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Electron Structure

Electron Structure. By Diane Paskowski. Most Important Ideas of Quantum Mechanics. Orbital is not an orbit Electron NOT moving around the nucleus DO NOT KNOW how it is moving. What does it all mean?. Wave function has no easily visualized physical meaning

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Electron Structure

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  1. Electron Structure By Diane Paskowski

  2. Most Important Ideas of Quantum Mechanics • Orbital is not an orbit • Electron NOT moving around the nucleus • DO NOT KNOW how it is moving

  3. What does it all mean? • Wave function has no easily visualized physical meaning • Square of the function indicates the probability of finding an electron near a particular point in space • Probabilty distribution

  4. Difficult to define precisely. Orbital is a wave function. Picture an orbital as a three-dimensional electron density map. Hydrogen 1s orbital: Radius of the sphere that encloses 90% of the total electron probability. 7.5 Relative Orbital Size

  5. Probability Distribution for the 1s Wave Function

  6. Radial Probability Distribution

  7. Quantum Numbers • Three “variables” contained within the wave function – n, l, ml • First, the value of n determines the possible values of l • The values of l, then determine the possible values of ml • The magnetic spin, ms, is not determined by the wave function

  8. Principal quantum number (n) – size and energy of the orbital Angular momentum quantum number (l) – shape of atomic orbitals Magnetic quantum number (ml) – orientation of the orbital in space relative to the other orbitals in the atom 7.6 Quantum Numbers(wave function variables)

  9. Table 7.1 The Angular Momentum Quantum Numbers (l) and Corresponding Letters Used to Designate Atomic Orbitals

  10. Table 7.2 Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom

  11. Figure 7.13 Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals (a) The Electron Probability Distribution (b) The Surface Contains 90% of the Total Electron Probability (the Size of the Oribital, by Definition)

  12. Figure 7.18 Orbital Energy Levels for the Hydrogen Atom

  13. 1s Orbital 7.7

  14. 2px Orbital 7.7

  15. The Boundary Surface Representations of All Three 2p Orbitals 7.7

  16. Orbital 7.7

  17. 3dxy Orbital 7.7

  18. Orbital 7.7

  19. The Boundary Surfaces of All of the 3d Orbitals

  20. Representation of the 4f Orbitals in Terms of Their Boundary Surfaces 7.7

  21. Electron spin quantum number (ms) – can be +½ or -½. Pauli exclusion principle - in a given atom no two electrons can have the same set of four quantum numbers. An orbital can hold only two electrons, and they must have opposite spins. 7.8 Electron Spin

  22. Figure 7.19 A Picture of the Spinning Electron

  23. Orbital Energies 7.9

  24. AComparison of the Radial Probability Distributions of the 2s and 2p Orbitals 7.9

  25. The Radial Probability Distribution of the 3s Orbital 7.9

  26. A Comparison of the Radial Probability Distributions of the 3s,3p, and 3d Orbitals 7.9

  27. Electron correlation problem: Since the electron pathways are unknown, the electron repulsions cannot be calculated exactly When electrons are placed in a particular quantum level, they “prefer” the orbitals in the order s, p, d, and then f. 7.9 Polyelectronic Atoms

  28. Figure 7.22 The Orders of the Energies of the Orbitals in the First Three Levels of Polyelectronic Atoms

  29. As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen like orbitals. 7.11 Aufbau Principle

  30. The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals. 7.11 Hund’s Rule

  31. The Orbitals Being Filled for Elements in Various Parts of the Periodic Table 7.11

  32. Orbital Diagrams and Electron Configuration Worksheet Using the following diagram, complete the orbital diagram and electron configuration of the elements assigned by the instructor. Atom Orbital diagram E configuration __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d 4p

  33. The electrons in the outermost principal quantum level of an atom. 1s22s22p6 (valence electrons = 8) The elements in the same group on the periodic table have the same valence electron configuration. 7.11 Valence Electrons

  34. Determine the expected electron configurations for each of the following. a) S b) Ba c) Eu 7.11 Exercise

  35. History of the Periodic Table • Many early scientists organized the elements in a table. • Mendeleev • organized the elements according to physical and chemical properties. • Correctly predicted the existence of previously undiscovered elements and their properties

  36. Figure 7.24 Mendeleev's Early Periodic Table, Published in 1872

  37. Ionization Energy Electron Affinity Atomic Radius 7.12 Periodic Trends

  38. Energy required to remove an electron from a gaseous atom or ion. In general, as we go across a period from left to right, the first ionization energy increases. In general, as we go down a group from top to bottom, the first ionization energy decreases. 7.12 Ionization Energy

  39. Ionization Energy • Why does the IE decrease down a group? • Why does the IE increase across a period? • The effective nuclear charge increases across a period and decreases down a column. • What is effective nuclear charge?

  40. The Values of First Ionization Energy for the Elements in the First Six Periods 7.12

  41. Explain why the graph of ionization energy versus atomic number (across a row) is not linear. Where are the exceptions? 7.12 Concept Check

  42. Which atom would require more energy to remove an electron? Why? Li Cs 7.12 Concept Check

  43. Which has the larger second ionization energy? Why? Lithium or Beryllium 7.12 Concept Check

  44. Successive Ionization Energies (KJ per Mole) for the Elements in Period 3 7.12

  45. Energy change associated with the addition of an electron to a gaseous atom. In general as we go across a period from left to right, the electron affinities become more negative. In general electron affinity becomes more positive in going down a group. 7.12 Electron Affinity

  46. Figure 7.32 The Electron Affinity Values for Atoms Among the First 20 Elements that Form Stable, Isolated X- Ions

  47. In general as we go across a period from left to right, the atomic radius decreases. In general atomic radius increases in going down a group. 7.12 Atomic Radius

  48. Figure 7.33 The Radious of an Atom (r) is Defined as Half the Distance Between the Nuclei in a Molecule Consisting of Identical Atoms

  49. Which should be the larger atom? Why? Na Cl 7.12 Concept Check

  50. Which is larger? The hydrogen 1s orbital The lithium 1s orbital Which is lower in energy? The hydrogen 1s orbital The lithium 1s orbital 7.12 Concept Check

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