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Chapter 6 Modern Atomic Theory

Chapter 6 Modern Atomic Theory. Review…. Dalton Thomson Rutherford Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton. Electromagnetic Radiation. Light travels in Light is a form of Form of energy that exhibits .

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Chapter 6 Modern Atomic Theory

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  1. Chapter 6Modern Atomic Theory

  2. Review… • Dalton • Thomson • Rutherford • Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

  3. Electromagnetic Radiation • Light travels in • Light is a form of • Form of energy that exhibits

  4. Electromagnetic Radiation • All waves can be described in 3 ways: • Amplitude – • Wavelength (l): • Frequency (n):

  5. Electromagnetic Radiation • Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of • Since light moves at constant speed there is a relationship between wavelength and frequency: Wavelength and frequency are inversely proportional

  6. Electromagnetic Spectrum

  7. Photoelectric Effect • The emission of • Albert Einstein (1905) used Planck’s equation to explain this phenomenon; • proposed that light consists of • Photon =

  8. Photoelectric Effect • He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore • Analogy: • 70 cents placed in soda machine: no soda • 30 cents more and you will get your soda

  9. Niels Henrik David Bohr • 1885-1962 • Physicist • Worked with Rutherford • 1912 • Studying line spectra • of hydrogen

  10. Niels Henrik David Bohr • 1913 – proposed new atomic structure • Electrons exist in • Electrons

  11. The Bohr Atom • Nucleus with • Electrons move in • When an electron moves from one state to another the energy lost or gained is in • Each line in a spectrum is produced when an electron moves from

  12. The Bohr Atom • Model didn’t seem to work with atoms with more than one electron • Did not explain chemical behavior of the atoms

  13. Now… • Light can be described as • What does this mean for the atom???

  14. LineSpectrum • Elements in gaseous states give off colored light • High temperature or high voltage • Always the same • Each element is unique • Spectra

  15. Line Spectrum • Ground state • Excited state

  16. Line Spectrum • Electron • Color of light emitted depends on

  17. Line Spectrum • Each band of color is produced by light of a different • Each particular wavelength has a definite • Each line must therefore be produced by emission of photons with

  18. Line Spectrum • Whenever an excited electron • The energy of this photon is equal to the difference

  19. Wave Matters… • Louis de Broglie (1924) • Proposed that electrons might have a • Used observations of normal wave activity

  20. Problems… • Wave theory does not explain • Heated iron gives off heat • 1stred glow yellow glow white glow • How elements such as barium and strontium give rise to green and red colors when heated

  21. Beginnings… • Max Planck (1858-1947) • Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy • Energy is released in

  22. Beginnings • A quantum is a finite quantity of energy that can be gained or lost by an atom • This constant, h, is the same for all electromagnetic radiation

  23. Bohr’s Equation • En = (-RH)(1/n2) • Where RH = 2.18 x 10-18J • And n = principal quantum number, 1 to infinity

  24. Jumping electrons… • If an electron moves from one energy level to another, the change in energy can be determined by the following equation: • E = Ef – Ei = hν • Or simply: E = hv • Where h=6.626 x 10-34 J s

  25. Then… by substitution… ( E RH ( 1 1 - = ν = ni2 nf2 h h

  26. Finally… Matter waves • All moving particles • Some is apparent, some not. • De Broglie’s equation h λ = mν

  27. Smart guy… • Erwin Schrodinger (1926) • Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves • Quantum Theory

  28. Uncertainty principle • Heisenberg:

  29. Quantum Theory • Describes mathematically the wave properties of electrons and other very small particles • Applies to all elements (not just H)

  30. Quantum Numbers • Numbers that specify the • Principle Quantum Number: • Symbolized by n,

  31. Energy Levels of Electrons • Principle energy levels • Designated by letter n • Corresponds to the • Each level divided into sublevels • 1st energy level has • 2nd energy level has • Etc.

  32. Orbitals • Electrons don’t • Orbital: region in space where • Each orbital sublevel can hold

  33. Orbitals Each sublevel (orbital) has a specific shape http://daugerresearch.com/orbitals/

  34. Quantum Numbers • Orbital Quantum Number: • Indicates the shape of an orbital • (subshell or sublevels) • s, p, d, f Principal Quantum # Orbital Quantum # 1 2 3 4

  35. Quantum Numbers • Magnetic Quantum Number: • Indicates the • Orbital position with respect to

  36. Orbitron • For a full view of the different orbital shapes, visit • http://www.shef.ac.uk/chemistry/orbitron/index.html

  37. Orbitals • Pauli exclusion principle: • Electrons can only spin • Shown with

  38. Rules for Orbital Filling • Pauli’s Exclusion Rule • No two electrons have • Hund’s Rule • Electrons will remain 1s 2s 2p 3s 3p

  39. Rules for Orbital Filling • Diagonal Rule • The order of filling once the d & f sublevels are being filled • Due to energy levels

  40. Rules for Orbital Filling

  41. Application of Quantum Numbers • Several ways of writing the address or location of an electron • Lowest energy levels are filled first • Electron Configuration: 12C: 32S:

  42. Application of Quantum Numbers • Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons • See examples on whiteboard

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