The document discusses electron configuration and trends in ionization energy. It explains that ionization energy decreases down a group and increases across a period due to changes in effective nuclear charge. Electrons fill atomic orbitals according to Aufbau principle and Hund's rule. The document provides examples of writing electron configurations and condensed configurations for various elements. Successive ionization energy data supports the electron configuration model.

1. IB Chemistry Power PointsTopic 12Atomic Structurewww.pedagogics.caElectronConfiguration

2. HL Topic 12.1 – Electron ConfigurationIonization 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.

6. Across a Period: shielding remains constant

7. atomic number increases so effective nuclear charge increases

8. ionization energy increasesDown a Group: 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 energyH+e-e-Li+e-e-2e-++++++++e-Na+++++++++2e-2e-8e-8e-e-Hydrogen (Z=1)Lithium (Z=3)2e-Sodium (Z=11)

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 subdividedElectron 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,f1st Term: Shell (n)- principle energy leveln = 3n = 3n = 2n = 2lone electronof Hydrogenn = 11sThe 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 = 3n = 21sThe first shell (1) has one subshell (s). The ssubshell has 1 spherical shaped orbitalorbitals are volumes of space where the probability of finding an electron is high

17. The Electronic Configuration of HydrogenenergyHydrogen 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.1sElectronic configuration1s1shell# of electrons presentsubshell1s Orbital Energy Level Diagram

18. The Electronic Configuration of Helium He: Atomic # of 2, 2 electrons in a neutral He atomH 1s1He 1s2He 1s 1sthe maximum number of electrons in an orbital is TWOif 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 1s2s Li 1s2p2s1sOrbital Energy Level DiagramLi 1s22s1Electronic configuration

20. 2s Be 1s Be 1s22s2Berylium (Be)Be: Z=4 Be has 4 electrons.Electronic configurationOrbital Energy Level DiagramBoron (B) has 5 electrons, the s subshell is full so the 5th electron occupies the first orbital in the p subshell2p 2sB 1s2p2s2ndshell1sB 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 < ps subshell: spherical1 orbitalzxyxzyp 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 isBut always written asC 1s22s22px1py12p2s2ndshell1sC 1s22s22p22p  2sC 1s

23. Can we relate the filling of the subshells with the ionization energy data?2p   2sN 1s1s22s22p32p   2sO 1s1s22s22p42p   2sNe 1s1s22s22p6

24. Ionization energy trendsDown 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 modelBe to B “dip”- because s shields p and lowers ENCN to O “dip” - because repulsions between electron pair in first full orbital (experimental evidence supporting Aufbau and Hund)

26. Electron Configurations and the Periodic TableSo 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 configurationsExample CALCIUM  1s22s22p63s23p64s2Principle energy level subshell # of e-Calcium can also be written shorthand as:[Ar]4s2

27. PracticeUse the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements Fluorine 56FeMagnesium - 22131IPotassium – 4275GeZirconium – 9041Ca2+

28. PracticeUse the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements Fluorine 1s22s2p556Fe 1s22s2p63s23p64s23d6Magnesium – 22 1s22s2p63s2131I 1s22s2p63s23p63d104s24p64d105s25p5Potassium – 42 1s22s2p63s23p64s175Ge 1s22s2p63s23p64s23d104p2Zirconium – 90 1s22s2p63s23p64s23d104p65s24d241Ca2+ 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] 4s23d104p5Read 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 IONSWhen 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 4s2The 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] 4s23d104p5Si ___________________________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