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14 November 2011

14 November 2011. Objective : You will be able to: describe evidence for the current theory of the electronic structure of atoms. Homework : p. 312 #3, 4, 5, 6, 7, 9, 16, 19, 25, 32. Electronic Structure of Atoms. Next Units:. Electron configuration Trends on the periodic table

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14 November 2011

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  1. 14 November 2011 • Objective: You will be able to: • describe evidence for the current theory of the electronic structure of atoms. • Homework: p. 312 #3, 4, 5, 6, 7, 9, 16, 19, 25, 32

  2. Electronic Structure of Atoms

  3. Next Units: • Electron configuration • Trends on the periodic table • Ionic/covalent bonding • Chemical reactivity

  4. In order to understand these things • we’ll study the electronic structure of atoms

  5. The Wave Nature of Light • electromagnetic radiation (a.k.a. light) is a form of energy with wave and particle characteristics. It moves through a vacuum at the speed of light • speed of light: 3.00x108 m/s

  6. To describe waves… • wavelength (λlamda): the distance between two adjacent peaks of a wave • frequency (v): the number of wavelengths that pass a given point in a second

  7. Electromagnetic Spectrum

  8. electromagnetic spectrum includes all wavelengths of radiant energy • visible spectrum: the part of the electromagnetic spectrum that is visible to the human eye (wavelengths between 400 and 700 nm)

  9. Quantized Energy and Photons • quantum (a.k.a. photon) is a specific particle of light energy that can be emitted or absorbed as electromagnetic radiation. • Energy of a photon E=hv • Energy is quantized – matter is allowed to emit or absorb energy in discrete amounts, whole number multiples of hv.

  10. How are these things related to electromagnetic radiation? E=hc/λ

  11. Example 1 Calculate the energy (in joules) of • a photon with a wavelength of 5.00x104 nm (infrared region) • a photon with a wavelength of 5.00x10-2 nm (x-ray region)

  12. Example 2 • What is the frequency and the energy of a single photon? • What is the energy of a mole of photons of light having a wavelength of 555 nm?

  13. Problem • The energy of a photon is 5.87x10-20 J. What is its wavelength, in nanometers?

  14. Homework • p. 312 #3, 4, 5, 6, 7, 9, 16, 19, 25, 32

  15. 15 November 2011 • Take Out Homework • Objective: You will be able to: • describe and explain experimental evidence for energy levels • Homework Quiz: The energy of a photon is 3.98x10-19 J. What color light do you observe?

  16. Agenda • Homework Quiz • Hand back tests • Line spectra and the Bohr model of the atom Homework: p. 313 #23, 24, 25, 26, 30, 31, 35, 36

  17. Line Spectra and the Bohr Model • atomic emission spectrum (a.k.a. line spectrum): a pattern of discrete lines of different wavelengths that result when the light energy emitted from energized atoms is passed through a prism • Each element produces a characteristic or identifiable pattern

  18. Demo • Emission spectra of common cations • Note: we don’t have a way to separate all the wavelengths of light into discrete lines of color, so we’re just seeing all those lines of color blended together. • http://www.youtube.com/watch?v=2ZlhRChr_Bw&feature=related

  19. So, why do we see these discrete lines of color? • Bohr model of the atom: energies are quantized. Electrons move in circular, fixed energy orbits around the nucleus. • Usually, electrons are in the most stable “ground” state. • When energy (a photon) is added, they “jump” up to the “excited” state • They fall back down, and release that photon.

  20. Multimedia • http://www.youtube.com/watch?v=45KGS1Ro-sc • http://www.colorado.edu/physics/2000/quantumzone/lines2.html

  21. Homework • p. 313 #23, 24, 25, 26, 30, 31, 35, 36

  22. 16 November 2011 • Objective: You will be able to: • explain how line spectra give evidence for the existence of energy levels • explain how quantum mechanics describes electron configuration

  23. Agenda • Homework Quiz • Go over homework • How do atoms emit photons? • Quantum mechanics: how do we describe where the electrons are?! • Writing orbital notation and electron configuration Homework: p. 313 #23-26, 30, 35, 48, 53, 60, 63,

  24. Energy levels

  25. Wave Behavior of Matter • Like light, electrons have characteristics of both waves and particles. Because a wave extends into space, its location is not precisely defined. • uncertainty principle: it is impossible to simultaneously determine the exact position and momentum of an electron. • we can only determine the probability of finding an electron in a certain region of space.

  26. Quantum Mechanics and Atomic Orbitals • quantum mechanical model: mathematical model that incorporates both the wave and particle characteristics of electrons in atoms. • quantum numbers: describe properties of electrons and orbitals • each electron has a series of four quantum numbers

  27. Table of Quantum Numbers

  28. Table of quantum numbers and orbital designations

  29. Pauli Exclusion Principle • Two electrons in an atom can’t have the same four quantum numbers • Two electrons per orbital, with opposite spins

  30. Representations of Orbitals • orbital: calculated probability of finding an electron of a given energy in a region of space

  31. p orbitals

  32. d orbitals

  33. orbital ≠ orbit

  34. 17 November 2011 • Objective: You will be able to: • write the orbital and electron configuration for any element • describe several exceptions to the orbital filling rules • Homework Quiz: Describe, as completely as you can in a paragraph or two, the evidence that convinced Neils Bohr of the existence of energy levels instead of a cloud of electrons.

  35. Agenda • Homework Quiz • Go over homework • Electron configuration notation • Problem Set Unit 4 Quiz Weds.

  36. Atoms with more than one electron • like hydrogen • electron-electron repulsions cause different sublevels to have different energies

  37. Order those orbitals fill

  38. Electron Configuration • distribution of electrons among various orbitals of an atom

  39. Rules for Writing E- Config. • at the ground state • Fill the lowest energy level first. Electrons in the same orbital must have opposite spins. Total number of electrons = atomic number • Only two electrons per orbital! • Do not pair electrons in a orbitals of the same energy until each orbital has one electron of the same spin (Hund’s rule) • Label each sublevel with the energy level number and letter of the sublevel

  40. Examples • phosphorus • calcium • iron

  41. Paired-ness of Electrons • Paramagnetic: an atom having one or more unpaired electrons • Ex: Li, B, C… • Diamagnetic: all electrons in an atom are paired. • Ex:

  42. Excited-State Configuration • has a higher energy than the ground-state electron configuration. • One or more electrons occupy higher energy levels than predicted by the rules • Ex: Iron in an excited state:

  43. Electron Configuration and the Periodic Table • Elements with similar electron configurations arranged in columns

  44. Examples • Write the electron configuration for palladium • Write the electron configuration for osmium

  45. Condensed Electron Config. • shows only the electrons occupying the outermost sublevels • preceded by the symbol for the noble gas in the row above the element • Example: calcium • Example: iodine

  46. Unusual Electron Configs. • Cr and Mo: ground state valence electrons are arranged s1d5 rather than s2d4 • a half filled d orbital is more stable than a more-than-half-filled d orbital • Cu, Ag and Au have s1d10 ground state configs because of the stability of a fill d orbital

  47. 21 November 2011 • Objective: You will be able to: • describe the electronic structure of an atom and make associated calculations.

  48. Agenda • Math with exponents (#6) • Problem set work time Homework: Problem set due tomorrow Quiz Mon. on all electronic structure, calculations, evidence for Bohr’s theory…

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