1. 3-Atomic Structure Overview Characteristics of Atoms
Interaction b/tw matter and light
Photoelectric Effect
Absorption and Emission Spectra
Electron behavior
Quantum numbers

2. Atomic Structure Atomic orbitals Periodic table Orbital energies
Electron configuration and the periodic table
Periodic table
Periodic properties
Energy

3. Characteristics of Atoms
Atoms possess mass
Atoms contain positive nuclei
Atoms contain electrons
Atoms occupy volume
Atoms have various properties
Atoms attract one another
Atoms can combine with one another to form molecules

4. Atomic Structure
Atomic structure studied through atomic interaction with light
Light: electromagnetic radiation
carries energy through space
moves at 3.00 x 108 m/s in vacuum
wavelike characteristics

5. Electromagnetic Spectrum

6. Visible Spectrum

7. Wavelength () & Frequency ()
amplitude
 = number of complete cycles to pass given point in 1 second

8. Energy c =  x  = 3.00 x 108 m/s long wavelength  low frequency
Low Energy
High Energy
short wavelength  high frequency

9. Energy Mathematical relationship: E = h E = energy
h = Planck’s constant: 6.63 x 10–34 J s
 = frequency in s–1

10. Energy Mathematical relationship: E = h c =  x  E =
Energy: directly proportional to frequency
inversely proportional to wavelength

11. Problems 3-1, 2, & 3
a) Calculate the wavelength of light with a frequency  = 5.77 x 1014 s–1
b) What is the energy of this light?
2. Which is higher in energy, light of wave-length of 250 nm or light of 5.4 x 10–7 m?
3. a) What is the frequency of light with an energy of 3.4 x 10–19 J?
b) What is the wavelength of light with an energy of 1.4 x 10–20 J?

12. Photoelectric Effect Light on metal surface Electrons emitted
Threshold frequency, o
If  < o, no photoelectric effect
If  > o, photoelectric effect
As  , kinetic energy of electrons 

13. Photoelectric Effect Einstein: energy  frequency
If  < o electron doesn’t have enough energy to leave the atom
If  > o electron does have enough energy to leave the atom
Energy is transferred from light to electron, extra is kinetic energy of electron
Ephoton = hphoton = ho + KEelectron
KEelectron = hphoton – ho
Animation

14. Problem 3-4
A given metal has a photoelectric threshold frequency of o = 1.3 x 1014 s1. If light of  = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.

15. Bohr Model Line spectra Light through a prism  continuous spectrum:
Ordinary
white
light

16. Bohr Model Line spectra Light from gas-discharge tube
through a prism  line spectrum:
H2 discharge tube

17. Line Spectra (emission)
White light H He Ne

18. Line Spectra (absorption)
Gas-filled tube
Light source

19. Bohr Model For hydrogen: C = 3.29 x 1015 s–1
Niels Bohr: Electron energy in the atom is quantized.
n = 1, 2, 3,….
RH = 2.18 x 10–18 J

20. Bohr Model E = Ef – Ei = h Eatom = Eelectron = h Line spectrum
Minus sign: free electron has zero energy
Line spectrum
Photoelectric effect:

21. Bohr Energy Levels

22. Electrons All electrons have same charge and mass
Electrons have properties of waves and particles (De Broglie)

23. Heisenberg Uncertainty Principle
Cannot simultaneously know the position and momentum of electron
x = h
Recognition that classical mechanics don’t work at atomic level.

24. Schrödinger Equation Erwin Schrödinger 1926
Wave functions with discrete energies
Less empirical, more theoretical
n En
n wave functions or orbitals
n2 probability density functions

25. Quantum Numbers Each orbital defined by 3 quantum numbers
Quantum number: number that labels state of electron and specifies the value of a property

26. Quantum Numbers Principal quantum number, n (shell)
Specifies energy of electron (analogous to Bohr’s n)
Average distance from nucleus
n = 1, 2, 3, 4…..

27. Quantum Numbers Azimuthal quantum number,  (subshell) = 0, 1, 2… n–1
n = 2,  = 0 or 1
n = 3,  = 0, 1, or 2
Etc. 1 2 3 4 s p d f g

28. Quantum Numbers Magnetic quantum number, m
Describes the orientation of orbital in space
m = –….+ 
If  = 2, m = –2, –1, 0, +1, +2

29. Problem 3-5 Fill in the quantum numbers in the table below. n  m 33s 2
–2, –1, 0, 1, 2 2p

30. Schrödinger Equation Wave equations: 
Each electron has  & E associated w/ it
Probability Density Functions: 2
-graphical depiction of high probability of finding electron

31. Probability Density Functions
Link to Ron Rinehart’s page
 energy
2 probability density function
s, p, d, f, g 1s 3s 2s
Node: area of 0 electron density

32. Probability Density Functions
Node: area of 0 electron density
nodes
Link to Ron Rinehart’s page

33. Electrons and Orbitals
Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers
Electron spin quantum number ms = ½
Electrons are spin paired within a given orbital

34. Electrons and Orbitals
 = 0, m = 0, ms = ½
2 electrons possible:
1,0,0,+½ and 1,0,0,–½
2 electrons per orbital
1s1 H
1s2 He

35. Electrons and Orbitals
 = 0, m = 0, ms = ½
2,0,0, ½
2 electrons possible
 = 1, m = –1,0,+1, ms = ½
2,1,–1, ½ ,1,0, ½ 2,1,+1, ½
6 electrons possible

36. Electron Configurations
1s 2 electrons possible
H 1e– 1s1 
He 2e– 1s2 

37. Electron Configurations
2s 2 electrons possible
Li 3e– 1s2 2s1 2s  1s 
Be 4e– 1s2 2s2 2s  1s 

38. Electron Configurations
2p  = 1, m = –1, 0, +1
3 x 2p orbitals (px, py, pz): 6 electrons possible 2p
B 5e– 1s2 2s2 2p1 2s 1s

39. Electron Configurations
2p  = 1, m = –1, 0, +1
3 x 2p orbitals (px, py, pz): 6 electrons possible 2p 
B 5e– 1s2 2s2 2p1 2s  1s 

40. Electron Configurations
2p  = 1, m = –1, 0, +1
C 6e– 1s2 2s2 2p2  1s  2s 2p
Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized

41. Problem 3-6
Write electron configurations and depict the electrons for N, O, F, and Ne.

42. Electron Configurations
3s, 3p, 3d  1s 2s 2p  3s 3p
Na 11e– 1s2 2s2 2p63s1

43. Electron Configurations
3s, 3p, 3d  1s 2s 2p 3s 3p
Mg 12e– 1s2 2s2 2p63s2

44. Electron Configurations
3s, 3p, 3d  1s 2s 2p  3s 3p
Al 13e– 1s2 2s2 2p63s23p1

45. Electron Configurations
3s, 3p, 3d  1s 2s 2p  3s 3p
Si 14e– 1s2 2s2 2p63s23p2

46. Electron Configurations
3s, 3p, 3d  1s 2s 2p  3s 3p
P 15e– 1s2 2s2 2p63s23p3

47. Electron Configurations
3s, 3p, 3d  1s 2s 2p  3s 3p
S 16e– 1s2 2s2 2p63s23p4

48. Electron Configurations
3s, 3p, 3d  1s 2s 2p 3s  3p
Cl 17e– 1s2 2s2 2p63s23p5

49. Electron Configurations
3s, 3p, 3d  1s 2s 2p 3s 3p
Ar 18e– 1s2 2s2 2p63s23p6

50. Electron Configurations
3d vs. 4s
Filling order 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
6s 6p 6d
7s 7p

51. Electron Configurations4p K 3d 4s  3p    3s  2p    2s  1s 

52. Electron Configurations4p Ca 3d 4s  3p    3s  2p    2s  1s 

53. Electron Configurations4p 3d  Sc 4s  3p    3s  2p    2s  1s 

54. Electron Configurations4p Ti 3d   4s  3p    3s  2p    2s 
Link to OSU site 1s 

55. Problem 3-7
Write the electron configurations for the transition metals V – Zn. Fill in the corresponding boxes to denote the electronic spin.