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Chapter 6

Chapter 6. Chemical & Physical Properties of the Elements and the Periodic Table. Review Quiz Chapter 6. Heats of (kJ/mol) conversion. ∆H summation formula. Valence Electrons. The valence electrons are the electrons in the outer energy level ( valence level ). Alkali Metals.

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Chapter 6

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  1. Chapter 6 Chemical & Physical Properties of the Elements and the Periodic Table

  2. Review Quiz Chapter 6 • Heats of (kJ/mol) conversion. • ∆H summation formula.

  3. Valence Electrons • The valence electrons are the electrons in the outer energy level (valence level).

  4. Alkali Metals

  5. Alkaline Earth Metals

  6. Transition Elements (Metals)

  7. Halogens

  8. Noble Gases

  9. Inner Transition Elements (Metals)

  10. The Representative Elements

  11. Covalent radius • Covalent radius is essentially the size of an atom.

  12. Covalent Radii (atomic radii) Atomic Radius

  13. Ionic Radius • Ionic Radius is the size of an ion.

  14. Isoelectronic Series • Substances are isoelectronic if they have the same electron configuration. • Name two isoelectronic species.

  15. Ionization Energy • Ionization energy is the energy needed to remove an electron from an atom or ion.

  16. First Ionization Energy • First Ionization energy is the energy needed to remove the first electron from an atom.

  17. Multiple Ionization Energies • Second Ionization energy is the energy needed to remove the second electron from an atom. • Third Ionization energy is the energy needed to remove the third electron from an atom. • Etc.

  18. Ionization Energies in kJ/mol Write the equation representing the first ionization energy of hydrogen.

  19. First Ionization Energy of H • H + 1312 kJ → H+ + e-

  20. Electron Affinity • The energy change that occurs when an electron is added to an atom.

  21. Write the equation representing the electron affinity of hydrogen.

  22. Electron Affinity of H • H + e-→ H- + 72 kJ

  23. Effective Nuclear Charge (Zeff) • You will find many of the notes for effective nuclear charge on a sheet in your notebook titled “Effective Nuclear Charge”. • The effective nuclear charge (Zeff) of an atom is basically how well it is able to hold on to its most loosely held electron.

  24. Effective Nuclear Charge (Zeff) • We can estimate the effective nuclear charge of an atom by using the following: • The nuclear charge (Z) • The shielding effect • Electron repulsions

  25. The Nuclear Charge (Z) • Based on the number of protons in the nucleus. • Example: Carbon vs. Nitrogen

  26. The Nuclear Charge (Z)

  27. The greater the number of protons in the nucleus the greater the effective nuclear charge.

  28. Nuclear Charge and Zeff

  29. Shielding Effect. • The shielding effect is when electrons between the nucleus and the outermost electrons in an atom shield or lessen the hold of the nucleus on the outermost electrons.

  30. Shielding Effect.

  31. Shielding Effect. • Shielding can be checked by writing the electron configuration.

  32. Example of the Shielding Effect He atom in the excited state with one electron in the 1s and one electron in the 2p. He+ ion in the excited state with one electron in the 2p. 1s12p1 2p1

  33. Shielding EffectEnergy Levels vs. Sublevels • Energy levels have the greatest effect on shielding. • Sublevels increase shielding but to a far lesser extent.

  34. Ionization Energies in kJ/mol

  35. Zeff can help us explain the ionization energies.

  36. Explain the first ionization energies of Be and B A

  37. Explain the first ionization energies ofBe and Mg

  38. Effective Nuclear Charge can be used to help explain atomic radius. Atomic Radius

  39. Explain the difference in atomic radii for Li and Be. Which are 1.52 and 1.11 angstroms respectively.

  40. Explain the difference in atomic radii for Li and Na. Which are 1.52 and 1.86 angstroms respectively.

  41. Effective Nuclear Charge can be used to help explain atomic radius. • Based on nuclear charge and shielding.

  42. Electron Repulsions:Paired vs. Unpaired Electrons • A paired electron has increased electron – electron repulsion. • It is easier (takes less energy) to remove a paired electron than it does to remove an unpaired electron. • We check the pairing of electrons in the outer sublevel by writing an orbital filling diagram.

  43. Nitrogen vs. OxygenFirst Ionization Energy

  44. Nitrogen vs. OxygenFirst Ionization Energy

  45. It is much harder to remove an electron from helium than it is Li. This is Illustrated by their respective ionization energies given below. Explain. • He = 2370 kJ/mol • Li = 520 kJ/mol Stability Schmability

  46. Penetration Effect • Electrons in a higher energy level can often penetrate (dive) through lower energy levels because of the attraction that the nucleus has on them. • Smaller sublevels can penetrate closer to the nucleus than larger sublevels.

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