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Periodic Trends

Periodic Trends . Valence Electrons. Valence Electrons are electrons in the outermost energy level. - s or p electrons only (even when d and f electrons are present they are not in the outermost energy level ). Electron Dot Diagrams.

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Periodic Trends

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  1. Periodic Trends

  2. Valence Electrons • Valence Electrons are electrons in the outermost energy level. • - s or p electrons only (even when d and f electrons are present they are not in the outermost energy level)

  3. Electron Dot Diagrams • Electron Dot Diagrams show the valence electrons of an element. • Draw the electron dot diagrams for the following: • Mg N F

  4. How many valence electrons are in the halogens? • 1 • 2 • 7 • 8

  5. How many valence electrons are in Group 17? • 1 • 2 • 7 • 8

  6. How many valence electrons are in zinc and chromium? • 1 • 2 • 7 • 8 • Unable to tell

  7. Size of the Atom • Think of a ball of aluminum foil. What happens with each layer? • As you go down the periodic table, the energy levels increase and the size of the radius of the atom increases. (each energy level is like another layer of Al foil)

  8. Which element has a larger atomic radius: Na or K? • Na • K • Unable to be determined

  9. Which element has a larger atomic radius: Br or Cl? • Br • Cl • Unable to be determined

  10. Atomic Radius cont. • As you move from left to right in the periodic table, what happens to the number of protons in the nucleus? • What effect do these protons have on the electrons? • What effect do the electrons have on each other?

  11. Effective Nuclear Charge • Effectivenuclear charge (Zeff) is the net positive charge attracting the nucleus • Zeff= # of protons – # of electrons between nucleus and the outer electron Note: you are subtracting any electron at a lower energy level than the valence electrons

  12. Effective Nuclear Charge • Effectivenuclear charge (Zeff) is the net positive charge attracting the nucleus • Electron Shielding (or Screening) – These inner electrons shield the valence electrons from receiving the entire attractive nuclear charge because they repel the valence electrons.

  13. Effective Nuclear Charge • Zeff = # of protons – # of electrons between nucleus and the outer electron • What is the effective nuclear charge of Magnesium?

  14. Effective Nuclear Charge • Zeff = # of protons – # of electrons between nucleus and the outer electron • What is the effective nuclear charge of Sulfur?

  15. What is the effective nuclear charge of neon? • 1 • 3 • 5 • 6 • 7 • 8

  16. What is the effective nuclear charge of nitrogen? • 1 • 3 • 5 • 6 • 7 • 8

  17. Effective Nuclear Charge and Atomic Radius • Within a period, as you go from left to right, Zeffincreases, and attracts the electrons more strongly. • As the electrons are more attracted to the nucleus, the atomic radius decreases. • Summary: as you go from left to right, the atomic radius generally decreases.

  18. Which element has a larger atomic radius: Li or Be? • Li • Be • Unable to be determined

  19. Which element has a larger atomic radius: Si or Ar? • Si • Ar • Unable to be determined

  20. Which element has a larger atomic radius: Be or Mg? • Be • Mg • Unable to be determined

  21. Which element has a larger atomic radius: Si or C? • C • Si • Unable to be determined

  22. Octet Rule • The octet rule states that all atoms attempt to become stable by having a full valence electron shell (generally 8 electrons, hence octet rule). • Atoms will gain, lose, or share electrons in order to attain this stability.

  23. Which group already has a full octet? • Alkali metals • Transition metals • Halogens • Noble Gases

  24. Ionization Energy • Electrons are held in atoms by their attraction to the positively charged nucleus. • To remove an electron requires energy. • Ionization energy is the energy required to remove the least tightly bound (or outermost) electron from an atom.

  25. Ionization Energy cont. • Compare Li and K. • How many valence electrons? • What is the relative size of the atoms? • Which has a higher ionization energy?

  26. Ionization Energy • As you go down a group, the ionization energy decreases because it takes less energy to remove an electron. • The least tightly bound electrons are further from the positive nucleus, and can therefore be removed more easily.

  27. Which group has a higher ionization energy? • Halogens • Alkali Metals • Alkaline Earth Metals • All the same

  28. Which element has a higher ionization energy? • Li • Be • F • Ne

  29. Which element has the lowest ionization energy? • Na • Mg • S • Ar

  30. Summary • Ionization energy generally decreases as you go down a group and from right to left in a period.

  31. A Quick Review …

  32. How many valence electrons does oxygen have? • 1 • 3 • 4 • 6 • 8

  33. Which has the biggest atomic radius? • K • Na • Li • Be • Cs • Ca • F

  34. What is the effective nuclear charge on lithium? • 1 • 2 • 3 • None of the above

  35. Which has the greatest ionization energy? • K • Na • Li • Be • Cs • Ca • F

  36. How many valence electrons are in group 14? • 1 • 4 • 7 • 14

  37. Second, Third, etc. Ionization Energy • Second ionization is the energy required to remove a second electron. • Ex. Sodium has a lower (first) ionization energy than Magnesium but Mg has a lower second ionization energy than Na. • Why?

  38. Which has a lower third ionization energy? • K • Ca • Ga • C

  39. Electronegativity • Electronegativity is the tendency of an element to attract electrons in a bond. • Therefore, elements that want to gain electrons will have higher electronegativity.

  40. Which has the greatest electronegativity? • K • Na • Li • Be • Cs • Ca • F

  41. Which group has the greatest electronegativity? • Alkali metals • Alkaline earth metals • Halogens • Noble Gases

  42. Trend Summaries The pattern for increasing electronegativity (except for noble gases). The pattern for increasing ionization energy. The pattern for increasing atomic radius.

  43. Ions • When an atom loses or gains electrons, it gains a charge. • An ion is a charged ion. • A positive ion is called a cation. • A negative ion is called an anion.

  44. Size of Ions • Cations (positive) – lose valence electrons in the outermost energy level. • They lose an energy level so they get smaller. • Anions (negative) – gain valence electrons but their Zeffdoes not change. • They get bigger because of shielding and a lower Zeff

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