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