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- 1. IB Chemistry Power Points<br />Topic 12<br />Atomic Structure<br />www.pedagogics.ca<br />Electron<br />Configuration<br />
- 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. Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.<br />
- 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. 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. Across a Period:<br /><ul><li> shielding remains constant
- 7. atomic number increases so effective nuclear charge increases
- 8. ionization energy increases</li></li></ul><li>Down a Group:<br /><ul><li> shielding increases AND atomic number increases
- 9. effective nuclear charge does not change significantly
- 10. valence electrons further from nucleus
- 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. 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. The organization of the Periodic table correlates directly to electron structure<br />
- 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. 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. 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. 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. 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. 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. 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 />

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