1. The Bohr and Quantum Models

2. Starts with the Study of…..Light
Light is….
Made up of electromagnetic radiation.
Waves of electric and magnetic fields at right angles to each other.

3. Parts of a wave l Wavelength l lambda
Wavelength = the distance between any point on a wave and the corresponding point on the next wave (crest to crest, trough to trough); m, nm, pm

4. Parts of a wave
Frequency (n) = number of cycles in one second; hertz, hz
1 hertz = 1 cycle/second
n - nu

5. Frequency = n

6. Kinds of EM waves There are many different l and n (p. 92)
Radio waves, microwaves, x rays and gamma rays are all examples.
Visible light is all our eyes can detect. G a m m a R a d i o R a y s w a v e s

7. Electromagnetic Spectrum

8. Energy is Quantized - Planck
Quantum – a packet of energy; smallest quantity of energy that can be emitted or absorbed

9. Einstein is next
Light is particulate
photon – a quantum of electromagnetic radiation

10. Which is it? Is energy a wave like light, or a particle?
Yes !
Concept is called the Wave -Particle duality (Dual Nature of Light)
What about the other way, is matter a wave?

11. Spectrum The range of frequencies present in light.
White light has a continuous spectrum.
All the colors are possible.
A rainbow.

12. Electromagnetic Spectrum

13. Hydrogen spectrum
Line spectrum – series of separated lines of different colors representing photons whose wavelengths are characteristic of one element
656 nm
434 nm
410 nm
486 nm

14. What this means
Only certain energies are allowed for the hydrogen atom.(Recall relationship b/w energy, wavelength and frequency)
Energy in the atom is quantized.
Can calculate Energy based on line spectra
Only the energy we can “see”

15. Quantum Theory:
Colors emitted can be used to identify elements (absorption of energy and color emitted is a fingerprint of an element)
Kind of like wearing your team colors.
            
Team Oxygen
Team Carbon

16. Let’s Play!!
White light - dispersion
Flame Tests

17. Niels Bohr

18. The Bohr Ring Atom Quantized energy – energy levels
Ground state – ion or atom at its lowest energy state
Excited state – atom or molecule at any other state besides ground state
Putting energy (photon) into the atom moves the electron away from the nucleus.

19. Bohr From ground state (lowest energy level) to excited state.
When it returns to ground state (or lower energy level) it gives off light of a certain energy (photon.)
Energy, frequency, wavelength – color of light

20. The Bohr Ring Atom
n = 4
n = 3
n = 2
n = 1

21. Bohr Planetary Model quantum model of the hydrogen atom
atom was like a solar system; protons in dense nucleus
electrons orbit (specific path electron travels)
ground and excited state

22. The Bohr Model Only works for hydrogen atoms. Electrons don’t orbit.
Energy quantization correct, but not circling like planets
Introduced concept of energy levels

23. The Quantum Model

24. Atomic Models: Old version = Bohr’s New version = Quantum Theory
Most accepted
Diagrams electrons of a atom based on probability of location at any one time
Old version = Bohr’s
Also known as the planetary atomic model
Describes electron paths as perfect orbits with definite diameters
Good for a visual

25. Basic Concepts of the Quantum Model
Atoms and molecules can exist only in certain energy states
Atoms absorb or emit radiation as they change energies
Orbitals NOT orbits
Energy states are described by sets of numbers called quantum numbers
MATHEMATICAL model

26. Heisenberg’s Uncertainty Principle
We can’t know how the electron is moving or how it gets from one energy level to another.
Heisenberg Uncertainty Principle:
It is impossible to know the exact position AND momentum of a particle simultaneously. (Marco Polo)

27. Electron Configuration
The distribution of electrons within their atoms orbitals
Distributed amongst energy levels, sublevels, and orbitals
Based on a set of stated principles

28. AUFBAU PRINCIPLE
Electrons are added one at a time to the lowest energy orbitals
“building up” principle
Reference sheet

29. Electron Configurations
# of electrons in sublevel
H 1s1
Energy Level
sublevel

30. How to Determine Electron Configurations and Orbital Diagrams
Determine total # of electrons
Use Aufbau Principle to “fill”
Make sure all electrons of the atom are accounted for!

31. 18 electrons
Argon
(18 electrons)
Energy

32. 18 electrons
Argon
(18 electrons)
Energy
1s22s22p63s23p6

33. Orbital Diagram – box, circle or line for each orbital in a given energy level, grouped by sublevel, with an arrow indicating electron AND its spin

34. HUND’S RULE
Hund’s Rule (BUS RULE)- The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in the orbital. C
1s s p2

35. Details
Valence electrons- the electrons in the outermost energy levels (not d).
Core electrons- the inner electrons.

36. TRY!
Calcium and Manganese

37. Calcium
(20 electrons)
Energy
1s22s22p63s23p64s2

38. Try Oxygen, Iron, Strontium, Molybdenum, Tungsten!
Orbital Diagrams AND Electronic Configuratiions

39. Using the Periodic Table!
Noble Gas Configuration!

40. Blocks and Sublevels P E R
I  O D S
1-7 1 2 3 4 5 6 7
d (n-1) 4 5

41. Label your blank periodic table. Read it “like a book”

42. WRITE the Electron Configuration using the periodic table:Kr Ca Fe Hg

43. Electron Configuration – Noble Gas Configuration
Electron Configuration demonstrates a periodic trend, so you can write shorthand electron configuration using the electron configuration of the noble gases in Group 18 of the periodic table.
Noble gases have stable configurations.

44. Noble Gas Configuration
When writing shorthand e- config for an element, refer to the noble gas in the energy level (period) just above the element.
Write the symbol of the noble gas in brackets.
Write out the remaining e-config based on the energy filling diagram.
Electron Configuration
Na = 1s22s22p63s1
Al = 1s22s22p63s23p1
Ne = 1s22s22p6
Shorthand Electron Configuration
Na = [Ne] 3s1
Al = [Ne] 3s23p1

45. Noble Gas Configuration
EX: Na
Step 1: Na is in period 3 so refer to the noble gas in period 2 which is Neon.
Step 2: Write Ne in brackets. [Ne]
Step 3: Now write remaining electrons in standard form. 3s1.
Step 4: Combine. [Ne]3s1

46. Nobel Gas Configuration
Now try: I Kr Na Cu

47. Electron Configuration with Ions
When we write the electron configuration of a positive ion, we remove one electron for each positive charge:
Na → Na+
1s2 2s2 2p6 3s1 → 1s2 2s2 2p6
When we write the electron configuration of a negative ion, we add one electron for each negative charge:
O → O2-
1s2 2s2 2p4 → 1s2 2s2 2p6

48. Electron Configuration with Ions
Now try:
Ca+2
Fe-3

49. Details
Elements in the same column have the same electron configuration.
Put in columns because of similar properties.
Similar properties because of electron configuration.
Noble gases have filled energy levels.

50. Exceptions Ti = [Ar] 4s2 3d2 V = [Ar] 4s2 3d3 Cr = [Ar] 4s1 3d5
Mn = [Ar] 4s2 3d5
Half filled orbitals.
Scientists aren’t sure of why it happens
same for Cu [Ar] 4s1 3d10

51. More exceptions Lanthanum La: [Xe] 6s2 5d1 Cerium Ce: [Xe] 6s2 4f1 5d1
Promethium Pr: [Xe] 6s2 4f3 5d0
Gadolinium Gd: [Xe] 6s2 4f7 5d1
Lutetium Pr: [Xe] 6s2 4f14 5d1
We’ll just pretend that all except Cu and Cr follow the rules.

52. Quantum Numbers
Quantum Numbers represent 4 solutions to Schroedinger’s Equation
Describe the distribution of electrons in space
BASED ON PROBABILITY!!!

53. 1. Principal Quantum Number (n)
Describes main energy level an electron occupies
n=1,2,3,4,5,6,7

54. 2. Angular Momentum Quantum (l)
sublevel
designates the shape of the region in space an electron occupies (orbital).
integer values from 0 to n-1
l = 0 is called s
Sublevels include s (l=0, p=1, d=2, f=3)
Not all energy levels have all sublevels

55. 3. Magnetic quantum number (Orbitals ml)
designates the spatial orientation of orbital
integer values between - l and + l
s: ml = 0
p: ml = -1, 0, +1 (i.e. px, py, pz)
s-1, p-3, d-5, f-7

56. S orbitals

57. P orbitals

58. P Orbitals

59. D orbitals

60. F orbitals

61. F orbitals

62. 4. Spin Quantum Number m s Electrons are negatively charged
Behave as if spinning on axis – repel so…
either +1/2 or -1/2

63. Pauli Exclusion Principle
NO TWO ELECTRONS MAY HAVE THE SAME FOUR QUANTUM NUMBERS
(Spin must be different!)