ)
Amplitude
Frequency (ν [nu]; I know some of you have used f, move on and get with chemistry!):
the number of cycles per second
measured in cycles per second (s-1) or Hz (Hertz)
Waves on a string

10. Visible Light Violet 400---460 7.5--6.5 5.0--4.3
color wavelength(nm) f(*1014 Hz) Energy (*10-19 J)
Violet
Indigo
Blue
Green
Yellow
Orange
Red

12. Some equations you need to know
= c / ν
and
E = hν
So….
E = hc / 
And…
 = h / mv*
When
= wavelength in m
c = speed of light, 3.00E8 m/s
ν (nu)= frequency in Hz
(cycles/sec or s-1 or 1/s)
E= energy in J
h= Plank’s constant, 6.626E-34 J*s [Joule(seconds)]
m= mass of particle in kg
V*= velocity in m/s

13. What the h? Planck’s Constant
What the h? Planck’s Constant
When metals are heated, they glow
1800s- physicists were trying to determine the relationship between the color (wavelength) and intensity of the glow
Max Planck (1900)- energy can be released or absorbed only in little chunks (packets) of energy “of some minimal size”

14. Max Planck and the h
The chunks of energy were dubbed “quantum” (“fixed amount”), which is the smallest amount that can be emitted or absorbed as EMR.
Proposed: E = hν
The energy (E) of a single quantum is equal to its frequency (ν) times a constant

15. Planck and the Nobel (Physics)
Planck determined that h= 6.626E-34 J-s
Energy is always released in multiples of hv (1hv, 2hv, 3hv, etc)
h is so small that we cannot see the effects of this in our daily lives
Analogous to…
Planck won the 1918 Nobel Prize in physics for his work

16. Einstein & Bohr: Perfect Together
Einstein, left
Bohr, above

17. Einstein: The Photoelectric Effect
Einstein discovered that one could cause electrons to be ejected from the surface of a metal if the energy of the light wave was strong enough
He treated the light needed to do this as a piece of matter- a photon, if you will
This ejection of e- is the photoelectric effect

18. The Photoelectric Effect
Only light of a certain energy could knock off an electron from the metal
Intense light of a weaker wavelength would not work, but even a low intensity of the correct wavelength would work
(the energy of the light is transferred to the kinetic energy of the electron)
Hmmm… light acting as a particle and as a wave…..

20. The photoelectric effect…
Online animations
PhET

21. Getting to Bohr….
Light of a given wavelength is monochromatic (one color)
Most common EMR sources are polychromatic, but we see only one color
These can be reduced to a spectrum when the different wavelengths are separated out

22. Spectral Emissions
Continuous spectrum: shows all colors of the rainbow

23. Bright line spectrum: only certain wavelengths are visible (the rest do not appear at all)
Different elements have different bright line spectrum when they are heated
Na is yellow
Ne is orange-red

24. Line spectrum Ne I2

26. Hydrogen Spectra Emission Spectra Absorption Spectra
Emission Spectra
Absorption Spectra

29. Color and what you see:
Absorption: the wavelengths that are absorbed by an object are not available for us to see, as we see the wavelengths of light that are reflected off of an object
This is not the same as those wavelengths that are emitted by an object that is emitting radiant energy.

30. Color and what you see…

31. Chlorophyll absorption spectra

32. Perception of color

33. Line spectra formation- go to…..

34. Bohr Model and Spectral Emissions
Bohr proposed that the emission of light energy from an (electrically or thermally) excited atom corresponds to the orbit of the electron around the nucleus of the atom
That energy can only be achieved by being a specific distance from the nucleus

35. What you’ve seen so far….
Model of a Nitrogen (z=7) atom

36. Bohr Model and moving electrons

37. Energy levels- Bohr Model
Electrons travel within set energy levels that have a particular energy associated with each level
After all, the e-s are moving around the nucleus
think KE here
Each shell has a number
Closest to the nucleus is n=1
For each successive level add 1 to n
n=2, n=3, ect….

38. Energy increases as the distance from the nucleus increases

39. Bohr Model and moving electrons

40. Electron config in energy level

41. SO… We know that the e-’s are free to move around the nucleus
We know that the e-’s are free to move around the nucleus
They also can move from one energy level to the next (and fall) back when energy is added
Move from ground state (“home” level) to a higher level (the “excited” state)
Returning back to the ground state releases energy

42. This emission is how we see colors:
This emission is how we see colors:
the wavelengths of EMR released from an atom when it has been excited by
Heat energy
Electrical energy
Chemical energy
Think glowing red hot metal, or fireworks

43. Determining Energy for n
To determine the energy for a given energy level, use the equation:
En=(-RH)(Z/n2)
Where n=1, 2, 3, 4….
And RH = 2.18E-18J,
So En=(-2.18E-18J)(Z/n2)

44. To determine E emitted or absorbed:
To determine the change in energy for a given energy transition:
ΔE=Ef-Ei
*Remember E=hν, so ΔE=hν
if ΔE=[(-2.18E-18J)(Z/n2)]f- [(-2.18E-18J)(Z/n2)]i

45. E changes continued
*Remember E=hν, so ΔE=hν to get the frequency of the light emitted or absorbed
If ΔE is positive
since Ef >Ei
E is absorbed
The e- was going from ground state to a excited state
If ΔE is negative
since Ef < Ei
E is released
The e- was going from excited state to a ground state

46. Also…life after Einstein and Bohr
We know that electrons have characteristics of both light (waves) and matter, so we say that they have a dual nature

47. De Broglie
De Broglie proposed that an electron moving about the nucleus had a wave-like behavior, so it has a particular wavelength associated with it. This wavelength depends upon the mass and velocity of the electron.
 = h / mv
mv = the momentum of the particle
Mass* velocity = p
momentum = p so p = mv
therefore  = h / p

48. This matter-wave idea applies to all matter, not just to electrons
However, the mass is so large, and the wavelength so small, that we cannot see it in macroscale objects
This matter-wave theory led to applications like the electron microscope

49. De Broglie wavelength

51. Heisenberg: The Uncertainty Principle
We can’t determine information about small scale objects the same way we can for large scale objects
Case in point: a ball rolling down a ramp- we can get position, direction, and speed at the same time
We can’t for electrons
Hence, the uncertainty principle

52. Heisenberg, cont’d
It is inherently impossible for us to simultaneously know both the exact momentum and exact location of an electron
This is because anything we do to determine the location or momentum of the electron moves it from its original path and location; this can’t be reduced past a certain minimal level
We can know only momentum or location- not both
We can talk probability of the location/ momentum of an electron

53. Schrodinger’s Wave Equation..
Continued wave-particle theory
Treated Hydrogen’s electron as a wave
Also worked for other elements’ electrons (not just hydrogen like Bohr’s model).
Known as the quantum mechanical model
Limits electrons to certain energy levels (values)
Does not try to describe an electron’s path around the nucleus

54. Schrodinger’s Wave Equation..
Very complex mathematical equation
Solutions to equations are know as wave functions
Is the probability of finding an electron in a certain space around the nucleus
Predicts a 3 dimensional region known as an atomic orbital
Looks like a fuzzy cloud
Density of cloud is determined by probability of finding the e- there
high probability = dense cloud (dark)

55. Sublevels and Orbitals
The electrons are spread out in orbitals that have varying
Shapes
Energy (distance from nucleus)
The orbitals are described in regards to their quantum numbers
Descriptions that are descriptive and hierarchical
There are 4 numbers that describe an orbital
Written as follows: (#, #, #, ±#)

56. Principal quantum number (n)
The first number (1, #, #,±#)
Describe the
distance from the nucleus of the orbital
The energy of the orbital
Values for n are integers
The smallest possible value is 1
As the distance from the nucleus (and therefore energy) increases, the number increases

57. Quantum numbers

58. There periodic table and n
The 7 periods on the periodic table correspond to n values
Each period has a unique n value
For the 1st period, n=1
For the 2nd period, n=2
And so on….

61. Angular Momentum (l) (this is a script l, as in llama)
Angular Momentum (l) (this is a script l, as in llama)
Is the shape of the sublevel
It is the second number in the description (#,1, #, ±#)
Range from _____________(although we never deal with anything above l=3)
s =0
p =1
d =2
f = 3

62. The s sublevel l = 0

63. p sublevel l = 1

64. d sublevel l = 2

65. Another look at d sublevel

66. f sublevel l = 3

67. General tutorials for electron configuration stuff
some slides in this PowerPoint are from this site already
See key equations and concepts (select from menu on the left), as well as the looking through the overview where to the tutorials are listed (links for just those are on the left, too)

69. Magnetic number (ml) Denote the orbital sublevel that is filled
Denote the orbital sublevel that is filled
It is the third number in the description (#,#,1, ±#)
s sublevesl have one orbital; a sphere has one orientation in space
p sublevels have three orbitals; 3 orientations in space
d sublevels have five orbitals; 5 orientations in space
f sublevels have seven orbitals; 7 orientations in space

70. “Flavors” of orbitals (ml)
s sublevels have one orbital; a sphere has one orientation in space

71. “Flavors” of orbitals (ml)
p sublevels have three orbitals; 3 orientations in space

72. “Flavors” of orbitals (ml)
d sublevels have five orbitals; 5 orientations in space

73. “Flavors” of orbitals (ml)
f sublevels have seven orbitals; 7 orientations in space

74. Spin
Spin is +½ or -½
Up or down (could say clock wise and counter-clock wise)

75. Summary, excluding spin

76. How we use this….
There is a specific order to how the e- fill the orbitals; it is not random
Although there are exceptions to the rules (last thing we do)

77. “The aufbau diagram shows the energy of each sublevel relative to the energy of other sublevels. Each box on the diagram represents an atomic orbital.” Excerpt From: Thandi Buthelezi, Laurel Dingrando, Nicholas Hainen & Cheryl Wistrom. “Chemistry.” McGraw-Hill Education, iBooks.

78. The principles of e- configuration
The Aufbau (next) Principle:
That e- fill the lowest energy sublevel before going to the next sublevel
The Pauli Exclusion Principle:
That e-s are paired according to opposite spins
Hund’s Rule:
e-s spread out in equal energy orbitals before pairing electrons

79. The first level to fill is the 1s level
It is the lowest energy sublevel
It holds two electrons
They are oppositely paired (up and down- ↑↓)
Each sublevel (each __) holds 2 electrons

80. Next… The second sublevel is the 2s sublevel
It also holds 2 electrons (because s holds 2, not because of the number),
also oppositely paired ↑↓

82. 1s2, 2s2,then comes 2p6 So, as it states above
1s fills, 2s fills ,then comes 2p
It holds up to six electrons
Because p orbitals hold 6 electrons

83. Next… From 2p, Notice, you follow the arrows
From 2p,
3s fills with 2e-, then onto
3p, with 6e- then
4s with 2e- followed by
3d with 10e- (because d holds 10e-)
Then 4p with 6e-
Notice, you follow the arrows
Remember, the number of electrons comes from the letter (the orbital’s momentum, m)

84. The sublevels of the orbitals are first filled, then you continue onto the next level (Aufbau)
Also be sure to place one electron in each sublevel prior to filling the level (↑ ↑ ↑ and not ↑↓ ↑ _) (Hund)
e-s must be paired with e-s of opposite spin (↑↓, not ↑↑ or ↓↓) (Pauli)

85. Putting it all together…
Carbon (neutral, so 6 electrons)
What this would look like:
↑↓ ↑↓ ↑ ↑ _
1s 2s 2p
(notice there are 6 arrows for 6 electrons)
This can also be written as 1s2 2s2 2p2
Notice the superscripts add up to 6

87. Nobel gases Last column of P.T. All end the same way with s2 p6
Last column of P.T.
All end the same way with s2 p6
Neon (8 electrons) would look like:
↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s p
Argon (18 electrons) would look like:
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
1s 2s p s p
Can use this trend as a shortcut for writing out long e- configurations

88. Nobel gas shortcut
Non-Nobel gas elements have inner e- configurations identical to certain Nobel gases
Ne is 1s2 2s2 2p6
Na is 1s2 2s2 2p6 3s1
Write Nobel gas in brackets
[Ne] instead of 1s2 2s2 2p6
After the brackets, write the electrons left over
[Ne] 3s1
* When using the shortcut, use the Nobel gas from the period (row) before the element in question*

89. There are some exceptions…
This is because some energy levels are very close together
electrons are able to move between close orbitals in order to minimize repulsion
Example: the 4s and 3d orbitals are very close in energy
So exceptions for some period 4 d block elements occur
Cr is not 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Because it takes less energy to split the electrons between the 5 sublevels than it does to put them together in the 4s and 3d

90. Valence electrons
The electrons in the outermost orbital (usually highest energy level)
Be is 1s2 2s1 has 1 valence electrons (2s1)
N is 1s2 2s2 2p3 has 5 valence electrons (2s2 + 2p3 )
Determine the chemical properties of elements
Used in forming chemical bonds

91. Valence electrons
Can represent valence electrons visually using electron dot structure
Dots stand for valence electrons
Placed one at a time on the four sides of the symbol (they may be placed in any sequence) and then paired up until all are shown