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Periodic Trends

Periodic Trends. OBJECTIVES: Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities. Trends in Atomic Size. First problem: Where do you start measuring from? The electron cloud doesn’t have a definite edge.

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Periodic Trends

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  1. Periodic Trends • OBJECTIVES: • Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities.

  2. Trends in Atomic Size • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. • We get around this by measuring more than 1 atom at a time.

  3. Atomic Size • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Diatomic means “two atoms” or a two atom molecule } Radius

  4. Trends in Atomic Size • Influenced by three factors: 1. Energy Level • Higher energy level is further away. 2. Charge on Nucleus • More charge pulls electrons in closer. • 3. Shielding Effect Electrons between nucleus and valance electrons

  5. Group Trends H • As we go down a group... • each atom adds another energy level, • so the atoms get bigger. Li Na K Rb

  6. Periodic Trends • As you go across a period, the radius gets smaller. • Electrons are in same energy level. • More nuclear charge. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

  7. Atomic Radius Trends

  8. Rb Overall K Na Li Atomic Radius (nm) Kr Ar Ne H Atomic Number 10

  9. Periodic Trend in Atomic Radius

  10. Ion Formation An atom can easily lose or gain electrons. The resulting ion is an atom that has an imbalance of charge or carries either a positive charge (a loss of electrons) or a negative charge (a gain of electrons). A positive ion is called a cation. A negative ion is called an anion.

  11. Trends in Ionization Energy • The amount of energy required to completely remove an electron from a gaseous atom. • Removing one electron makes a 1+ ion. • The energy required to remove the first electron is called the first ionization energy.

  12. Ionization Energy • The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.

  13. Ionization Energy How many valance electrons does Magnesium have? How many is it willing to give up?

  14. First Ionization Energy

  15. Second Ionization Energy

  16. Third ionization Energy Mg3+ 1s22s22p5

  17. What Determines IE • The greater the nuclear charge, the greater IE. • Greater distance from nucleus decreases IE • Filled and half-filled sublevel have lower energy, so achieving them is easier, lower IE. • Shielding effect on valence electrons

  18. Group Trends • As you go down a group, first IE decreases because... • the valence electron is further away • and there is more shielding.

  19. Periodic Trends • As you go across a period, • All the atoms in the same period have the same highest energy level. • They have similar shielding. • But, increasing nuclear charge reduces atomic radius (bringing valence closer) • so IE generally increases from left to right. • Exceptions are at full and 1/2 full sublevels.

  20. He • He has a greater IE than H. • same shielding • greater nuclear charge H First Ionization energy Atomic number

  21. He • Li has lower IE than H • more shielding • further away • outweighs greater nuclear charge H First Ionization energy Li Atomic number

  22. He • Be has higher IE than Li • same shielding • greater nuclear charge H First Ionization energy Be Li Atomic number

  23. He • B has lower IE than Be • same shielding • greater nuclear charge • By removing an electron we leave the s orbital full H First Ionization energy Be B Li Atomic number

  24. He C H First Ionization energy Be B Li Atomic number

  25. He N C H First Ionization energy Be B Li Atomic number

  26. He N • Breaks the pattern, because removing an electron leaves 1/2 filled p orbital O C H First Ionization energy Be B Li Atomic number

  27. He F N O C H First Ionization energy Be B Li Atomic number

  28. Ne He F N • Ne has a lower IE than He • Both are full, • Ne has more shielding • Greater distance O C H First Ionization energy Be B Li Atomic number

  29. Ne He • Na has a lower IE than Li • Both are s1 • Na has more shielding • Greater distance F N O C H First Ionization energy Be B Li Na Atomic number

  30. First Ionization energy Atomic number

  31. Driving Force • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full energy level. • Atoms behave in ways to achieve noble gas configuration.

  32. Trends in Ionic Size • Cations form by losing electrons. • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration. Na Na+1

  33. Ionic Size • Anions form by gaining electrons. • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration. Cl Cl-1

  34. Ion Group Trends • Adding energy level • Ions get bigger as you go down. Li1+ Na1+ K1+ Rb1+ Cs1+

  35. Ion Periodic Trends • Across the period, nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ C4+ Be2+

  36. Ionic Radius Trends

  37. Size of Isoelectronic Ions • Iso- means the same • Iso electronic ions have the same # of electrons • Al3+ Mg2+ Na1+ Ne F1- O2- and N3- • all have 10 electrons • all have the configuration: 1s22s22p6

  38. Size of Isoelectronic Ions • Positive ions that have more protons would be smaller. N3- O2- F1- Ne Na1+ Al3+ Mg2+

  39. Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to itself. Concept proposed by Linus Pauling 1901-1994

  40. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair is the sharing? • Large electronegativity means the atom pulls the electrons toward it. • Atoms with large electron affinity have larger electronegativity.

  41. H 2.2 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.1 O 3.5 F 4.1 Na 1.0 Mg 1.2 Al 1.5 Si 1.7 P 2.1 S 2.4 Cl 2.8 K 0.9 Ca 1.0 Ga 1.8 Ge 2.0 As 2.2 Se 2.5 Br 2.7 Rb 0.9 Sr 1.0 In 1.5 Sn 1.7 Sb 1.8 Te 2.0 I 2.2 Cs 0.9 Ba 1.0 Tl 1.4 Pb 1.5 Bi 1.7 Po 1.8 At 1.9 Electronegativity • Relative ability of atoms to attract electrons of bond.

  42. Electronegativity

  43. Ionization Energy, Electronegativity, and Electron Affinity INCREASE

  44. Atomic Size Increases Ionic size increases

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