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Chapter 4
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Chapter 4

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  1. Chapter 4 Arrangement of Electrons in Atoms Chemistry chapter 4

  2. Light • A kind of electromagnetic radiation • A form of energy that exhibits wavelike behavior as it travels through space Chemistry chapter 4

  3. Electromagnetic Spectrum Chemistry chapter 4

  4. Frequency • The number of wave peaks that occur in a given time period • Represented by the letter f or the greek letter nu (n). • One wave per second is 1 Hertz (Hz). Chemistry chapter 4

  5. Wavelength • The distance between peaks • Measured in length units such as m or nm. • Represented by the letter l (lambda) Chemistry chapter 4

  6. Frequency and wavelength Chemistry chapter 4

  7. The speed of light • c is the speed of light (and all electromagnetic waves) • It’s value is 3.0 x 108 m/s • Frequency and wavelength are related by the equation • c = lf Chemistry chapter 4

  8. Photoelectric effect • The emission of electrons from a metal when light shines on the metal • Only works when the light is above a certain frequency. Chemistry chapter 4

  9. Planck’s Hypothesis • Quantum – the smallest amount of energy that can be lost or gained by an atom. • Quanta of radiant energy are called photons. • E=hf • h= 6.63 x 10-34 J∙s Chemistry chapter 4

  10. Wave-Particle Duality • Introduced by Einstein in 1905 • Light exhibits particle-like properties as well as wave-like ones. Chemistry chapter 4

  11. Photon • Particle of electromagnetic radiation that has zero mass and carries a quantum of energy. Chemistry chapter 4

  12. Photoelectric effect • Einstein explained • In order for an electron to be knocked loose, it must be struck by a single photon with a high enough energy. • This requires a high enough frequency. Chemistry chapter 4

  13. Ground State • The lowest energy state of an atom Chemistry chapter 4

  14. Excited State • A state in which an atom has a higher potential energy than its ground state. • When an atom returns to its ground state, it gives of the extra energy as EM radiation. Chemistry chapter 4

  15. Continuous spectrum • Continuous range of frequencies (colors) • Expected from classical theory Chemistry chapter 4

  16. Line-emission spectrum • Separated bands of light • Have different frequencies • Produced by passing light through a thin slit. • Different elements have different spectra. Chemistry chapter 4

  17. Chemistry chapter 4

  18. Bohr model • Used quantum theory to explain line spectra. • Electrons only exist in specific energy states called orbitals. • Since the change in energy from one state to another is fixed, only certain frequencies are emitted. Chemistry chapter 4

  19. Bohr model • Successful for the hydrogen atom • Needs tweaking for other atoms Chemistry chapter 4

  20. De Broglie hypothesis • 1923 • Louis De Broglie’s dissertation • Planck’s quantum theory implied that light, which had formerly been thought of as a wave, behaves as a particle. • De Broglie hypothesized that the reverse is true. Chemistry chapter 4

  21. Interference • Occurs when waves overlap • Results in a reduction of energy in some areas and an increase of energy in others Chemistry chapter 4

  22. A two-slit light diffraction-interference pattern A two-slit electron diffraction-interference pattern Chemistry chapter 4

  23. Diffraction pattern from X-rays with l = 7.1 x 10-11 m Chemistry chapter 4

  24. Diffraction pattern from 600 eV electrons with l = 5.0 x 10-11 m Chemistry chapter 4

  25. Diffraction pattern from .0568 eV neutrons with l = 1.2 x 10-10 m Chemistry chapter 4

  26. Chemistry chapter 4 An electron microscope

  27. An electron micrograph of DNA Chemistry chapter 4

  28. Heisenberg • Pointed out that it is impossible to know both the exact position and the exact momentum of an object at the same time. • By measuring one, we change the other. Chemistry chapter 4

  29. Measuring position • If we measure the position of an object by hitting it with a photon of energy, the collision with the photon changes its momentum. Chemistry chapter 4

  30. Measuring momentum • If we measure an objects momentum by observing its collision with another object, we have altered its position. Chemistry chapter 4

  31. Schrödinger • Heisenberg treated the electron as a particle. • Schrödinger treated it as a wave • Formulated a difficult wave equation with solutions called wave functions Chemistry chapter 4

  32. Quantum theory • Describes mathematically the wave properties of electrons and other very small particles Chemistry chapter 4

  33. Probability • Wave functions only give the probability of finding an electron at a given place • Electrons don’t travel in neat orbits • “God doesn’t play dice” - Einstein Chemistry chapter 4

  34. Orbital • A 3D region around the nucleus that indicates the probable location of an electron Chemistry chapter 4

  35. Discuss • What is quantum theory? • What radical new idea did de Broglie introduce? • What is interference? Chemistry chapter 4

  36. Omaha zip codes • 681 - Chemistry chapter 4

  37. Four quantum numbers • Specify the properties of atomic orbitals and the properties of electrons in orbitals • Each electron in an atom has a unique set of quantum numbers Chemistry chapter 4

  38. Principal quantum number, n • Indicates the main energy level occupied by the electron • Start numbering with 1 at the level closest to the nucleus. Chemistry chapter 4

  39. Electrons in a given level • The greatest number of electrons that can be in a given level is calculated by the formula 2n2. • So, in the first level there can be 2 ∙ 12 = 2 electrons. Chemistry chapter 4

  40. The angular momentum quantum number, l • Indicates the shape of the orbital • Allowed values: 0, 1, 2, … n – 1 Chemistry chapter 4

  41. s orbital • When l = 0 • spherical Chemistry chapter 4

  42. p orbital • When l = 1 • Dumbbell shaped Chemistry chapter 4

  43. d orbital • When l = 2 • More complex Chemistry chapter 4

  44. f orbital • When l = 3 • Too complex for this class Chemistry chapter 4

  45. Sublevels • Each orbital is designated by the principal quantum number and the orbital letter. • Examples: • 1s • 2p • 4f Chemistry chapter 4

  46. Magnetic quantum number, m • Indicates the orientation of an orbital around the nucleus • Allowed values: -l to l Chemistry chapter 4

  47. s orbital • Each s sublevel only has one s orbital • m = 0 Chemistry chapter 4

  48. p orbitals • Each p sublevel has 3 different p orbitals • m = -1, m = 0, m = 1 Chemistry chapter 4

  49. d orbitals • There are 5 different d orbitals in each d sublevel • m = -2, m = -1, m = 0, m = 1, m = 2 Chemistry chapter 4

  50. Spin quantum number • Has only two values, + ½ and – ½ • Indicate the two fundamental spin states of an electron • Spin up or spin down Chemistry chapter 4