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Chapter 6: Electronic Structure of Atoms

Chapter 6: Electronic Structure of Atoms. Light is a form of electromagnetic radiation (EMR) :. an oscillating charge, such as an electron, gives rise to electromagnetic radiation:. Electric Field. Magnetic Field. Chapter 6: Electronic Structure of Atoms.

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Chapter 6: Electronic Structure of Atoms

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  1. Chapter 6: Electronic Structure of Atoms Light is a form of electromagnetic radiation (EMR): • an oscillating charge, such as an electron, gives rise to electromagnetic radiation: Electric Field Magnetic Field

  2. Chapter 6: Electronic Structure of Atoms • Both the Electric and the Magnetic field propagate through • space • In vacuum, both move at the speed of light(3 x 108 m/s)

  3. Chapter 6: Electronic Structure of Atoms • Electromagnetic radiation is characterized by • wavelength (), or frequency () and • amplitude (A) l A = intensity l l

  4. Chapter 6: Electronic Structure of Atoms Frequency (n) measures how many wavelengths pass a point per second: 4 xl ÷ 1 s = 4 s-1= 4 Hz 1 s

  5. Chapter 6: Electronic Structure of Atoms Electromagnetic radiation travels at the speed of light: c = 3 x 108 m s-1 Relation between wavelength, frequency, and amplitude: c =l n

  6. Chapter 6: Electronic Structure of Atoms The Electromagnetic Spectrum 400 nm 750 nm

  7. wavelength (l) frequency (n) energy (E) Chapter 6: Electronic Structure of Atoms RedOrangeYellowGreenBlueUltraviolet

  8. Chapter 6: Electronic Structure of Atoms A certain type of laser emits green light of 532 nm. What frequency does this wavelength correspond to?

  9. Chapter 6: Electronic Structure of Atoms Photoelectric Effect Albert Einstein (1879-1955) e- e- e-

  10. Chapter 6: Electronic Structure of Atoms Photoelectric Effect Albert Einstein (1879-1955) e- e- e- e- • Light of a certain minimum frequency is required to dislodge electrons from metals

  11. Max Planck (1858 - 1947) Chapter 6: Electronic Structure of Atoms frequency E = h  Planck’s constant = 6.63 x 10-34 J s

  12. Chapter 6: Electronic Structure of Atoms • Energy of light is related to its frequency, not intensity E = h  • light comes in “units” or packets of “h” • Intensity is related only to the number of “units” • The h “unit” is called a quantum of energy • A quantum of light (EMR) energy = photon

  13. Chapter 6: Electronic Structure of Atoms Electromagnetic Radiation stream of particles (photons) wave or E = h n Whether light behaves as a wave or as a stream of photons depends on themethod used to investigate it !

  14. Chapter 6: Electronic Structure of Atoms Relationship between Energy, Wavelength, and Frequency:

  15. Chapter 6: Electronic Structure of Atoms What is the energy of a photon of light of 532 nm? = 3.74 x 10-19 J

  16. Chapter 6: Electronic Structure of Atoms Understanding light in terms of photons helped understand atomic structure many light sources produce a continuous spectrum

  17. Chapter 6: Electronic Structure of Atoms Thermally excited atoms in the gas phase emit line spectra continuous spectrum (all wavelengths together: white light) line spectrum (only some wavelengths: emission will have a color)

  18. Rydberg constant 1.097 x 107 m-1 positive integers (e.g. 1,2,3, etc) Chapter 6: Electronic Structure of Atoms Photograph of the H2 line spectrum (Balmer series) in the visible region (1825-1898) Johann Balmer (1825-1898)

  19. Chapter 6: Electronic Structure of Atoms Thermally excited atoms in the gas phase emit line spectra continuous spectrum (all wavelengths together: white light) line spectrum (only some wavelengths: emission will have a color)

  20. Chapter 6: Electronic Structure of Atoms Niels Bohr was the first to offer an explanation for line spectra Bohr Model of the Hydrogen Atom • Only orbits of defined energy and radii are permitted in the hydrogen atom • Energy is absorbed or emitted by the electron as the electron moves from one allowed orbit into another. Energy is absorbed or emitted as a photon of E = hn

  21. (1885-1962) Chapter 6: Electronic Structure of Atoms Niels Bohr was the first to offer an explanation for line spectra electron orbits n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 nucleus Bohr’s Model of the Hydrogen Atom

  22. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms Bohr’s Model of the Hydrogen Atom Energy absorption of a photon e Ground State nucleus

  23. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms Bohr’s Model of the Hydrogen Atom Energy “excited state” e Ground State nucleus

  24. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms Bohr’s Model of the Hydrogen Atom Energy e Ground State nucleus

  25. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms Bohr’s Model of the Hydrogen Atom Energy e Ground State emission of a photon nucleus

  26. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms Which of these transitions represents an absorption process? a [b and c are emission processes] (a) (b) (c) Energy Which of these transitions involves the largest change in energy? c Which of these transitions leads to the emission of the longest wavelength photon? b smallest distance = lowest energy = long l Ground State Does this wavelength correspond to a high or low frequency? long l = low energy = low frequency nucleus

  27. Transitions corresponding to the Balmer series n=3 → n=2 n=4 → n=2 n=6 → n=2 n=5 → n=2 Chapter 6: Electronic Structure of Atoms

  28. Chapter 6: Electronic Structure of Atoms

  29. Energy of electron in a given orbit: n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms n = Principal Quantum Number (main energy levels) h=Planck’s constant, c=speed of light, RH = Rydberg constant

  30. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms For an electron moving from n = 4 to n = 2: DE = - 4.09 x 10-19 J energy decreases = emission of a photon

  31. n = 6 n = 5 n = 4 n = 3 n = 2 n = 1 Chapter 6: Electronic Structure of Atoms The energy of the photon emitted is: E = 4.09 x 10-19 J What wavelength (in nm) does this energy correspond to? l = 486 x 10-9 m = 486 nm

  32. n=3 → n=2 n=4 → n=2 n=6 → n=2 n=5 → n=2 Chapter 6: Electronic Structure of Atoms Balmer Series l = 486 nm

  33. Chapter 6: Electronic Structure of Atoms The Uncertainty Principle Werner Heisenberg (1901-1976) and Niels Bohr

  34. Chapter 6: Electronic Structure of Atoms The Uncertainty Principle It is impossible to know both the exact position and the exact momentum of a subatomic particle

  35. Chapter 6: Electronic Structure of Atoms Quantum Mechanics and Atomic Orbitals Erwin Schrödinger (1887-1961)

  36. Chapter 6: Electronic Structure of Atoms Quantum Mechanics and Atomic Orbitals • Schrödinger proposed wave mechanical model of the atom • Electrons are described by a wave function, ψ • The square of the wave function, ψ2, provides information on • the location of an electron (probability density or electron density)

  37. Chapter 6: Electronic Structure of Atoms Quantum Mechanics and Atomic Orbitals • the denser the stippling, the • higher the probability of finding • the electron • shape of electron density • regions depends on energy of • electron

  38. z y x Chapter 6: Electronic Structure of Atoms Bohr’s model: n = 1 orbit electron circles around nucleus Schrödinger’s model: orbital n = 1 or electron is somewhere within that spherical region

  39. Chapter 6: Electronic Structure of Atoms Bohr’s model: • requires only a single quantum number (n) to describe an orbit Schrödinger’s model: • requires three quantum numbers (n, l, and m) to describe an orbital n: principal quantum number l : second or azimuthal quantum number ml: magnetic quantum number

  40. - energy of electron in a given orbital: Chapter 6: Electronic Structure of Atoms Schrödinger’s model: (1) n = principal quantum number (analogous to Bohr model) - the higher n, the higher the energy of the electron - is always a positive integer: 1, 2, 3, 4 ….

  41. - lis normally listed as a letter: Value of l: 0 1 2 3 letter: spdf Chapter 6: Electronic Structure of Atoms Schrödinger’s model: (2)l = azimuthal quantum number - takes integral values from 0 to n-1 e.g. n = 3 l = 0, 1, 2 - ldefines the shape of an electron orbital

  42. p-orbital (1 of 3) d-orbital (1 of 5) f-orbital (1 of 7) Chapter 6: Electronic Structure of Atoms Schrödinger’s model: z y x s-orbital

  43. Chapter 6: Electronic Structure of Atoms Schrödinger’s model: (3) ml = magnetic quantum number - takes integral values from -lto +l, including 0 e.g. l = 2 ml= -2, -1, 0, 1, 2 - mldescribes the orientation of an electron orbital in space

  44. n=3 shell 4f subshell Chapter 6: Electronic Structure of Atoms 3 n = 1 2 4 l = 0 0, 1 0, 1, 2 0, 1, 2, 3 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f ml = 0; -1,0,1; 0 -1,0,1 -2,-1,0,1,2 -3,-2,-1,0,1,2,3 0; 0; -1,0,1; -2,-1,0,1,2; # orbitals in subshell 1 1 3 3 5 1 3 5 1 7 Total # of orbitals in shell 1 4 9 16

  45. Chapter 6: Electronic Structure of Atoms 3s-room 3p-room 3deluxe-room 3rd floor 2s-room 2promotion-room 2nd floor standard-room 1st floor

  46. Chapter 6: Electronic Structure of Atoms Orbital energy levels in the Hydrogen Atom

  47. Chapter 6: Electronic Structure of Atoms Representation of Orbitals 1s 2s 3s

  48. Chapter 6: Electronic Structure of Atoms Representation of Orbitals s-orbital

  49. Chapter 6: Electronic Structure of Atoms Representation of Orbitals p-orbitals

  50. Chapter 6: Electronic Structure of Atoms Representation of Orbitals d-orbitals

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