Содержание
- 2. Picture of the Atom Electromagnetic radiation and Atomic Spectra The Nature of Electron and Atomic Orbitals
- 3. 3.1.1 Atomic concept, 3.1.2 Subatomic particles, 3.1.3 Atomic structure: first ideas Part I
- 4. Dalton Atomic Theory Elements are made of tiny particles called atoms 2. The atoms of a
- 5. J.J. Thomson’s Cathode Tube Charge-to-mass ratio
- 6. The Atom : J. J. Thomson (1856-1940) e/m = -1.76 x 108 C/g Experiment date 1898-1903
- 7. The Atom based on Thomson’s experiment A ray of particles is produced between two metallic electrodes.
- 8. Mass of electron Mass of a single electron e= -1.6x10-19 C m = 9.11 x 10-31
- 9. Rutherford Experiment Ernest Rutherford – 1911 With Thomson Model : a particles should travel through the
- 10. Rutherford Experiment
- 11. The Nucleus Ernest Rutherford – 1911 Conclusion : Dense positive center with electrons far from the
- 12. Modern View
- 13. 3.2. Electromagnetic Radiation and Quantization 3.2.1: Electromagnetic Radiation 3.2.2: Quantization 3.2.3: The Atomic Spectrum of Hydrogen
- 14. Spectrum
- 15. Electromagnetic radiation Light X-ray MRI Microwave Travel like a wave Travel with the speed of light
- 16. Electromagnetic Radiation Electromagnetic Radiation = a way for energy to travel. 2 oscillating fields (H and
- 17. ELECTROMAGNETIC RADIATION
- 18. Electromagnetic Radiation - Characteristics λ = wavelength = distance between two peaks or two troughs in
- 19. Radio in the 909kHz. What wavelength does it correspond to? λ = c/ν = 330 m
- 20. Nature of Matter At the end of the 19th century : Matter ≠ Energy Matter =
- 21. Photoelectric effect When UV radiation hits a metal surface, electrons are ejected – photoelectric effect. (in
- 22. E = h x ν E = 6.63 x 10-34 (J•s) x 3.00 x 10 8
- 23. Dual Nature of Light Energy – Mass relationship : A particle but also a wave :
- 24. De Broglie 1924 λ = h/mν λ Proportional to h/mν H :Planck Constant M : masse
- 25. Diffraction What is the wavelength for an electron? Me = 9.11x10-31 kg Ve = 1.0x107 m/s
- 26. How to test the wave properties of an electron?
- 27. How to test the wave properties of an electron?
- 28. Diffraction When X-rays are scattered by ordered atoms ? Diffraction pattern.
- 29. Conclusion All matter exhibits both particulate and wave properties. Large particles : mainly particle Small particles
- 30. Atomic Spectrum of Hydrogen When a high energy discharge is passed through H2 ? H-H breaks
- 31. Table 3.4. The atomic spectrum of hydrogen
- 33. Atomic Spectrum of Hydrogen Why do we have a line spectrum for H ? Only certain
- 34. 3.3.2: The Bohr Model
- 35. The Bohr Model General Idea : The electron in a hydrogen atom moves around the nucleus
- 36. The Bohr Model Example : Energy emitted from n=6 to ground state : The negative sign
- 37. Wave Function and Atomic Orbitals 3.5.1 Wave properties of matter, Heisenberg uncertainty principle 3.5.2 Wave-functions and
- 38. De Broglie All moving particles have wave properties λ= h mu = Wavelength h = Planck
- 39. 2.2 SCHRONDINGER EQUATION Enter
- 40. Quantum Mechanical Description of the Atom Heisenberg – de Broglie – Schrödinger Only certain circular orbits
- 41. The Schrödinger equation The probability distributions and allowed energy levels for electrons in atoms and molecules
- 42. Schrodinger Wave Equation
- 43. Kinetic Energy of the Electron Motion Potential Energy of the Electron. The result of electrostatic attraction
- 44. Kinetic Energy Potential Energy
- 45. Cartesian and Spherical Coordinate
- 46. The wavefunction Atomic wavefunctions are usually expressed in spherical polar coordinates – give value of Ψ
- 47. Homework-2 Please solve problems ; Chapter 3 6, 9, 10, 12, 14, 16 and 17 Due
- 48. Wave Equation for the Hydrogen Atom – R(r) is radial part of wavefunction Describes electrons density
- 49. Quantum numbers : Quantum numbers : n = principal quantum number : size and energy of
- 50. Radial and Angular Wave Function for 1s derived from Schrodinger Equation
- 51. Plot of Radial Wave Function = f(r)
- 52. s orbitals Size : 1s Energy : 1s Surface of 0 probability = nodal surface /
- 53. Physical Meaning of Orbitals The wave function has no easy physical meaning. The square of the
- 54. a1 = 52.9pm radius at n =1 for hydrogen
- 55. a1 = 52.9pm radius at n =1 for hydrogen
- 56. p orbitals Two lobes separated by a node. Sine function : + and - ? same
- 57. d orbitals 2 different shapes : dxz,dyz,dxy, dx2-y2 and dz2
- 58. f orbitals Very complex shapes
- 59. Schrödinger Equation Each solution ψ of the Schrödinger equation has a specific value for E. A
- 60. Heisenberg uncertainty principle There is a fundamental limitation to just how precisely we can know both
- 61. The Hydrogen Atom : summary The quantum mechanical model : electron = wave Series of wave
- 62. Polyelectronic Model Schrödinger equation can be solved exactly only for hydrogen. Schrödinger equation cannot be solved
- 63. Self-Consistent Field Method http://www.youtube.com/watch?v=UVkTuOwfOh0
- 64. https://www.youtube.com/watch?v=A6DiVspoZ1E Review this link at home
- 65. Many Electron Atoms Electron spin, Aufbau principle, Anomalies in electronic configuration, Structure of Periodic table Part
- 66. Electron Spin and Pauli Principle A 4th quantum number describe the electron : ms : electron
- 67. History of the Periodic Table Dmitri Mendeleev : ми́трий Менделе́ев One of first to arrange known
- 68. The Aufbau Principle Principle to populate orbitals.
- 69. Valence electrons Valence electrons = electrons from the outermost principal quantum level of an atom. Group
- 70. Rules
- 71. Rules After 4s2, we fill 3d. Mn : [Ar]4s23d5 – Fe [Ar]4s23d6 Additional Rules: The (n+1)
- 72. Rules Element above 118 are generally unstable G contain 9 orbitals l = n-1 = 4
- 73. Rules
- 74. Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number
- 75. Pauli Exclusion Principle Pauli Exclusion principle ; no two electrons in an atom can have the
- 76. Penetration Effect Why do we fill s, p and then d?. Core electrons : 1s, 2s
- 77. Penetration Effect The penetration effect also explains why 4s is filled before 3d. Potassium : 1S22S22P63S23P64S1
- 78. 4s 5g 5s 3p 3s 2p 2s Slater rules provide an approximate Guide explain why certain
- 79. https://en.wikipedia.org/wiki/Effective_nuclear_charge
- 80. Slater’s Rules The rules were devised semi-empirically by John C. Slater and published in 1930 Identify
- 81. Slater’s Rules The rules were devised semi-empirically by John C. Slater and published in 1930 Rules
- 82. Slater’s Rules for determining S for a specific electron The shielding constant (S) ns and np
- 83. Slater’s Rules for determining S for a specific electron Rule -3b: Each electron in n-1 group
- 84. Z* for Na = Z – S = 11 – 8.8 = 2.2
- 85. Slater’s Rules for determining S for a specific electron Rule -4a: Each electron in nd and
- 86. Nickel Use slater rules to calculate the shielding constant S and effective nuclear charge of 3d
- 87. Solution Rule-1 : the electron configuration is written using slater’s groupings: (1s2)(2s2, 2p6)(3s2, 3p6)(3d8)(4s2) To calculate
- 88. Solution Rule-1 : the electron configuration is written using slater’s groupings: (1s2)(2s2, 2p6)(3s2, 3p6)(3d8) (4s2) To
- 89. Comparison of The effective nuclear charge 3d electrons The effective nuclear charge Z*=28-20.45=7.55 4s electrons Z*=28-23.95=4.05
- 90. Periodic Properties of Atoms : Ionization Energy Ionization Energy : Energy required to remove an electron
- 91. Trend in Atomic Properties : Ionization Energy IE core electrons >> IE >> from left to
- 94. Trend in Atomic Properties : Ionization Energy Li: 1s22s1 (3 electrons) Be: 1s22s2 (3 electrons) Expected
- 95. Trend in Atomic Properties : Ionization Energy IE goes down along a group. The removed electron
- 96. Trend in Atomic Properties : Atomic Radius Atomic Radius: half the distance between the nuclei in
- 97. END END
- 98. Alkali Metals – 1A Low melting point Lose easily an electron Strong reducer Li > K
- 99. ρsinφ Δρ Angle = Arc Length radius of the circle
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