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Quantum Chemistry. Dr. Ron Rusay. Atomic Structure and Periodicity. Electromagnetic Radiation The Nature of Matter The Atomic Spectrum of Hydrogen The Bohr Model The Quantum Mechanical Model of the Atom Quantum Numbers Orbital Shapes and Energies Electron Spin and the Pauli Principle

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Quantum Chemistry


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    1. Quantum Chemistry Dr. Ron Rusay

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

    7. Frequency & Wave length

    8. Waveshttp://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  108 m/s.

    9. Waveshttp://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  108 m/s.

    10. Wavelength and frequency  = c /  •  = frequency (s1) •  = wavelength (m) • c = speed of light (m s1)

    11. QUESTION

    12. 1 B) 4.12 10 s The smaller the frequency of light, the longer the wavelength. ANSWER 5 – ´

    13. Planck’s Constant Transfer of energy is quantized, and can only occur in discrete units, called quanta. • 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

    14. Planck’s Equation (Interactive)

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

    16. Energy and Mass • Energy has mass • E = mc2 • E = energy • m = mass • c = speed of light

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

    18. Wavelength and Mass de Broglie’s Equation •  = wavelength, in m • h = Planck’s constant, 6.626  1034J .s = kg m2 s1 • m = mass, in kg •  = frequency, in s1

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

    20. Absorption & Emission

    21. Emissions: Flame Tests

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

    23. Atomic Emission Spectrum of H2

    24. 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

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

    26. Energy Changes in the Hydrogen Atom • E = Efinal state Einitial state

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

    28. 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 lor)= 0 to n  1) : relates to shape of the orbital. • 3. Magnetic QN (integer m l orm= + l to  l) : relates to orientation of the orbital in space relative to other orbitals. • 4. Electron Spin QN : (ms = +1/2, 1/2) : relates to the spin state of the electron.

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

    30. Periodic Table ClassificationsElectron 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)

    31. 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.

    32. QUESTION

    33. B) 4 Orbitals are designated by . For = 2, has m four values. ANSWER n m l l

    34. QUESTION

    35. A) 0 For = 3, can be 0, 1, or 2. An orbital has an n = 3. ANSWER l f l

    36. Quantum Numbers : l, mlOrbital Shape & Orientation

    37. Magnetic Spin ms

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