1. Thompson’s Experiment

2. Rutherford’s Experiment:

3. Explanation

5. Diffraction:- the change in direction of a wave as
it passes the edge of an object

6. The Wave Nature of Light
Light as electromagnetic waves: polarization, interference, diffraction, reflection, and refraction

8. Electromagnetic Spectrum

11. Atomic Line Spectra (Line Emission Spectra)
Hydrogen has the simplest atomic emission spectrum ( 1880s).

12. The Spectrum of Atomic Hydrogen
Compare the absorption spectrum to the emission spectrum of H
If we pass light from a continuous source (eg from a hot object) through a substance, then the continuous spectrum has some of the wavelengths removed! -- absorption spectrum

14. Light as Particles

15. Bohr Atom

16. Hidrojen İçin Bohr Atom Modeli (1913)
Elektronlar belirli yörüngelerde bulunabilirler

17. Bohr Model of Hydrogen Atoms
Assumptions
Quantized energy levels – total energy for each level is the sum of the kinetic energy of the electron plus its potential energy.
Electrons do not radiate while in their orbits, but do when they move from one orbit to another.

18. Energy E = 0 eV Paschen Series (IR) n = 3 Balmer Series (visible)
Lyman Series
(ultraviolet)
E1 = eV
n = 1
Lyman
Balmer
Paschen
Example Data

19. Line Spectra
The Lyman and Balmer series of lines in the hydrogen spectrum correspond to transitions that the electrons make between higher and lower energy levels.
The Bohr model only has one quantum number, n, which represents the energy level.

20. Electron in the Hydrogen Atom

21. Elektronlar gerçekten de çekirdek etrafında belirli yörüngelerde mi dönerler?
Bunu tespit etmek mümkün mü?

22. Elektronun konumu ve momentumunun ölçülmesi
Işının her çarpışında elektronun da konumu değişir. Bu sebeple ışıkla bir elektronunun konumu ve hızı hassas bir şekilde belirlenemez.

23. Heisenberg, Werner 1901–76, Alman Fizikçi 1932 Nobel Fizik Ödülü
Eğer bir taneciğin nerede olduğunu kesin olarak biliyorsak, aynı anda taneciğin nereden geldiğini ve nereye gittiğini kesin şekilde bilemeyiz.

24. Peki elektronlar nerede ve nasıl hareket ediyorlar?

25. Erwin Schrödinger 1927 yılında
Elektronlar, zamanlarının büyük bir çoğunluğunu orbital denen bölgelerde geçirirler.
Değişik şekillerde orbitaller mevcuttur.

26. Atomik Orbitaller Şekli Orbital sayısı e sayısı s küresel 1 2
p halter
d karışık
f karışık
Herbir orbital 2 elektron içerir

27. Quantum Mechanics

28. The First Shell
The innermost shell (1) only contains an s orbital.   1s

29. Atomic Orbitals, s-type

30. S orbitalleri

31. Atomic Structure

32. Quantum Mechanics

33. Electron Configuration in p Orbital

34. Atomic Orbitals, p-type

35. p Orbitals
Rather than being a sphere, the "p" orbital has two lobes pointed in opposite direction away from the nucleus.  
One p orbital points along each the x, y, and z axis.
There are three p orbtitals in every shell except the first.

36. The Second Shell 2py 2px 2pz 2s
The second shell contains an s type orbital as well as a new kind of orbital called a "p" orbital.
2py
2px
2pz 2s

37. Atomic Orbitals, d-type

38. d Orbitals
There are 5 types of d orbitals. Four of the five have four lobes at 90o to one another. The fifth looks like a donut around a p-orbital
Image from: HMChem

39. The Third Shell The 3rd shell has: 1 s orbital (3s)
3 p orbitals (3px, 3py, 3pz)
5 d orbitals (3dxy, 3dyz, 3dxz, 3dx2-y2, 3dz2)

41. f orbitals
f-orbitals have 6 lobes and are very challenging to envision
Image from: HMChem

42. The First Two Shells
Picture from and more info on How Atoms Work

43. The Fouth Shell The 4th shell has: 1 s orbital (4s)
3 p orbitals (4px, 4py, 4pz)
5 d orbitals (4dxy, 4dyz, 4dxz, 4dx2-y2, 4dz2)
7 f orbitals
What pattern do you see in the relationship between shells and orbitals?
Electrons are rarely if ever observed in orbitals beyond f so, we most chemists are not too worried about them.

44. Principle quantum number
n = 1, 2, 3,… describes orbital size and energy
Angular momentum quantum number
l = 0 to n describes orbital shape
Magnetic quantum number
ml = l, l-1…-l describes orientation in space
of the orbital relative to the other
orbitals in the atom
Spin quantum number
ms = +1/2 or -1/ describes the direction of spin
of the e- on its axis
Pauli Exclusion Principle: "no two electrons in an
atom can have the same set of quantum numbers",
or, only two electrons (of opposite spin) per orbital.

45. Write a valid set of quantum numbers for each of the
following sub-shells:
(a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2
n = 2, l = 0, ml = 0, ms = ± 1/2
2 combinations

46. Write a valid set of quantum numbers for each of the
following sub-shells:
(a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2
n = 2, l = 0, ml = 0, ms = ± 1/2
2 combinations
(b) 2 p n = 2, l = 1, ml = -1, ms = - 1/2
n = 2, l = 1, ml = -1, 0 or 1, ms = ± 1/2
6 combinations

47. Write a valid set of quantum numbers for each of the
following sub-shells:
(a) 2 s n = 2, l = 0, ml = 0, ms = - 1/2
n = 2, l = 0, ml = 0, ms = ± 1/2
2 combinations
(b) 2 p n = 2, l = 1, ml = -1, ms = - 1/2
n = 2, l = 1, ml = -1, 0 or 1, ms = ± 1/2
6 combinations
(c) 3 d n = 3, l = 2, ml = -2, ms = - 1/2
n = 3, l = 2, ml = -2, -1, 0, 1, or 2, ms = ± 1/2
10 combinations

48. How many orbitals in a subshell?
l = 1, px, py, pz
l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5

49. How many orbitals in a subshell?
l = 1, px, py, pz
l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5
2 l + 1 orbitals per subshell

50. How many orbitals in a subshell?
l = 1, px, py, pz
l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5
2 l + 1 orbitals per subshell
How many orbitals in a shell?
n = 1, 1s
n = 2, 2s, 2px, 2py, 2pz
n = 3, 3s, 3px, 3py, 3pz, 3dxy,, 3dxz,, 3dyz ,, 3dx2-y2, 3dz2 9

51. How many orbitals in a subshell?
l = 1, px, py, pz
l = 2, dxy,, dxz,, dyz ,, dx2-y2, dz2 5
2 l + 1 orbitals per subshell
How many orbitals in a shell?
n = 1, 1s
n = 2, 2s, 2px, 2py, 2pz
n = 3, 3s, 3px, 3py, 3pz, 3dxy,, 3dxz,, 3dyz ,, 3dx2-y2, 3dz2 9
n2 orbitals per principal quantum level

52. Hydrogen atom-
all orbitals within a shell have the same energy
electrostatic interaction between e- and proton

53. Hydrogen atom-
all orbitals within a shell have the same energy
electrostatic interaction between e- and proton
Multi-electron atoms-
the energy level of an orbital depends not only on the
shell but also on the subshell
electrostatic interactions between e- and proton and other e-

54. Orbital Energies
3dxy
3dxz
3dyz
3dx2-y2
3dz2
3px
3py
3pz 3s
Energy
2px
2py
2pz 2s 1s

55. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle

56. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)

57. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)
3. Try to attain maximum number of unpaired
e- spins in a given sub-shell (Hund's Rule)

58. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)
3. Try to attain maximum number of unpaired
e- spins in a given sub-shell (Hund's Rule)
H (Z = 1) 1s1
2s 2p
Energy 1s

59. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)
3. Try to attain maximum number of unpaired
e- spins in a given sub-shell (Hund's Rule)
N (Z = 7) 1s2, 2s2, 2p3, 2p 2s
Energy 1s

60. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)
3. Try to attain maximum number of unpaired
e- spins in a given sub-shell (Hund's Rule)
B (Z = 5) 1s2, 2s2, 2p1 2p 2s
Energy 1s

61. Electronic Configuration: Filling-in of Atomic Orbitals
Rules: 1. Pauli Principle
2. Fill in e-'s from lowest energy orbital
upwards (Aufbau Principle)
3. Try to attain maximum number of unpaired
e- spins in a given sub-shell (Hund's Rule)
F (Z = 9) 1s2, 2s2, 2p5 2p 2s
Energy 1s

62. Hydrogen
2s 3s 4s
1s 2p 3p 4p
3d 4d 4f
Multi-electron atoms
1s s 3s 4s 5 s
2p 3p 4p

63. 1s s px 2py 2pz
H 1s1
He 1s2
Li 1s2, 2s1
Be 1s2, 2s2
B 1s2, 2s2, 2px1
C 1s2, 2s2, 2px1, 2py1
N 1s2, 2s2, 2px1, 2py1, 2pz1
O 1s2, 2s2, 2px2, 2py1, 2pz1
F 1s2, 2s2, 2px2, 2py2, 2pz1
Ne 1s2, 2s2, 2px2, 2py2, 2pz2

64. H 1s1 He 1s2
Li [He], 2s1 Be [He], 2s2

65. H 1s1 He 1s2
Li [He], 2s1 Be [He], 2s2
B [He], 2s2, 2p1 Ne [He], 2s2, 2p6
Na [He], 2s2, 2p6, 3s1  [Ne], 3s1

66. H 1s1 He 1s2
Li [He], 2s1 Be [He], 2s2
B [He], 2s2, 2p1 Ne [He], 2s2, 2p6
Na [He], 2s2, 2p6, 3s1  [Ne], 3s1
Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2
Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2

67. H 1s1 He 1s2
Li [He], 2s1 Be [He], 2s2
B [He], 2s2, 2p1 Ne [He], 2s2, 2p6
Na [He], 2s2, 2p6, 3s1  [Ne], 3s1
Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2
Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2
P [Ne], 3s2, 3p3 S [Ne], 3s2, 3p4
Cl [Ne], 3s2, 3p5 Ar [Ne], 3s2, 3p6

68. H 1s1 He 1s2
Li [He], 2s1 Be [He], 2s2
B [He], 2s2, 2p1 Ne [He], 2s2, 2p6
Na [He], 2s2, 2p6, 3s1  [Ne], 3s1
Mg [He], 2s2, 2p6, 3s2  [Ne], 3s2
Al [Ne], 3s2, 3p1 Si [Ne], 3s2, 3p2
P [Ne], 3s2, 3p3 S [Ne], 3s2, 3p4
Cl [Ne], 3s2, 3p5 Ar [Ne], 3s2, 3p6
outermost shell - valence shell
most loosely held electron and are the most important
in determining an element’s properties

69. K [Ar], 4s1 Ca [Ar], 4s2
Sc [Ar], 4s2, 3d1 Ti [Ar], 4s2, 3d2

70. K [Ar], 4s1 Ca [Ar], 4s2
Sc [Ar], 4s2, 3d1 Ca [Ar], 4s2, 3d2
Zn [Ar], 4s2, 3d10 Ga [Ar], 4s2, 3d10, 3p1
Kr [Ar], 4s2, 3d10, 3p6

71. K [Ar], 4s1 Ca [Ar], 4s2
Sc [Ar], 4s2, 3d1 Ca [Ar], 4s2, 3d2
Zn [Ar], 4s2, 3d10 Ga [Ar], 4s2, 3d10, 3p1
Kr [Ar], 4s2, 3d10, 3p6
Anomalous electron configurations
d5 and d10 are lower in energy than expected
Cr [Ar], 4s1, 3d not [Ar], 4s2, 3d4
Cu [Ar], 4s1, 3d10 not [Ar], 4s2, 3d9

72. Electron Configuration of Ions
Electrons lost from the highest energy occupied orbital
of the donor and placed into the lowest unoccupied orbital
of the acceptor (placed according to the Aufbau principle)

73. Electron Configuration of Ions
Electrons lost from the highest energy occupied orbital
of the donor and placed into the lowest unoccupied orbital
of the acceptor (placed according to the Aufbau principle)
Examples:
Na [Ne], 3s1 Na+ [Ne] + e-
Cl [Ne], 3s2, 3p5 + e- Cl- [Ne], 3s2, 3p6
Mg [Ne], 3s2 Mg2+ [Ne]
O [He], 2s2, 2p4 O2- [He], 2s2, 2p6

74. Modern Theories of the Atom - Summary
Wave-particle duality of light and matter
Bohr theory
Quantum (wave) mechanical model
Orbital shapes and energies
Quantum numbers
Electronic configuration in atoms