1. Quantum Chemistry Dr. Ron Rusay

2. Atomic Structure and Periodicity ð ð Electromagnetic Radiation ð ð The Nature of Matter ð ð The Atomic Spectrum of Hydrogen ð ð The Bohr Model ð ð The Quantum Mechanical Model of the Atom ð ð Quantum Numbers ð ð Orbital Shapes and Energies ð ð Electron Spin and the Pauli Principle ð ð Polyelectronic Atoms ð ð The History of the Periodic Table ð ð The Aufbau Principles and the Periodic Table ð ð Periodic Trends in Atomic Properties ð ð The Properties of a Group: The Alkali Metals

3. Quantum Theory  Based on experimental observations of light and particles  Development progressed through rigorous mathematical computations  It bridges physics and chemistry  It is described generally as quantum mechanics

4. Electromagnetic Radiation (“Light”)  Energy that exhibits wave-like behavior.  In a vacuum, electromagnetic energy travels through space at the speed of light.  It is described by the Electromagnetic Spectrum.

5. Nature of EM Energy

6. Demonstrating Light’s Wave Nature

7. Frequency & Wave length

9. Waves http://chemistry.beloit.edu/BlueLight/waves/index.html http://chemistry.beloit.edu/BlueLight/waves/index.html  Waves have 4 primary characteristics:  1.Wavelength: distance between two peaks in a wave.  2.Frequency: number of waves per second that pass a given point in space.  3.Amplitude: the height of the wave.  4. Speed: speed of light is 2.9979  10 8 m/s.

10. Waves http://chemistry.beloit.edu/BlueLight/waves/index.html http://chemistry.beloit.edu/BlueLight/waves/index.html  Focus on 2 of the primary characteristics:  1.Wavelength: distance between two peaks in a wave.  2.Frequency: number of waves per second that pass a given point in space.  3.Amplitude: the height of the wave.  4. Speed: speed of light is 2.9979  10 8 m/s.

11. Wavelength and frequency  = c /  = c /  = frequency (s  1 )  = wavelength (m)  c = speed of light (m s  1 )

12. QUESTION

13. Planck’s Constant   E = change in energy, in J  h = Planck’s constant, 6.626  10  34 J s  = frequency, in s  1  = wavelength, in m  c = speed of light Transfer of energy is quantized, and can only occur in discrete units, called quanta.

15. Planck’s Equation (Interactive)

16. Electromagnetic Energy  EM Spectrum : Chem Connections http://chemistry.beloit.edu/Stars/EMSpectrum/index.html

17. Energy and Mass  Energy has mass  E = mc 2  E = energy  m = mass  c = speed of light

18. Energy and Mass ”Duality” (Hence the dual nature of light.)

19. Wavelength and Mass  = wavelength, in m  h = Planck’s constant, 6.626  10  34 J. s = kg m 2 s  1  m = mass, in kg  = frequency, in s  1 de Broglie’s Equation

20. Atomic Spectrum of Hydrogen  Continuous spectrum: Contains all the wavelengths of light.  Absorbtion vs.Emission  http://chemistry.beloit.edu/BlueLight/pages/elements.html http://chemistry.beloit.edu/BlueLight/pages/elements.html  Line (discrete) spectrum: Contains only some of the wavelengths of light. http://chemistry.beloit.edu/BlueLight/pages/color.html

22. Absorption & Emission

23. Emissions: Flame Tests

25. Electromagnetic Energy  Visible Light / Color : ChemConnections  http://chemistry.beloit.edu/Stars/applets/emission/index.html http://chemistry.beloit.edu/Stars/applets/emission/index.html  The Perception of Colors  http://chemconnections.org/organicchem227/227assign-06.html#vision http://chemconnections.org/organicchem227/227assign-06.html#vision

26. Atomic Emission Spectrum of H 2

29. The Bohr Model  E = energy of the levels in the H-atom  z = nuclear charge (for H, z = 1)  n = an integer “The electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.” X

30. The Bohr Model  Ground State: The lowest possible energy state for an atom (n = 1).

31. Energy Changes in the Hydrogen Atom   E = E final state  E initial state

33. Heisenberg Uncertainty Principle  The more accurately we know a particle’s position, the less accurately we can know its momentum or vice versa. Quantum Entanglement/Superposition Schrödinger’s Cat: Alive or Dead? Can something be in two places at the same time? In quantum microstates, YES. Science, 272, 1132 (1996)

34. Quantum Numbers (QN) for Electrons (Solutions for the Schrödinger Equation:  =  ) Where:  = Wave function  1.Principal QN ( integer n = 1, 2, 3,...) : relates to size and energy of the orbital.  2.Angular Momentum QN ( integer l or )= 0 to n  1) : relates to shape of the orbital.  3.Magnetic QN (integer m l or m = + l to  l) : relates to orientation of the orbital in space relative to other orbitals.  4.Electron Spin QN : (m s = + 1 / 2,  1 / 2 ) : relates to the spin state of the electron.

35. Electron Probability = |  | 2 |  | 2 =  (double integral of wave function  ) “ORBITAL”:

36. Periodic Table Classifications Electron Configurations & Quantum Numbers  Representative Elements (A Groups): s (l=0) and p (l=1) (N, C, Al, Ne, F, O)  Transition Elements: d (l=2) orbitals (Fe, Co, Ni, etc.)  Lanthanide and Actinide Series (inner transition elements): f (l=3) orbitals (Eu, Am, Es)

37. Valence Electrons Valence electrons are the outermost electrons in the highest principal quantum level of an atom. They are found in the s- and p- orbitals and are the bonding electrons. Inner electrons are called core electrons.

39. QUESTION

40. QUESTION

41. Quantum Numbers : l, m l Orbital Shape & Orientation

42. Magnetic Spin m s

43. Electron Probability = |  | 2 |  | 2 =  (double integral of wave function  )

48. Atomic Orbitals  See the following Web page: Identify the unknown orbitals by comparing their shapes to the known orbitals and assign quantum numbers to each orbital. http://chemconnections.org/general/chem120/atomic-orbitals/orbitals.html

50. Multi-electron Atoms Electron Configuration

52. Aufbau Principle  As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.

53. Full electron configuration (Spectroscopic notation) --->

54. QUESTION

55. Pauli Exclusion Principle  In a given atom, no two electrons can have the same set of four quantum numbers ( n, l, m l, m s ).  Therefore, an orbital can hold only two electrons, and they must have opposite spins.

56. QUESTION

57. Hund’s Rule orbital diagrams  The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. Orbital Diagram ->

60. Periodic Table Classifications Electron Configurations  Representative Elements (A Groups): fill s and p orbitals (Na, Al, Ne, O)  Transition Elements: fill d orbitals (Fe, Co, Ni)  Lanthanide and Actinide Series (inner transition elements): fill 4f and 5f orbitals (Eu, Am, Es)

61. Valence Electrons Valence electrons are the outermost electrons in the highest principal quantum level of an atom. They are found in the s- and p- orbitals and are the bonding electrons. Inner electrons are called core electrons.

62. QUESTION

63. QUESTION

64. Two ways of showing the formation of lithium fluoride: LiF; [Li + and F - ] using electron configurations & diagrams

65. QUESTION

66. Paramagnetism & Diamagnetism Electron Configuration & Magnetic Properties Diamagnetic materials have all electrons paired and are not attracted to a magnetic field.Diamagnetic materials have all electrons paired and are not attracted to a magnetic field. Paramagnetic materials have unpaired electrons and the magnetic attraction (magnetism) is generally proportional to the number of unpaired electrons. (Note: not all metals follow this rule.)Paramagnetic materials have unpaired electrons and the magnetic attraction (magnetism) is generally proportional to the number of unpaired electrons. (Note: not all metals follow this rule.)

67. Electron Diagrams Magnetic Properties #1 = H 2 O(l) # 2 = Fe 2 O 3 (s) # 3 = FeO(s) #4= Fe(s)

68. Transition Metal Ions (B Groups) Oxidation Numbers (States)

69. Isoelectronic atoms and ions have the same electron configurations

70. Apparatus Used to Measure Paramagnetism NOTE: O 2 is paramagnetic, N 2 is not! Also, Ferromagnetic effects are much, much stronger than Paramagnetic

71. Electron Diagrams Magnetic Properties Ground state configurations of nitrogen (N) and oxygen (O) have 3 and 2 unpaired electrons in their electron diagrams respectively, what can be going on in the video?Ground state configurations of nitrogen (N) and oxygen (O) have 3 and 2 unpaired electrons in their electron diagrams respectively, what can be going on in the video? Ground state diagrams do work very well for the Transition metals but not many others because of bonding, which forms pairs of electrons. (molecular orbitals vs. atomic orbitals).Eg. water, nitrogen and oxygen.Ground state diagrams do work very well for the Transition metals but not many others because of bonding, which forms pairs of electrons. (molecular orbitals vs. atomic orbitals).Eg. water, nitrogen and oxygen.

72. Molecular Orbital Diagrams

73. Summary: Information from the Periodic Table  1.Can obtain Group A valence electron configurations  2.Can determine individual electron configurations. This information can be used to:  a.Predict the physical properties and general chemical behavior of the elements.  b.Identify metals and nonmetals.