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Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY. Bohr Model. First model of the electron behavior Vital to understanding the atom Does not work for atoms with more than 1 electron. Collision of Ideas. Matter. Dalton. Thompson. Rutherford. Bohr. ?.

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Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

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  1. Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

  2. Bohr Model • First model of the electron behavior • Vital to understanding the atom • Does not work for atoms with more than 1 electron

  3. Collision of Ideas Matter Dalton Thompson Rutherford Bohr ? De Broglie Einstein Plank Maxwell Newton Light

  4. The Photoelectric Effect Duality of Light • Wave behavior • Particle behavior 1905

  5. 1923 de Broglie’s Novel Notion Light was “known” (thought) to be a wave, but Einstein showed that it also acts particle-like Electrons were particles with known mass & charge What if …… electrons behaved as waves also

  6. Evidence for de Broglie’s Notion Diffraction pattern obtained with firing a beam of electrons through a crystal. This can only be explained if the electron behaves as a wave! Nobel Prize for de Broglie in 1929

  7. Electron Characteristics • Extremely small mass • Located outside the nucleus • Moving at very high speeds • Have specific energy levels • Standing wave behavior

  8. Baseball vs Electron A baseball behaves as a particle and follows a predictable path. BUT An electron behaves as a wave, and its path cannot be predicted. All we can do is to calculate the probability of the electron following a specific path.

  9. What if a baseball behaved like an electron? • Characteristic wavelength • baseball  10-34 m • electron  0.1 nm All we can predict is…..

  10. Werner Heisenberg(1901-1976) The Uncertainty Principle • Proposed that the dual nature of the electron places limitation on how precisely we can know both the exact location and speed of the electron • Instead, we can only describe electron behavior in terms of probability. speed position

  11. Erwin Schrodinger(1887-1961) Wave Equation & Wave Mechanics • In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron • When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level • Can be applied to more complex atoms too!

  12. Solutions to Schrodinger’s Wave Equation Gives the most probable location of electron in 3-D space around nucleus (probability map) - most probable location called an orbital - orbitals can hold a maximum of 2 e-

  13. “Most Successful Theory of 20th Century” Matter Dalton Thompson Rutherford Quantum Mechanics Bohr De Broglie Heisenberg Einstein Schrödinger Plank Maxwell Wave Mechanics Newton Light

  14. Quantum Mechanics ModelDescribes the arrangement and space occupied by electrons in atoms Electron’s energy is quantized Quantum Mechanics Mathematics of waves to define orbitals (wave mechanics)

  15. Bohr Model v. Quantum Mechanics BohrQ. Mech. Energy Electron Position/Path

  16. Dartboard Analogy Suppose the size of the probability distribution is defined as where there is a % chance of all hits being confined.

  17. Quantum Mechanics Model The electron's movement cannot be known precisely. We can only map the probability of finding the electron at various locations outside the nucleus. The probability map is called an orbital. The orbital is calculated to confine 90% of electron’s range.

  18. Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) = distance from nucleus 1, 2, 3, … SUBSHELLS (l) = shape of region of probability s, p, d, f ORBITALS (ml) = orientation in space

  19. Arrangement of Electrons in Atoms • There is a relationship between the quantum number (n) and its the number of subshells. Principal quantum number (n) = number of subshells

  20. Representing s Orbitals

  21. Comparison of 1s and 2s Orbitals The 2s orbital is similar to the 1s orbital, but larger in size. ”Larger” means that the highest probability for finding the electron lies farther out from the nucleus. Each can hold a maximum of electrons.

  22. Probability Maps of the Three 2p Orbitals The 2p orbital is in the n = energy level. There are 2p orbitals oriented in three directions. Each orbital can hold a maximum of electrons. The maximum number of electrons in the 2p sublevel is . Adding all 2p orbitals would result in a sphere.

  23. Probability Maps of the Five 3d Orbitals The five 3d orbitals are generally oriented in different directions. Adding all five orbitals, would result in a sphere. The five orbitals, taken together, make up the d subshell of the n = 3 shell. Each orbital can hold a maximum of two electrons. This sublevel has a maximum of electrons.

  24. Probability Maps of 7 f Orbitals

  25. Arrangement of Electrons in AtomsElectron Spin Quantum Number- ms Each orbital can be assigned no more than 2 electrons! And each electron spins in opposite directions.

  26. Electron Spin Quantum Number Diamagnetic: NOT attracted to a magnetic field Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

  27. Summary: 4 QUANTUM NUMBERS n ---> shell 1, 2, 3, 4, ... l ---> sublevel s, p, d, f ml ---> orbital -l ... 0 ... +l ms ---> electron spin +1/2 and -1/2

  28. Pauli Exclusion Principle- No two electrons in the same atom can have the same set of 4 quantum numbers. Determine the quantum numbers for the outer two valence electrons in the lithium atom.

  29. Aufbau Principle-Electrons fill open lower energy levels sequentially lower energy to higher energy

  30. spdf notation for H, atomic number = 1 1 no. of s 1 electrons value of l value of n Writing Electron Configurations Two ways of writing configs. One is called thespdf notation.

  31. Broad Periodic Table Classifications • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O) • Transition Elements: filling dorbitals (Fe, Co, Ni) • Lanthanide and Actinide Series(inner transition elements): filling 4fand 5forbitals (Eu, Am, Es)

  32. Writing Orbital Notations Two ways of writing configs. Other is called theorbital box notation. One electron has n = 1, l = 0, ml = 0, ms = + 1/2 Other electron has n = 1, l = 0, ml = 0, ms = - 1/2

  33. Energy ordering of orbitals for multi-electron atoms Different subshells within the same principal shell have different energies. The more complex the subshell, the higher its energy. This explains why the 3d subshell is higher in energy than the 4s subshell.

  34. Rules for Filling Orbitals Bottom-up (Aufbau’s principle) Fill orbitals singly before doubling up (Hund’s Rule) Paired electrons have opposite spin (Pauli exclusion principle)

  35. Cobalt Symbol Atomic Number Full Configuration Valence Configuration Shorthand Configuration

  36. Orbital diagram and electron configuration for a ground state lithium atom

  37. Orbital diagram and electron configuration for a ground state carbon atom Hund’s Rule- electrons in the same sublevel will spread out into their own orbital before doubling up.

  38. Silicon's valence electrons

  39. Selenium's valence electrons

  40. Core electrons and valence electrons in germanium

  41. Outer electron configuration for the elements

  42. The periodic table gives the electron configuration for As

  43. Valence Electrons by Group

  44. Ion charges by group

  45. Periodic Law All the elements in a group have the same electron configuration in their outermost shells Example: Group 2 Be 2, 2 Mg 2, 8, 2 Ca 2, 2, 8, 2

  46. Question Specify if each pair has chemical properties that are similar (1) or not similar (2): A. Cl and Br B. P and S C. O and S

  47. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends 1. Atomic and ionic size 2. Electron affinity 3. Ionization energy 4. Metallic Character

  48. Effective Nuclear Charge, Z* • Z* is the nuclear charge experienced by the outermost electrons. Screen 8.6. • Explains why E(2s) < E(2p) • Z* increases across a period owing to incomplete shielding by inner electrons. • Estimate Z* by --> [ Z - (no. inner electrons) ] • Z = number of electrons • Charge felt by 2s e- in Li Z* = 3 - 2 = 1 • Be Z* = 4 - 2 = 2 • B Z* = 5 - 2 = 3 and so on!

  49. Effective Nuclear Charge Figure 8.6 Electron cloud for 1s electrons

  50. Effective Nuclear Charge, Z* • Atom Z* Experienced by Electrons in Valence Orbitals • Li +1.28 • Be ------- • B +2.58 • C +3.22 • N +3.85 • O +4.49 • F +5.13 Increase in Z* across a period

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