2011 hl ib chemistry - topic 12


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2011 hl ib chemistry - topic 12

  1. 1. IB Chemistry Power Points<br />Topic 12<br />Atomic Structure<br />www.pedagogics.ca<br />Electron<br />Configuration<br />
  2. 2. HL Topic 12.1 – Electron Configuration<br />Ionization energy is defined as the amount of energy required to remove a mole of electrons from a mole of gaseous atoms of a particular element.<br />E(g)  E+(g) + e-<br />
  3. 3. Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.<br />
  4. 4. Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.<br />WHY?<br />
  5. 5. Effective nuclear charge is the net positive charge felt by an electron in an atom. <br />The basics of electrostatics imply that each and every electron is attracted to each and every proton in the nucleus and repelled by every other electron. However . . . <br />Electrons between the valence electron and the nucleus provide a shielding effect – weakening the electrostatic force on the valence electron. <br />
  6. 6. Across a Period:<br /><ul><li> shielding remains constant
  7. 7. atomic number increases so effective nuclear charge increases
  8. 8. ionization energy increases</li></li></ul><li>Down a Group:<br /><ul><li> shielding increases AND atomic number increases
  9. 9. effective nuclear charge does not change significantly
  10. 10. valence electrons further from nucleus
  11. 11. so weaker electrostatic force and lower ionization energy</li></li></ul><li>H+<br />e-<br />e-<br />Li+<br />e-<br />e-<br />2e-<br />+<br />+<br />+<br />+<br />+<br />+<br />+<br />+<br />e-<br />Na+<br />+<br />+<br />+<br />+<br />+<br />+<br />+<br />+<br />2e-<br />2e-<br />8e-<br />8e-<br />e-<br />Hydrogen (Z=1)<br />Lithium (Z=3)<br />2e-<br />Sodium (Z=11)<br />
  12. 12. This diagram shows how 1st ionization energy decreases down a group and that trends in ionization energy also occur across a period. Look at the H, Li, Na, K, Rb, Cs values!<br />.<br />
  13. 13. Looking at just the trend across the 1st period, what does the graph imply?<br />The theory is . . . <br />Across a period, number of p+ increases so effective nuclear charge increases. <br />As a result, the valence electrons are more strongly held, and are closer to the nucleus (radius decreases)<br />This does not explain the drop in ionization energy (decreased stability) observed between Be and B and between N and O.<br />
  14. 14. NEW IDEA – suborbitals (or subshells)<br />Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels<br />
  15. 15. Old Idea expanded: 2,8,8,2 configuration with energy shells 1,2,3,4 . . . but each shell subdivided<br />Electron arrangement in atoms can be described by terms called quantum numbers – no two electrons can have the same quantum number!<br />2nd Term: subshell<br /> - designated by s, p,d,f<br />1st Term: Shell (n)<br />- principle energy level<br />n = 3<br />n = 3<br />n = 2<br />n = 2<br />lone electron<br />of Hydrogen<br />n = 1<br />1s<br />The first energy shell (1) has one subshell (s).<br />
  16. 16. 2nd Term: subshell<br /> - designated by s, p, d, f<br /> - designates the sub-energy level<br /> within the shell.<br />- refers to the shape(s) of the volume of space in which the electron can be located.<br />n = 3<br />n = 2<br />1s<br />The first shell (1) has one subshell (s). <br />The ssubshell has 1 spherical shaped orbital<br />orbitals are volumes of space where the probability of finding an electron is high<br />
  17. 17. The Electronic Configuration of Hydrogen<br />energy<br />Hydrogen has one electron located in the first shell (1). (Aufbau principle)<br />The first shell has only one subshell (s). The ssubshell contains 1 spherical orbital.<br />1s<br />Electronic configuration<br />1s1<br />shell<br /># of electrons present<br />subshell<br />1s <br />Orbital Energy Level Diagram<br />
  18. 18. The Electronic Configuration of Helium <br /> He: Atomic # of 2, 2 electrons in a neutral He atom<br />H 1s1<br />He 1s2<br />He 1s <br />1s<br />the maximum number of electrons in an orbital is TWO<br />if there are 2 electrons in the same orbital they must have an opposite spin. <br />This is called Pauli’s Exclusion Principle<br />
  19. 19. Lithium (Li)<br />Li: Z=3 Li has 3 electrons.<br />The 2nd shell (n= 2) has 2 subshells which are s and p. <br />The s subshell fills first! (Aufbau Principle)<br />2ndshell <br />1s<br />2s <br />Li 1s<br />2p<br />2s<br />1s<br />Orbital Energy Level Diagram<br />Li 1s22s1<br />Electronic configuration<br />
  20. 20. 2s <br />Be 1s<br /> <br />Be 1s22s2<br />Berylium (Be)<br />Be: Z=4 Be has 4 electrons.<br />Electronic configuration<br />Orbital Energy Level Diagram<br />Boron (B) has 5 electrons, the s subshell is full so the 5th electron occupies the first orbital in the p subshell<br />2p <br />2s<br />B 1s<br />2p<br />2s<br />2ndshell<br />1s<br />B 1s22s22p1<br />
  21. 21. Subshells so far<br /> - designated by s, and p<br /> - refers to the shape(s) of<br /> the volume in which the electron<br /> can be located.<br /> - also designates an energy level<br /> within the shell.<br /> - relative energy: s < p<br />s subshell: spherical<br />1 orbital<br />z<br />x<br />y<br />x<br />z<br />y<br />p subshell: pair of lobes, 3 orbitals, each holds 2 electrons<br />
  22. 22. Carbon (C)<br />C: Z=6 C has 6 electrons.<br />The 6th electron occupies an empty p orbital. This illustrates “Hund’s Rule” – electrons do not pair in orbitals until each orbital is occupied with a single electron.<br />The electron configuration is<br />But always written as<br />C 1s22s22px1py1<br />2p<br />2s<br />2ndshell<br />1s<br />C 1s22s22p2<br />2p  <br />2s<br />C 1s<br />
  23. 23. Can we relate the filling of the subshells with the ionization energy data?<br />2p   <br />2s<br />N 1s<br />1s22s22p3<br />2p   <br />2s<br />O 1s<br />1s22s22p4<br />2p   <br />2s<br />Ne 1s<br />1s22s22p6<br />
  24. 24. Ionization energy trends<br />Down a group : ionization energy decreases<br />- ENC constant but atoms larger so easier to ionize <br />Across a period : ionization energy increases<br />- increasing ENC therefore smaller size (e- closer to nucleus)<br />so harder to ionize<br />
  25. 25. Explaining the “dips” – support for s and p orbital model<br />Be to B “dip”<br />- because s shields p and lowers ENC<br />N to O “dip” <br />- because repulsions between electron pair in first full orbital (experimental evidence supporting Aufbau and Hund)<br />
  26. 26. Electron Configurations and the Periodic Table<br />So far, we have seen how the subshell model provides and explanation for the patterns in ionization energy we see in the periodic table.<br />You have also seen how to write electron configurations<br />Example CALCIUM  1s22s22p63s23p64s2<br />Principle energy level subshell # of e-<br />Calcium can also be written shorthand as:<br />[Ar]4s2<br />
  27. 27. Practice<br />Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements<br /> <br />Fluorine <br />56Fe<br />Magnesium - 22<br />131I<br />Potassium – 42<br />75Ge<br />Zirconium – 90<br />41Ca2+<br />
  28. 28. Practice<br />Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements<br /> <br />Fluorine 1s22s2p5<br />56Fe 1s22s2p63s23p64s23d6<br />Magnesium – 22 1s22s2p63s2<br />131I 1s22s2p63s23p63d104s24p64d105s25p5<br />Potassium – 42 1s22s2p63s23p64s1<br />75Ge 1s22s2p63s23p64s23d104p2<br />Zirconium – 90 1s22s2p63s23p64s23d104p65s24d2<br />41Ca2+ 1s22s2p63s23p6<br />
  29. 29. The organization of the Periodic table correlates directly to electron structure<br />
  30. 30. Condensed electron configurations – for example the electron configuration of bromine can be written [Ar] 4s23d104p5<br />Read questions carefully – many IB questions require you to write the FULL electron configuration<br />
  31. 31. Electron configuration of ions:<br />In general, electrons will be removed from orbitals (ionization) in the reverse order that the orbitals were filled. In other words, electrons vacate higher energy orbitals first.<br />The exception: TRANSITION METAL IONS<br />When these ions form, electrons are removed from the valence shell sorbitals before they are removed from valence dorbitals when transition metals are ionized.<br />For example: Cobalt has the configuration [Ar] 4s23d7 OR [Ar] 3d7 4s2<br />The Co2+ and Co3+ ions have the following electron configurations. <br /> Co2+: [Ar] 3d7Co3+: [Ar] 3d6<br />
  32. 32. Condensed electron configurations – for example the electron configuration of bromine can be written [Ar] 4s23d104p5<br />Si ___________________________<br />S2- ___________________________<br />Rb+ ___________________________<br />Se ___________________________<br />Ar ___________________________<br />Nb ___________________________<br />Zn2+ ___________________________<br />Cd ___________________________<br />Sb ___________________________<br />
  33. 33. You are responsible for configurations up to Z = 54 (Xe). The table works well for this with the exception of Cr and Cu<br />
  34. 34. Chromium’s configuration is:<br />[Ar]4s13d5<br />Copper’s configuration is:<br />[Ar]4s13d10<br />These configurations are energetically more stable than the expected arrangements. KNOW THEM!<br />
  35. 35. Successive ionization<br />energy data supports the electron configuration model<br />189367.7<br />169988<br />35458 31653 25661 21711 18020 13630 10542.5<br />3rd<br />7732.7<br />2nd<br />1450.7<br />1st<br />737.7<br />
  36. 36. Review: the principles involved<br />Aufbau Principle: electrons will fill the lowest energy orbitals first<br />Hund’s Rule: the most stable arrangement of electrons in orbitals of equal energy is where there is the maximum number of unpaired electrons all with the same spin.<br />Pauli’s Exclusion Principle: A maximum of two electrons can occupy a single orbital. These electrons will have opposite spins.<br />