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