10.4 Energy Levels of Electrons

1. 10.4 Energy Levels of Electrons Electrons move in definite energy levels; these are labeled 1 - 7

2. Sublevels in each of the 7 main energy levels • Each level has sublevel(s) which are probability shapes that show where the electrons may be at any one time. Also known as orbitals.

3. S orbital can hold up to 2 electrons

4. P orbital can hold up to 6 electrons

5. D orbital can hold up to 10 electrons

6. F orbital can hold up to 14 electrons

7. How do they fill? • Aufbau chart shows how electrons fill into the main energy levels and the sublevels or orbitals

8. Energy Levels and Sublevels • 1s • 2s 2p • 3s 3p 3d • 4s 4p 4d 4f • 5s 5p 5d 5f • 6s 6p 6d 6f • 7s 7p 7d 7f

9. Aufbau Diagram or Chart • 1s START HERE and follow • 2s 2p the arrows! • 3s 3p 3d • 4s 4p 4d 4f • 5s 5p 5d 5f • 6s 6p 6d 6f • 7s 7p 7d 7f

10. Electron Configuration • 4 Be • 1s2 2s2 • 15 P • 1s22s22p63s23p3 • 25 Mn • 1s22s22p63s23p64s23d5

11. Now begin working HW Handout #1

12. Q. O. D.—Show how you solved it! _______1. The element with atomic number 53 contains a) 53 neutrons b) 53 protons C) 26 neutrons & 27 protons d) 26 protons & 27 neutrons 108 _______2. The number of neutrons in an atom of 47 Ag is a) 47 b) 108 c) 155 d) 61 27 _______3. The number of electrons in an ion of 13 Al3+ is a) 13 b) 10 c) 27 d) 14 _______4. What is the relative atomic mass of boron if two stable isotopes of boron have the following mass and abundance: 10.0129 amu (19.91%) & 11.0129 (80.09%) a) 10.81 amu b) 10.21 amu c) 10.62 amu d) 10.51 amu

13. Hund’s Rule • Hund’s rule states that electrons will fill a subshell unpaired until it cannot occupy a subshell unpaired and has to pair. (Applies only to p, d, and f)

14. Pauli Exclusion Principle • According to the Pauli Exclusion Principle, states that when two electrons must occupy the same subshell they will have OPPOSITE spin. One will be symbolized by an UP arrow and the other by a DOWN arrow. Such as this example.

15. Pauli Exclusion Principle • So they invented spin (+1/2 or -1/2) called spin up and spin down. Has nothing to do with the direction of the electron--we don’ t know how they move just where they may be at with 90% chance of finding it inside the energy level and orbital designated.

16. Orbital Diagrams S orbitals get one box P orbitals get 3 boxes (2 e- per box) D orbitals get 5 boxes and f gets 7

17. Orbital Diagrams (cont.) • Insert electrons (using arrows into each box according to Hund’s and Pauli) 2 p3

18. Answer • 2p3 (arrows can all point up or down) • Now try 4f10

19. Answer to 4f10 • Arrows may point up or down if they are in boxes individually; however, if there are 2 electrons in a box, one must point up and one down.

20. Now begin working on the lab activity Electron Configurations

21. #5.

22. #9.

23. Electrons and the Periodic Table History of the Table Periodic Law Important People

24. 1. Mendeleev • Mendeleev was a Russian chemist who arranged the known elements in vertical columns in order of increasing mass and noticed a pattern in physical and chemical properties

25. 2. Mosley • Mosley was a British physicist who determined the atomic number (number of protons) of the atoms of elements and then arranged the elements according to their atomic number.

26. 3. Arrangement of Periodic Table • The current periodic table is arranged in order of increasing atomic number.

27. 11. Periodic Law • According to Mendeleev, “"The properties of the elements are a periodic function of their atomic masses" • According to Mosley, the periodic table was arranged according to atomic number and patterns repeat periodically

28. 4. Periods • Periods of the periodic table are the horizontal rows across

29. 5. Groups • Groups or Families are vertical columns on the periodic table. • Currently we have 18 groups. We will use the 1-18 designations.

30. 6. Make it into a true statement • Elements of the same group are very similar.

31. 7. Change to a True Statement • The characteristic properties of the elements in a period change from group to group.

33. Property Summary

34. Property Summary

36. Transition Elements • Groups 3 - 12 are also called the Heavy Metals

37. Inner Transition • Rare Earth elements that are located in the bottom two rows (away from the rest of the table) of the periodic table

38. Groups with names • Group 1 = Alkali Metals • Group 2 = Alkaline Earth Metals • Group 18 = Inert or Noble Gases • Group 17 = Halogens

39. Periodic Table and Electron Configuration • The light metals compose the s block. • The transition elements are the d block. • The nonmetals are p block. • The inner transition (rare earth) metals are the f block.

40. Periodic Table 1 2 3 4 5 6 7 4 5 6 7 3 4 5 6 4f 5f p block Noble (inert) gases s block f block d block

41. Shorthand Electron Configuration Using Noble Gases to shorten the electron configuration

42. He 1s2 Ne 1s22s22p6 Ar = 1s22s22p63s23p6 Kr = 1s22s22p63s23p64s23d104p6 Xe = 1s22s22p63s23p64s23d104p65s24d105p6 Rn = 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6 Except He, do you see a trend in all of the noble gas configurations? What do they all end in? Complete the electron configurations for the Noble Gases (Hint: Group 18)

43. Shorthand Notation • We use the noble gases in shorthand notation • Find the closest noble gas that has an atomic number LESS than that of the element

44. Example • Ex. K • What is K’s atomic number? • 19 • Closest noble gas? • Ar • What is Ar’s atomic number? • 18 = 1s22s22p63s23p6 • = [Ar] 4s1 = Means the first 18 electrons are arranged like argon and the last electron is called the VALENCE ELECTRON (outermost shell)

45. Example for you to try • You try Ba • Ba = [Xe] 6s2 • Try Pb • Pb = [Xe] 6s24f145d106p2