1. Arrangement of the Electrons Chapter 4
(Electron Configurations)

2. Electron Behavior
Scientists began to understand how electrons acted by observing the way that light interacts with matter.

3. In your notebook, answer the question:
How would you describe what light is to someone else?

4. Spectrum of Light!
Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space.
Electromagnetic Spectrum-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation.

5. Parts of the wave Wavelength-lambda (λ)
The distance between corresponding points on adjacent waves. Units: m, nm, cm, or Å
Frequency-nu (ν)
The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/s = 1/sec = s-1

6. RESULT-An electromagnetic wave.
When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other.
RESULT-An electromagnetic wave.
EM waves do not require a medium to propagate through
When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other RESULT-An electromagnetic wave. Click

7. Thursday 10/29 Today you will need: Agenda: Notes Calculator
Periodic Table
Agenda:
Review Properties of Waves and EM Spectrum
Flame Test Demo
Frequency and Wavelength Calculations
Goal: Know the relationship between λ and v and how to calculate them!

8. Parts of the wave Wavelength-lambda (λ)
The distance between corresponding points on adjacent waves. Units: m, nm, cm, or Å
Frequency-nu (ν)
The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/s = 1/s = s-1

9. Spectrum of Light!
Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space.
Electromagnetic Spectrum-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation.

10. Relationship between λ and ν
c = λ∙ν
λ = wavelength (m)
ν = frequency (Hz)
c = speed of light= 3.0 x 108 m/s (constant)
λ and ν are _______________ related.
Calculate the frequency for violet light.

12. Practice Problem
Truck-mounted helium-neon laser produces red light whose wavelength (λ ) is 633 nanometers. Determine the frequency (v).
*Remember that c=3.0x108m/s.
*Use the formula c= λ . v

13. Frequency = 4.7x1014 Hz (cycles per second)
c= λ . v
c =3.0x108 m/s
c= λ . v
v=c / λ
λ = 633nm= 6.33x10-7m
v = 3.0x108 m/s = 0.47x 1015s-1 = 4.7x1014 s-1
6.33x10-7m
Frequency = 4.7x1014 Hz (cycles per second)

14. B. EM Spectrum  = ?  = c   = 434 nm = 4.34  10-7 m
EX: Find the frequency of a photon with a wavelength of 434 nm.
GIVEN:
 = ?
 = 434 nm = 4.34  10-7 m
c = 3.00  108 m/s
WORK:
 = c 
 = 3.00  108 m/s
4.34  10-7 m
 = 6.91  1014 Hz

15. Calculate the frequency for the yellow- orange light of sodium.
Calculate the frequency for violet light.

16. Relationship between Energy and ν
E = h∙ν
E = energy (joule)
h = Planck’s constant = 6.63 x J∙s
ν = frequency (Hz)
E and ν are ______________ related.
Calculate the energy for the yellow-orange light for sodium.
Calculate the energy for the violet light.
A plot of light frequencies with KE of escaping electrons of different metals revealed a slope of 6.63x10-34Js, or said another way, the kinetic energy increases linearly with frequency proportional to planck’s constant.

17. MoNday 11/2 Welcome! Please turn into the box: -Chp. 4 Outline
-EM Spectrum Practice (if you didn’t Fri)
-Energy Calculations (if you didn’t Fri)

18. Agenda Chp. 4 Reading Quiz Pass Back Chp. 3 Quiz
Wavelength, Frequency, Energy Problem
Evidence that light is a wave!
If Time: Evidence that light is a particle!

19. Light Calculations
Pick your favorite (or any, really) radio station. Knowing that radio station frequencies are measured in Mega Herz and that 1 MHz = 106Hz, calculate:
A)The wavelength (in m) of your radio wave
B) The Energy associated with your station
Discuss: What do you know/notice about the relationship between wavelength and energy?

20. Tuesday 11/3
Today!
Understand that there is experimental evidence that supports the idea that light is both a wave AND a particle.

21. Evidence That Light is a Wave!

22. Young’s Double Slit Experiment
Constructive Interference:
Before
During
Destructive Interference:

23. Young’s Double Slit Experiment

24. Problems with Light as a wave

25. Light as waves and particles (the Particle Theory of light)
2 problems that could not be explained if light only acted as a wave.
1.) Emission of Light by Hot bodies:
Characteristic color given off as bodies are heated: red  yellow  white
If light were a wave, energy would be given off continually in the infrared (IR) region of the spectrum.

26. The second problem………
2.) Absorption of Light by Matter = Photoelectric Effect
Light can only cause electrons to be ejected from a metallic surface if that light is at least a minimum threshold frequency . The intensity is not important.
If light were only a wave intensity would be the determining factor, not the frequency!

27. Wed 11/4
Please turn your photoelectric effect assignment into the box.
Tomorrow is the Flame Test Lab.
We will meet in Rm. 126
To participate, you MUST be wearing proper lab attire.

29. Einstein

30. Max Planck (1900’s) Particle Theory of Light
When an object loses energy, it doesn’t happen continuously but in small packages called “quanta”.
“Quantum”-a definite amount of energy either lost or gained by an atom.
“Photon”-a quantum of light or a particle of radiation.

31. Line SpectrA
Excited State: Higher energy state than the atom normally exists in.
Ground State: Lowest energy state “happy state”
Line Spectrum: Discrete wavelengths of light emitted.
2 Types:
1.) Emission Spectrum: All wavelengths of light emitted by an atom.
2.) Absorption Spectrum: All wavelengths of light that are not absorbed by an atom. This is a continuous spectrum with wavelengths removed that are absorbed by the atom. These are shown as black lines for absorbed light.
Continuous Spectrum: All wavelengths of a region of the spectrum are represented (i.e. visible light)

32. Hydrogen line Spectrum & niel’s Bohr
Hydrogen’s spectrum can be explained with the wave-particle theory of light.
Niel’s Bohr (1913)
1.) The electron travels in orbits (energy levels) around the nucleus.
2.) The orbits closest to the nucleus are lowest in energy, those further out are higher in energy.
3.) When energy is absorbed by the atom, the electron moves into a higher energy orbit. This energy is released when the electron falls back to a lower energy orbit. A photon of light is emitted.

33. Drawing Bohr Models

34. Bohr Models Bohr models are used to predict reactivity in elements.
Reactivity refers to how likely an element is to form a compound with another element.
When looking at Bohr models, we look at its valence electrons (the electrons on the last energy level) to determine reactivity.

35. Drawing Bohr Models Draw the nucleus.
Write the number of neutrons and the number of protons in the nucleus.
Draw the first energy level.
Draw the electrons in the energy levels according to the rules (on the next slide). Make sure you draw the electrons in pairs.
Keep track of how many electrons are put in each level and the number of electrons left to use.

36. Valence Electrons Electron Shell Number of Electrons 1 2 8 3 4 18 5 6
Each electron shell can hold a certain number of electrons
Electron shells are filled from the inside out
Noble Gases have full outer electron shells
All other elements have partially filled outer electron shells
Electron Shell
Number of Electrons 1 2 8 3 4 18 5 6 32 7

37. Guided Practice
In order to draw Bohr models of these elements, you must first determine the number of protons, neutrons, and electrons. Once you have found this information, follow the directions to draw your model. 6 6 6 6 C
Carbon
12.011
Protons: _____ Neutrons: _____ Electrons: ______
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Bohr Model: 2 4

38. Guided Practice 16 S
Sulfur
32.066 16 16 16
Protons: _____ Neutrons: _____ Electrons: ______
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Bohr Model: 3 6

39. Guided Practice Protons: _____ Neutrons: _____ Electrons: ______3 Li
Lithium
6.941
Protons: _____ Neutrons: _____ Electrons: ______
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Bohr Model: 3 4 3 2 1

40. Guided Practice 10 10 10 10 Ne
Neon
20.180
Protons: _____ Neutrons: _____ Electrons: ______
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Bohr Model: 2 8

41. Warm-Up Protons: _____ Neutrons: _____ Electrons: ______ 15 16 15
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Draw the Bohr Model: 15 16 15 15 P
Phosphorus
30.974 3 5

42. Warm UP 11 12 11 Protons: _____ Neutrons: _____ Electrons: ______Na
Sodium
22.990
Protons: _____ Neutrons: _____ Electrons: ______
How many energy shells will this have? ____
How many valence (outer) electrons does this element have? ____
Draw the Bohr Model: 3 1

43. BOHR’S Model Only worked for Hydrogen!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!

44. Hydrogen Spectrum
Lyman Series-electrons falling to the 1st orbit, these are highest energy, _____ region.
Balmer Series- electrons falling to the 2nd orbit, intermediate energy, _______ region.
Paschen Series-electrons falling to the 3rd orbit, smallest energy, ______ region.

45. Bohr’s equation for Hydrogen
En = (-RH) 1/n2
En = energy of an electron in an allowed orbit (n=1, n=2, n=3, etc.)
n = principal quantum number (1-7)
RH = Rydberg constant (2.18 x J)
When an electron jumps between energy levels: ΔE =Ef – Ei
By substitution: ΔE = hν = RH(1/ni2 - 1/nf2)
When nf > ni then ΔE = (+)
When nf < ni then ΔE = (-)

46. New Theory Needed to explain more complex atoms!
DeBroglie (1924)-Wave properties of the electron was observed from the diffraction pattern created by a stream of electrons.
Schrodinger (1926)-Developed an equation that correctly accounts for the wave property of the electron and all spectra of atoms. (very complex)

47. Schrodinger Equation for Hydrogen
DO NOT WRITE THIS IN YOUR NOTES!

49. Quantum Theory (current theory of the atom)
Rather than orbits  we refer to orbitals. These are 3-dimensional regions of space where there is a high probability of locating the electron.
Heisenberg Uncertainty Principle-it is not possible to know the exact location and momentum (speed) of an electron at the same time.
Quantum Numbers-4 numbers that are used to identify the highest probability location for the electron.

50. Quantum numbers 1.) Principal Quantum Number (n)
States the main energy level of the electron and also identifies the number of sublevels that are possible.
n=1, n=2, n=3, etc. to n=7
2.) Azimuthal Quantum Number (l)
Values from 0 to n-1
Identifies the shape of the orbital
l = 0 s sphere 1 orbital
l = 1 p dumbbell 3 orbitals
l = 2 d 4-4 leaf clovers & 1-dumbbell w/doughnut5 orbitals
l = 3 f very complex 7 orbitals

51. Quantum numbers (cont.)
3.) Magnetic Quantum Number (ml)
Values from –l  l
States the orientation in space (x, y, z)
ml = 0 s only 1 orientation
ml = -1, 0, +1 p 3 orientations
ml = -2,-1,0,+1,+2 d 5 orientations
ml = -3,-2,-1,0,+1+2,+3 f 7 orientations
4.) Spin Quantum Number (ms)
Values of +1/2 to -1/2
States the spin of the electron.
Each orbital can hold at most 2 electrons with opposite spin.