Chapter 5

1. Chapter 5 Electrons in Atoms

2. Rutherford’s Model • Discovered the nucleus • Small dense and positive • Electrons moved around in Electron cloud

3. Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Energy separates one level from another.

4. Bohr’s Model Nucleus Electron Orbit Energy Levels

5. Bohr’s Model Nucleus Electron Orbit Energy Levels

6. Bohr’s Model } • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels Fifth Fourth Third Increasing energy Second First Nucleus

7. The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • Quanta - the amount of energy needed to move from one energy level to another. • Quantum leap in energy. • Schrödinger derived an equation that described the energy and position of the electrons in an atom • Treated electrons as waves

8. The Quantum Mechanical Model • a mathematical solution • It is not like anything you can see.

9. The Quantum Mechanical Model • Does have energy levels for electrons. • Orbits are not circular. • It can only tell us the probability of finding an electron a certain distance from the nucleus.

10. The Quantum Mechanical Model • The electron is found inside a blurry “electron cloud” • An area where there is a chance of finding an electron.

11. Atomic Orbital’s and Quantum Numbers • Principal Quantum Number (n) = indicates the main energy level occupied by the electron. • Positive integers 1,2,3,… • Within each energy level the complex math of Schrödinger's equation describes several shapes. • These are called atomic orbitals • Regions where there is a high probability of finding an electron. • The total number of orbital’s that exist in a main energy level is equal to n2

12. Angular Momentum Quantum Number • (l) indicates the shape of the orbital. • Except at E1, orbitals of different shapes (sublevels) exist for a given value of n. • The number of orbital shapes possible is equal to n. • l = 0, 1, 2, … n-1 (all positive integers)

13. Shapes of Orbitals • n = 1 l = 0 one orbital s • n = 2 l = 0, 1 two orbitals s, p • n = 3 I = 0, 1, 2 three orbitals s, p, d

14. S orbitals • One s orbital for every energy level • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals.

15. P orbitals • Start at the second energy level • 3 different directions • 3 different shapes (dumbell) • Each can hold 2 electrons

16. P Orbitals

17. D orbitals • Start at the third energy level • 5 different shapes • Each can hold 2 electrons

18. F orbitals • Start at the fourth energy level • Have seven different shapes • 2 electrons per shape

19. F orbitals

20. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6 2 5 10 3 d 7 14 4 f

21. Magnetic Quantum Number • ml • the orientation of the orbital in 3-D space. (x, y, z) • the values of ml range from –l to +l • Ex: n=1 l=0 ml=0 n=2 l=0, 1 ml= -1,0, 1 n=3 l= ? ml= ?

22. Spin Quantum Number • ms • electrons are not stationary particles, they spin • they can only spin in two directions, clockwise and counterclockwise (designations we have assigned them) • the values of ms are +1/2 or – 1/2

23. The Address of an Electron • No two electrons have the same 4 quantum numbers. • what I know from the quantum numbers of an electron: • 1, 0, 0, +1/2 • first principal energy level, s orbital, (x,y,z) axis, spinning clockwise • 3, 1, -1, -1/2 • Third principal energy level, p orbital, x axis, spinning counterclockwise

24. Aufbau Principle • German for Building Up • When the we build an atom with its various electrons we start with the lowest energy level and build up.

25. 7s 7p 7d 7f 7g 7h 7i 6s 6p 6d 6f 6g 6h 5s 5p 5d 5f 5g 4s 4p 4d 4f all orbitals 3s 3p 3d beyond f are 2s 2p theoretical 1s The easy way to remember filling order

26. 7p 6d 6p 5d 5p 5f 4f 4p 4d 3p 3d 2p 7s 6s 5s 4s Increasing energy 3s 2s 1s

27. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

28. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

29. Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . • Let’s determine the electron configuration for Phosphorus • Need to account for 15 electrons

30. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first to electrons go into the 1s orbital • Notice the opposite spins • only 13 more

31. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more

32. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more

33. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more

34. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes • 3 unpaired electrons • 1s22s22p63s23p3

35. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s The easy way to remember • 1s2 • 2 electrons

36. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 • 4 electrons

37. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 • 12 electrons

38. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 • 20 electrons

39. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 • 38 electrons

40. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 • 56 electrons

41. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 • 88 electrons

42. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 • 118 electrons

43. Rewrite when done • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 • Group the energy levels together • 1s2 2s2 2p63s2 3p6 3d10 4s2 4p6 4d10 4f145s2 5p6 5d105f146s2 6p6 6d107s2 7p6

44. Exceptions to Electron Configuration

45. Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Filled and half-filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order of d orbitals

46. Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p63d24s2 • Vanadium - 23 electrons 1s22s22p63s23p63d34s2 • Chromium - 24 electrons • 1s22s22p63s23p63d44s2 is expected • But this is wrong!!

47. Chromium is actually • 1s22s22p63s23p63d54s1 • Why? • This gives us two half filled orbitals.

48. Chromium is actually • 1s22s22p63s23p63d54s1 • Why? • This gives us two half filled orbitals.

49. Chromium is actually • 1s22s22p63s23p63d54s1 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principle applies to copper.

50. Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p63d94s2 • But the actual configuration is • 1s22s22p63s23p63d104s1 • This gives one filled orbital and one half filled orbital. • Remember these exceptions • d4s2  d5 s1 • d9s2  d10s1