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Electrons in Atoms

Electrons in Atoms. Golden Valley High School Chapter 13. Created by David Galaz, modified by Stacey Cool. Summarize the development of Atomic Theory through Bohr’s model Understand the importance of quantized energies for electron placement.

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Electrons in Atoms

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  1. Electrons in Atoms Golden Valley High School Chapter 13 Created by David Galaz, modified by Stacey Cool Chapter 13

  2. Summarize the development of Atomic Theory through Bohr’s model • Understand the importance of quantized energies for electron placement Chapter 13

  3. Write the electron configurations of various element using the Aufbau principle, Pali exclusion principle and .Hund’s rule • Explain exceptions to the Aufbau Principle. Chapter 13

  4. Dalton’s Atomic Theory (4 parts) Good stuff: Atoms combine in whole number ratios Chemical reactions occur when atoms are joined, separated or rearranged 13.1 – Models of the Atom Chapter 13

  5. Dalton’s Atomic Theory (4 parts) Not so good stuff: Part 1: Atoms canbe separated: subatomic particles Part 2: All atoms are notidentical: isotopes 13.1 – Models of the Atom Chapter 13

  6. Thomson Atomic Theory Post discovery of the electron Also known as: the Plum Pudding model 13.1 – Models of the Atom Chapter 13

  7. Thomson Atomic Theory Good stuff: accounts for both electrons and protons Not so good stuff: Nothing about the arrangement or the number of p+, e-, or no 13.1 – Models of the Atom Chapter 13

  8. 13.1 – Models of the Atom • Ernest Rutherford • Post discovery of the nucleus • A.K.A. Nuclear atomic model Chapter 13

  9. 13.1 – Models of the Atom • Ernest Rutherford • Proposed two things: • Nuclear atom, with e- surrounding the positive center • Dense nucleus, with rest of atom mostly empty space Chapter 13

  10. 13.1 – Models of the Atom • Ernest Rutherford • Problem: No explanation as to why e- do not collapse onto the nucleus. Chapter 13

  11. 13.1 – Models of the Atom • Bohr • AKA Planetary model • Proposed: e- arranged in orbits around the nucleus Chapter 13

  12. 13.1 – Models of the Atom • Bohr • First to answer the question: • e- in a particular path have a fixed energy level • No loss / gain of energy, so no collapse Chapter 13

  13. 13.1 – Models of the Atom • Energy level of an e-: • The region around the nucleus where the e- is likely to be moving • Fixed energy level of electrons • Similar to rungs on the ladder Chapter 13

  14. 13.1 – Models of the Atom • Lowest energy level = lowest step • Higher the energy level, the farther an e- is from the nucleus • e- can change position or energy levels Chapter 13

  15. 13.1 – Models of the Atom • Quantum of energy is the amount of energy required to move an electron from its present energy level to the next higher one Chapter 13

  16. 13.1 – Models of the Atom Quantum Mechanical Model • 1926 Schrodinger • Modern description of the atom • Based on statistical probability of the location of an electron • Does not define the e- path, it estimates the probability Chapter 13

  17. A.K.A. fuzzy cloud Dense area = high probability 13.1 – Models of the Atom Chapter 13

  18. 13.1 – Models of the Atom • Energy levels are designated by principal quantum numbers (n) • Range = 1 – 7 • Greater the #, the farther from the nucleus Chapter 13

  19. 13.1 – Models of the Atom Energy Sublevels • Located within each principal quantum number (n) • Up to 4 sublevels: s, p, d, f • Maximum number of e- that occupy (n) = 2n2 Chapter 13

  20. Energy Sublevels Chapter 13

  21. Atomic orbitals2e- per orbitals Chapter 13

  22. Chapter 13

  23. 13.2 – Electron Configuration • Definition: Electron arrangement around the nucleus of an atom Rules to determine (3) Aufbau principle Pauli exclusion Hund’s rule Chapter 13

  24. Electrons enter orbitals of lowest energy first Orbitals within a sublevel are of equal energy Sublevels located horizontally, lowest  highest, left  right Aufbau Principle Chapter 13

  25. Pauli Exclusion Principle • An atomic orbital may describe at most two electrons • Two electrons in the same orbit must have opposite spins • (clockwise/counter clockwise) • Vertical arrows represent an electron’s spin Chapter 13

  26. Hund’s Rule • Definition: • When electrons occupy orbitals of equal energy (p,d,f), one e- enters each orbital until all the orbitals contain one e- with parallel spins Chapter 13

  27. Pauli Exclusion Principle • An atomic orbital may describe at most two electrons • Two electrons in the same orbit must have opposite spins • (clockwise/counter clockwise) • Vertical arrows represent an electron’s spin Chapter 13

  28. Chapter 13

  29. Atomic orbitals2e- per orbitals Chapter 13

  30. s & p = row number d = row number -1 • f = row number -2 Chapter 13

  31. Hydrogen Helium Lithium Carbon Fluorine Neon Determine the e- configuration & draw orbital diagrams • Silicon • Vanadium • Strontium Chapter 13

  32. http://www.youtube.com/watch?v=Vb6kAxwSWgU Chapter 13

  33. Review Slip • 1. Write the electron configuration for copper • 2. Rate how you feel about your knowledge? • (1 to 5) One being: HuH? Five being: I can teach this course. • 3. Do you have a question you need answered? Chapter 13

  34. 13.3 – Physics & the QMM • Electromagnetic waves • Speed of light (in a vacuum) • Symbol = (c) • Value = 3.0 x 108 m/s Chapter 13

  35. 13.3 – Physics & the QMM Energy Wave – there are 6 parts • Origin: Beginning and ending on the base line • Crest: highest point of a wave • Trough: Lowest point of a wave Chapter 13

  36. 13.3 – Physics & the QMM Energy Wave – there are 6 parts • Amplitude: Height of the wave from the origin to the crest • Wavelength: Distance between crests • Symbol: lamda ( λ ) Chapter 13

  37. 13.3 – Physics & the QMM Energy Wave – there are 6 parts • Frequency: Number of wave cycles to pass a given point per unit of time • Symbol: Greek letter nu ( ν ) *Frequency and Wavelength are inversely related Chapter 13

  38. Frequency: Number of wave cycles to pass a given point per unit of time. Units = Hertz (Hz) 1 Hz = 1s-1 Formula: ν = c λ Example problems: Calculate the frequency, if the wave length is equal to 8.0 x 105 m. What is the frequency of light when λ = 1.0 x 108 m? 13.3 – Physics & the QMM Chapter 13

  39. What is the wave length, when the frequency is equal to 7.3 x 10-4Hz? Calculate the wavelength if ν = 3.5 x 10–9 Hz. ν = cλ Chapter 13

  40. Longest (λ) to shortest Radio waves Micro waves Infrared Visible light Ultra violet light X – rays Gamma rays Greatest (ν) to smallest Electromagnetic RadiationSeven parts Chapter 13

  41. Visible light (7 parts) Longest (λ) to shortest Red Orange Yellow Green Blue Indigo Violet Greatest (ν) to smallest Spectrum of Colors Chapter 13

  42. Section 12.7: The Quantum ConceptObjective: Why metals change color when heated? 1900s, German physicist Quantitatively Color changes occur with small discrete energy changes The amount of radiant energy (E) absorbed / emitted is proportional to the frequency (ν) of radiation • Formula for amount of energy ( E) = h x ν • Planck’s constant = h = 6.6262 x 10-34 J s Chapter 13

  43. Calculate the amount of energy absorbed when the wave length is equal to 5.0 x 10-8 m. Calculate the amount of energy emitted from the southern lights, if the wave length is equal to 3.5 x 10-7m. Chapter 13

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