1. Chapter 4.1 Radiant Energy
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. Components of a Wave Frequency () nu Units: Hertz (Hz) or 1/s
How often a wavelength passes a given point in time.

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

5. Wavelength vs. Frequency
Inversely proportional.
As wavelength increases, frequency decreases.

6. Chapter 4.1 Radiant Energy
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

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

8. Chapter 4.1 Radiant Energy
Problems
What is the wavelength of EMR that has a frequency of 7.50 x 10 12Hz?

9. Chapter 4.1 Radiant Energy
Problems:
1. Determine the frequency of light with a wavelength of x 10-7 cm.
2. What is the wavelength of U.V light that has a frequency of 4.50 x Hz?
3. What is the wavelength and color of light, that has a frequency of 6.00 x KHz?

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

11. Chapter 4.2 Quantum Theory
Photoelectric Effect

12. Chapter 4.2 Quantum Theory
Photoelectric Effect

13. Chapter 4.2 Quantum Theory
Finite quantity of energy that can be gained or lost by an atom.
Planck’s 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 

14. 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

16. Chapter 4.2 Quantum Theory

17. Chapter 4.2 Quantum Theory
Problems:
What is the energy of U.V. light with a frequency of 4.50 x Hz?

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

19. Chapter 4.2 Quantum Theory
Equations:

20. Chapter 4.2 Quantum Theory
Problems:
1. What is the energy of a photon of green light with a frequency of 5.80 x /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 nm.

21. Summary
– restricted the amount of energy that an object emits or absorbs as a quantum.
– used Planck’s theory and explained the photoelectric effect.
– light travels as tiny particles, photons.
Planck
Einstein
Compton

22. Another Look at the Atom
Chapter 4-3
Another Look at the Atom

23. Bohr’s 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, … ∞
Only electrons dropping from a Higher Level to a Lower one emit EMR
A Number of Possibilities for electron drops

24. Hydrogen’s 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. The Lines become more closely spaced as the levels increase
The Bohr model explained spectral lines but not how atoms bonded.
Ultimately Displaced

27. 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 – Plank’s Con.
m – mass
v - velocity

28. Small Objects – Large Wavelengths
Large Objects –
Small Wavelengths
200 g 30 m/s -  = cm
Undetectable
Small Objects – Large Wavelengths
9.11 x 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

29. STOP!

30. Heisenberg’s 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.

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
s orbital
p 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
N=1 1s
N=2
2s, 2p
2s(one) + 2p(three)
N=3
3s, 3p, 3d
3s(one) + 3p(three) + 3d (five)
N=4
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 2 4 8 3 9 18 16 32

45. Chapter 4.4 Quantum Numbers
Orbital
Max # electrons s 2 p 6 d 10 f 14

46. 4.5 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.
Hund’s Rule – orbitals of equal energy each receive one electron (equal spin) before any receive two.
Pauli’s 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 ____
Electron is represented by an ½ Arrow 
+ ½ ()
- ½ () 

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 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

50. Short-Hand Notation Eliminates the lines & arrows
Superscripts are used to illustrate the number of electrons in the sublevel
Same order of sublevels

51. Electron-Configuration Notation
Write the electron-configuration for the following: Cs Kr Br Po 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. Identifying Electrons
Paired electrons – when 2 electrons are within the same orbital.
Unpaired electrons – when a single electron is within an orbital.

54. Identifying Electrons
How many unpaired electrons does the following elements have? Na O B