Electron Configuration of the Elements

1. Electron Configuration of the Elements

2. Hydrogen Emission Spectrum When hydrogen gas (H2) is placed in a CRT and a high voltage electrical current passed through it, the tube glows a violet colour. Johann Balmer 

3. When viewed through a spectroscope (prism), we observe four discrete lines—and NOT a continuous spectrum:

4. When an electron in the ground state of H absorbs energy, it gets promoted into a higher energy level. The electron is unstable in this higher energy level.

5. When the electron falls back to the ground state, energy is given off. This explains the “bands” of light emitted from a hydrogen discharge tube.

7. Here’s another way to look at it:

8. Hydrogen Emission Spectrum Electrons can only exist in certain energy levels (n) n = 1, n = 2, n = 3, n = 4, etc Energy levels in atom are quantized. This means that only certain E levels are allowed.

9. Each E level has one or more sublevels called orbitals • An orbital is a region of space where there is a high probability of finding an electron

10. Each orbital can hold a maximum of two electrons. • Electrons in an orbital will have opposite spin, designated ↑ (clockwise spin) or ↓ (counterclockwise spin).

11. high “probability” ? Heisenberg Uncertainty Principle We cannot simultaneously know the position and the momentum of an electron

12. Back to orbitals . . . • For n = 1 there is only one sublevel, called an s orbital. • Since this orbital is in the first energy level, it is called a 1s orbital. • s orbitals are spherical.

13. For n = 2 there are two sublevels: • 2s orbital (one of these) • 2p orbital (three of these) • a p orbital looks like this • A set of three p orbitals looks like this • We refer to the individual p orbitals as px, py, pz.

14. Let’s put these orbitals together . . .

15. For n = 3 (the third energy level) there are three sublevels: 3s orbital (one of these) 3p orbital (three of these) 3d orbital (five of these—see next slide) NB. Each orbital holds a maximum of 2 electrons

16. The d-orbitals

17. A funky look at d-orbitals

18. Your “bottom line” with d-orbitals: • There are five of them in each set. • eg. there are five 3d orbitals; five 4d orbitals, etc • 2 electrons in each, for a maximum of 10 electrons • How many columns are in the Transition Metal block (d-block) in the periodic table? 10 columns in the transition metals (5x2).

19. For n = 4 (the fourth energy level) there are four sublevels: 4s orbital (one of these) 4p orbital (three of these) 4d orbital (five of these) 4f orbital (seven of these—see next slide)

20. f-orbitals

21. How do electrons fill orbitals? Aufbau Principle • aka “Building-up” Principle • Electrons occupy orbitals beginning from the lowest energy orbital (i.e. the orbital closest to the nucleus) Start by filling 1s orbital • How many electrons per orbital? Each orbital can hold a maximum of two electrons—of opposite spin, don’t forget

22. Here is the order in which orbitals are filled . . .

23. Note the peculiarity . . . 3s is followed by 3p, which is followed by 4s, which is followed by 3d. There are others . . . (help is on the way)

24. How do the electrons of 7N fill the orbitals? 1s2 2s2 2p3 Overall for 7N: 1s2 2s2 2p3

25. Hund’s Rule More stable than . . .

26. Hund’s Rule • When filling p, d, f orbitals, pair electrons only when necessary

27. Aufbau Principle Mnemonic Device

28. Let’s write some electron configurations . . . 1H 1s1 2He 1s2 3Li 1s2 2s1 4Be 1s2 2s2 5B 1s2 2s2 2p1 6C 1s2 2s2 2p2 ↓ ↓ 10Ne 1s2 2s2 2p6

29. 11Na 11Na1s2 2s2 2p6 3s1 12Mg 1s2 2s2 2p6 3s2 13Al 1s2 2s2 2p6 3s2 3p1 ↓ ↓ 18Ar 1s2 2s2 2p6 3s2 3p6 19K 1s2 2s2 2p6 3s2 3p64s1 20Ca 1s2 2s2 2p6 3s2 3p64s2 21Sc 1s2 2s2 2p6 3s2 3p64s2 3d1

30. In the Periodic Table, what is the connection between the outermost electron configuration and family (column)? Alkali metals end in s1 Alkali earth metals end in s2 Halogens end in p5 Noble Gases end in p6

31. Groups (families) in PT

32. Putting it all together . . . • To write the electron configuration of any element, use the periodic table (play “Battleship”) and the Aufbau Principle mnemonic device.

33. Write the complete electron configuration for 1s2 2s2 2p6 3s2 3p64s2 3d10 4p65s2 4d10 5p1 ____________________________________ [Kr] preceding noble gas only 49In [Kr] 5s2 4d10 5p1

34. Practice Write the electron configuration for each of the following 15P 15P 1s2 2s2 2p6 3s2 3p3 33As [use noble gas core abbreviated form] 33As [Ar] 4s2 3d10 4p3

35. more practice . . . 13Al 13Al 1s2 2s2 2p6 3s2 3p1 26Fe 26Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 44Ru [ ] 44Ru [Kr] 5s2 4d6 52Te [ ] 52Te [Kr] 5s2 4d10 5p4

36. Exceptional Electron Configurations 1. Write the expected electron configuration of 24Cr [Ar]4s2 3d4 Actual electron configuration is [Ar]4s1 3d5 Special stability associated with half-filled p, d, f orbitals

37. Now write the electron configuration for 42Mo 42Mo [Kr] 5s14d5 Notice any similarity with Cr?

38. 2. Write the expected electron configuration of 29 Cu: [Ar]4s2 3d9 Actual electron configuration is [Ar]4s1 3d10 In this way Cu has completely filled 3rd energy level (Copper is a very stable metal)

39. Now write the electron configuration for silver (47Ag) and gold (79Au). Use the noble gas core abbreviated forms. 47Ag [Kr] 5s1 4d10 79Au [Xe] 6s14f14 5d10