1. Slide #3: Use Spectrascopes Slide #21: Breaking Bad: colored flames
Teacher Demos:
Slide #3: Use Spectrascopes
Slide #21: Breaking Bad: colored flames

2. Chapter 4: Electrons

3. Electromagnetic Radiation
Electromagentic Radiation(EM)- energy, in the form of photons, that moves in waves as it travels through space.
Examples of EM Radiation are:
1 gamma rays 6 microwaves
2 x-rays 7 radar waves
3 ultraviolet radiation 8 radio waves
4 visible light
5 infra-red radiation

4. Electromagnetic Spectrum

5. Electromagnetic Spectrum

6. Electromagnetic Radiation
Figure 7.1

7. Wave measurements

8. Wave measurements Crest- highest point of a wave cycle.
Trough – lowest point of a wave cycle.
Wavelength (l)– distance between two consecutive points on a wave.
Amplitude – height of the wave, from the axis.
Frequency (n) – the number of wave cycles per second.

9. Photons All forms of Electromagnetic Radiation are made up of photons.
Photon (quantum)- a tiny packet of light energy that travels through space in electric and magnetic waves.
Photons behave as both waves and particles.(Duality of Light Theory)

10. Waves III(SKIP)

11. Visible Light All photons with a wavelength between 360 nm -740 nm.
Ultra violet(UV) visible light Infra-red(IR) v i b g y o r nm nm

12. Electromagnetic Radiation
Waves have a frequency
Use the Greek letter “nu”, , for frequency, and units are “cycles per sec”
All radiation: c =  •  where c = speed of light = 3.00 x 108 m/sec
Long wavelength --> small frequency
Short wavelength --> high frequency

13. Important Conversion 1 meter = 1 x 109 nanometers (memorize!)
Make each conversion:
364 nm = ______ meter
9.88 x meter = ______nm
2000 nm = ______meter
m = ________nm

14. Practice Problems c =  • 
Find the wavelength of a photon(light) if it has a frequency of 8.1 x 1012 Hz.
Find the frequency of a photon of light with a wavelength of 350 nm.
Find the wavelength of a 6 x 1014 Hz photon and indicate its color.
What is the frequency of a photon with
a wavelength of 895 nm?

15. Bohr Atomic Model

16. How are photons produced?
(n= infinity)
(n=3) (excited states)
(n=2)
Energy
(n=1) (ground state)
Electron

17. How are photons produced?
Bohr Model of an atom a hydrogen electron moves around the nucleus only in certain allowed orbits (energy levels).
If an electron in the ground state (lowest energy level), it absorbs energy it goes to an excited state.
All excited electrons always return to the ground state.
As electrons return to the ground state, they release any extra energy in the form of photons.

18. How are photons produced?

19. How are photons produced?

20. Line Spectra(pre-AP)
Each element has atoms with different numbers of electrons.
As these electrons drop to lower energy levels, they emit photons of unique wavelengths.
Each element emits its own set of colored lines, called an emission line spectrum.
These spectra are used to identify elements in unknown samples or composition of stars.

21. Hydrogen atom spectra series

22. Line spectra series
Lyman Series- e- drops from n = X to n=1 lights are in Ultraviolet Region Balmer Series- e- drops from n = X to n =2 lights are in Visible Region Paschen Series- e- drops from n = X to n = 3 lights are in Infrared Region

23. Calculating the energy in a photon
E = hn
Energy = Planck’s Constant x frequency
Find the energy in a 1.8 x 1014 Hz photon.
Find the frequency of a 5.05 x J photon.
Find the energy of a 485 nm photon.

24. Photon calculations 1) Find the energy of a 3.50 x 1014Hz photon.
2) What is the frequency of a 6.13 x J photon?
A photon has a wavelength of 525 nm. What is the frequency?
4) What is the energy of this photon?
A photon has an energy of 1.05 x J. Is this photon visible?
Which line series does this light belong to?

25. Independent Practice Work out section review 4.1 of resource binder.
Quiz Tomorrow! Hurrrrrrrray!!!!!!!

26. Electron orbitals
Orbital – a 3-dimensional space around the nucleus which can hold up to 2 electrons, with opposite spin.
*electrons are found in their orbitals 99.9% of the time.
Orbitals have different shapes: s, p, d, f

27. Orbitals have different shapes:
Orbital shape
s spherical
p dumbbell
d clover or
dumbbell/donut
f too complex

28. s-orbitals: spheres

29. p-orbitals: dumbbell

30. d-orbitals: clover (double dumbbell) or dumbbell/donut

31. f-orbitals: complex

32. Quantum Numbers
The location of each electron in an atom can be determined by assigning each electron a four-number code called quantum numbers.

33. Important electron laws:
Heisenberg Uncertainty Principle : the more accurately one knows the position of an electron the less accurately one can predict its momentum(speed) AND vice versa.

34. Important electron laws:
Hund’s Rule: electrons fill empty orbitals first.
Aufbau Principle: electrons occupy orbitals in a certain order: nearest to the nucleus first.
Pauli Exclusion Principle: No two electrons in the same atom can have the same set of four quantum numbers.

35. 4 Quantum Numbers:
1 Principal Quantum Number (n), gives the energy level (aka shell) of the electron.
2 Orbital Quantum Number (l) = gives the shape of orbital of the electron.
3 Magnetic Quantum Number(ml) = gives the orientation(direction) of the orbital.
4 Spin Quantum Number(ms) = gives the direction of spin of the electron.

36. Allowable ranges for quantum numbers
n (energy level) = 1  infinity
l (type of orbital = 0  n-1
ml (direction) = -l  +l
ms (spin) = -1/2 or +1/2

37. Determine whether each set of quantum numbers is valid or invalid:
1) 3, 2, -1, +1/2
2) 1, 2, 0, -1/2
3) 4, 3, -3, +1/4
4) 88, 67, -55, -1/2
5) -2, 1, -1, +1/2
6) 1, 0, 1, -1/2
7) 2, 0, 0, +1/2
8) 4, 2, -3, +1/2
9) 3, 1, 2, -1/2
10) 4, 2, -2, +1/2
Valid
Invalid

38. Energy sublevels around an atom
energy electron
sublevels # of orbitals capacity s p d f g*
*orbitals exist, but not used most of the time.

39. Orbital Notation Aufbau Order:
(Nucleus)1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f
**This is the order in which electrons fill!!!!!
You must learn the order!! Don’t worry there is always an easier way to memorize these things. For the Aufbau Order there are 2 ways: With arrows or with the periodic table! Smartboard activate!

40. Electron Configurations and the Periodic Table
Figure 8.7

41. Orbital Notation Write the orbital notation for each atom:
Nitrogen, N(7 electrons)
Sodium, Na(11 e-)
Iron, Fe(__ e-)
Antimony, Sb(__ e-)
Gold, Au(__ e-)

42. Electron Configurations
Write the electron configurations for each atom:
Nitrogen, N(7 electrons)
Sodium, Na(11 e-)
Iron, Fe(__ e-)
Antimony, Sb(__ e-)
Gold, Au(__ e-)

43. Quantum Numbers
Write the four quantum numbers for the last electron to fill each atom:
Nitrogen, N(7 electrons)
Sodium, Na(11 e-)
Iron, Fe(__ e-)
Antimony, Sb(__ e-)
Gold, Au(__ e-)

44. Valence electrons

45. Valence electrons Give the number of valence electrons for each atom:
Nitrogen, N(7 electrons)
Sodium, Na(11 e-)
Iron, Fe(__ e-)
Antimony, Sb(__ e-)
Gold, Au(__ e-)

46. Lewis Dot Structures Give the Lewis Dot Structure for each atom:
Nitrogen, N(7 electrons)
Sodium, Na(11 e-)
Iron, Fe(__ e-)
Antimony, Sb(__ e-)
Gold, Au(__ e-)

47. Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic field all electrons are spin paired.
Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.