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Aufbau Principle

Aufbau Principle

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Aufbau Principle

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  1. Aufbau Principle e- fill lowest energy levels 1st. Half-filled and filled states are preferred. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 6f 6g 6h 7s 7p 7d 7f 7g 7h 7i 4f 5p 4d 5s 4p 3d Increasing energy 4s 3p 3s 3 Li Lithium 1s22s1 1 H Hydrogen 1s1 11 Na Sodium [Ne]3s1 8 O Oxygen 1s22s22p4 27 Co Cobalt [Ar]3d74s2 34 Se Selenium [Ar]3d104s24p4 5 B Boron 1s22s22p1 29 Cu Copper [Ar]3d104s1 13 Al Aluminum [Ne]3s23p1 24 Cr Chromium [Ar]3d54s1 22 Ti Titanium [Ar]3d24s2 19 K Potassium [Ar]4s1 16 S Sulfur [Ne]3s23p4 31 Ga Gallium [Ar]3d104s24p1 2 He Hydrogen 1s2 12 Mg Magnesium [Ne]3s2 6 C Carbon 1s22s22p2 14 Si Silicon [Ne]3s23p2 32 Ge Germanium [Ar]3d104s24p2 35 Br Bromine [Ar]3d104s24p5 30 Zn Zinc [Ar]3d104s2 4 Be Beryllium 1s22s2 20 Ca Calcium [Ar]4s2 28 Ni Nickel [Ar]3d84s2 9 F Fluorine 1s22s22p5 17 Cl Chlorine [Ne]3s23p5 10 Ne Neon 1s22s22p6 36 Kr Krypton [Ar]3d104s24p6 15 P Phosphorus [Ne]3s23p3 18 Ar Argon [Ne]3s23p6 21 Sc Scandium [Ar]3d14s2 7 N Nitrogen 1s22s22p3 26 Fe Iron [Ar]3d64s2 33 As Arsenic [Ar]3d104s24p3 29 Cu Copper [Ar]3d104s1 24 Cr Chromium [Ar]3d54s1 23 V Vanadium [Ar]3d34s2 25 Mn Manganese [Ar]3d54s2 2p 2s 1s nucleus

  2. Irregular e- configurations of Cr and Cu • Half filled and filled sublevels are preferred. Chromium: A 4s electron moves to a 3d sublevel to half fill its entire 3d sublevel Copper: A 4s electron moves to FILL its 3d sublevel

  3. Orbital filling table s1 s2p1 s2p2 s2p3 s2p4 s2p5 s2 s2 s2p6

  4. Mendeleev’s Periodic Table (1872) Dmitri Mendeleev 1st to publish an organized table of elements. • Grouped elements by similar chemical properties. • Arranged elements by increasing mass.

  5. Modern Russian Table

  6. Henry Moseley (1913) Used X-Ray diffraction to determine how many protons are in an atom of an element. • Grouped elements by similar chemical properties. • Arranged elements by increasing atomic number.

  7. Period The Periodic Table Group or Family 1 2 3 4 5 6 7

  8. Nonmetals Metals Properties of Metalloids Metalloids straddle the border between metals and nonmetals on the periodic table. Metalloids

  9. Periodic Table with Group Names Period Group or Family 1 2 3 Noble Gases Halogens 4 Transition Metals Alkali Metals Alkaline Earth Metals Boron Family Carbon Family Nitrogen Family Oxygen Family 5 6 7 Inner Transition Metals

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

  11. Determination of Atomic Radius: Half of the distance between nuclei in covalently bonded diatomic molecule (e- shared) "covalent atomic radii"

  12. Trends in Atomic Size • Influenced by three factors: 1. Energy Level • Higher energy level is further away. 2. Charge on nucleus (Zeff = effective nuclear charge) • More charge pulls electrons in closer. 3. Shielding effect • Electrons within level and from previous levels block the effects of the (+) nucleus. e <-> e repulsion

  13. Periodic TableTrend forAtomic Radii Size of atom decreases from left to right. Size of atom increases from top to bottom.

  14. Periodic Trends in Atomic Radius • Radius decreases across a period Effective nuclear charge, Zeff, is increased. Due to: more overall charge (p+ and e-)  shielding from lower levels is constant.

  15. Group trends H Li • As we go down a group... • each atom has another energy level (more shielding) • so the atoms get bigger (Zeff is less effective). Na K Rb

  16. Periodic Trends • As you go across a period, the radius gets smaller. (more effective nuclear charge, Zeff) • Electrons are in same energy level. • More nuclear charge. atomic #:(p+ e-) • Outermost electrons are closer. z = 11 17 18 12 13 14 15 16 Na Mg Al Si P S Cl Ar

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

  18. 1stIonization Energy - the energy required to remove the 1st electron from a mole of atoms. 2372 kJ/mol He H 1312 kJ/mol 2081 kJ/mol Li Be Ne O N F 520 kJ/mol B C 1521 kJ/mol Ar Na 496 kJ/mol 1351 kJ/mol K Kr 419 kJ/mol Increasing 1st Ionization energy 1170 kJ/mol Rb 403 kJ/mol Xe 1037 kJ/mol Cs Rn 376 kJ/mol Fr 375 kJ/mol

  19. Ionization Energy - the energy required to remove an electron from an atom • Increases for successive electrons taken from • the same atom • Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove • Tends to decrease down a group Outer electrons are farther from the nucleus

  20. Table of 1st Ionization Energies

  21. Ionization of Magnesium Mg + 738 kJ  Mg+ + e- 1st ionization energy Mg+ + 1451 kJ  Mg2+ + e- 2nd ionization energy Mg2+ + 7733 kJ  Mg3+ + e- 3rd ionization energy

  22. Symbol First Second Third 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne 11810 14840 3569 4619 4577 5301 6045 6276

  23. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

  24. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across • a period • Electronegativities tend to decrease down a • group or remain the same

  25. Periodic Table of Electronegativities Increasing electronegativity

  26. Summation of Periodic Trends

  27. Ionic Radii • Positively charged ions formed when • an atom of a metal loses one or • more electrons Cations • Smaller than the corresponding • atom • Negatively charged ions formed • when nonmetallic atoms gain one • or more electrons Anions • Larger than the corresponding • atom

  28. Ion radius Atom radius

  29. Shielding • The electron in the outermost energy level experiences more inter-electron repulsion (shielding). • Second electron has same shielding, if it is in the same period

  30. Group trends • As you go down a group, first IE decreases because... • The electron is further away. • More shielding.

  31. Periodic trends • All the atoms in the same period have the same energy level. • Same shielding. • But, increasing nuclear charge • So IE generally increases from left to right. • Exceptions at full and 1/2 full orbitals.

  32. He • He has a greater IE than H. • same shielding • greater nuclear charge H First Ionization energy electrons 1s 2s 2p nucleus Orbital diagram Atomic number

  33. He • Li has lower IE than H • Outer electron further away • outweighs greater nuclear charge H First Ionization energy electrons Li nucleus 1s 2s 2p Orbital diagram Atomic number

  34. He • Be has higher IE than Li • same shielding • greater nuclear charge H First Ionization energy Be electrons Li nucleus 1s 2s 2p Orbital diagram Atomic number

  35. He • B has lower IE than Be • same shielding • greater nuclear charge • p orbital is slightly more diffuse and its electron easier to remove H First Ionization energy Be electrons B Li nucleus 1s 2s 2p Atomic number

  36. He C H First Ionization energy Be electrons B Li nucleus 1s 2s 2p Atomic number

  37. He N C H First Ionization energy Be B Li nucleus 1s 2s 2p Atomic number

  38. He N • Breaks the pattern, because the outer electron is paired in a p orbital and experiences inter-electron repulsion. O C H First Ionization energy Be B Li nucleus 1s 2s 2p Atomic number

  39. He F N O C H First Ionization energy Be B Li nucleus 1s 2s 2p Atomic number

  40. 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 nucleus 1s 2s 2p Atomic number

  41. 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

  42. First Ionization energy Atomic number

  43. Table of 1st Ionization Energies