1. Ms. Cleary Chem 11

2. A model A representation or explanation of a reality that is so accurate and complete that it allows the model builder to predict events. Scientific Method leads to model building Gather data, develop a model, formulate a hypothesis, test and modify the model.

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. Amounts of energy separate one level from another.

4. Bohr’s Model Nucleus Electron Orbit Energy Levels Nucleus Electron Orbit Energy Levels

5. Bohr postulated that: Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is away from the nucleus An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)

6. How did he develop his theory? He used mathematics to explain the visible spectrum of hydrogen gas http://www.mhhe.com/physsci/chemistr y/essentialchemistry/flash/linesp16.swf http://www.mhhe.com/physsci/chemistr y/essentialchemistry/flash/linesp16.swf

7. Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low energy High energy Low Frequency High Frequency Long Wavelength (700 nm) Short Wavelength (400 nm) Visible Light Energy and Visible Light

8. The line spectrum electricity passed through a gaseous element emits light at a certain wavelength Can be seen when passed through a prism Every gas has a unique pattern (color)

9. Line spectrum of various elements

10. Bohr’s Triumph His theory helped to explain periodic law Halogens are so reactive because it has one e- less than a full outer orbital Alkali metals are also reactive because they have only one e- in outer orbital

11. Drawback Bohr’s theory did not explain or show the shape or the path traveled by the electrons. His theory could only explain hydrogen and not the more complex atoms

12. Further away from the nucleus means more energy. There is no “in between” energy Energy Levels First Second Third Fourth Fifth Increasing energy }

13. Complete Bohr Diagrams for the Following: Mg Li Ne F

14. The Quantum Mechanical Model Energy is quantized. It comes in chunks. A quanta is the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy. Schrödinger derived an equation that described the energy and position of the electrons in an atom

15. Atomic Orbitals Principal Quantum Number (n) = the energy level of the electron. 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

16. Orbitals Electrons spin around the nucleus creating an electron cloud. The electron clouds come in 4 different shapes, called orbitals. The four orbitals are called s, p, d, and f.

17. Each orbital is capable of holding different numbers of electrons: Orbital# of Electrons s2 p6 d10 f14

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

19. P orbitals Start at the second energy level 3 different directions 3 different shapes Each orbital can hold 2 electrons

20. The p Sublevel has 3 p orbitals

21. The D sublevel contains 5 D orbitals The D sublevel starts in the 3 rd energy level 5 different shapes (orbitals) Each orbital can hold 2 electrons

22. The F sublevel has 7 F orbitals The F sublevel starts in the fourth energy level The F sublevel has seven different shapes (orbitals) 2 electrons per orbital

23. Summary Starts at energy level

24. 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

25. Electron Configurations First Energy Level only s sublevel (1 s orbital) only 2 electrons 1s 2 Second Energy Level s and p sublevels (s and p orbitals are available) 2 in s, 6 in p 2s 2 2p 6 8 total electrons

26. Levels Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s 2 3p 6 3d 10 18 total electrons Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, and 14 in f 4s 2 4p 6 4d 10 4f 14 32 total electrons

27. Electron Configurations Electron configurations are a shorthand for writing exactly what was in the energy level diagrams. Electron configuration for O is: 1s 2 2s 2 2p 4 period orbital # of electrons Electron configuration for Ar is: 1s 2 2s 2 2p 6 3s 2 3p 6

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

29. Electron Configuration Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to.

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

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

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

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

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

35. Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Half filled orbitals have a lower energy. Makes them more stable. Changes the filling order

36. Write these electron configurations Titanium - 22 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Chromium - 24 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected But this is wrong!!

37. Chromium is actually 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why? This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to copper.

38. Copper’s electron configuration Copper has 29 electrons so we expect 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This gives one filled orbital and one half filled orbital. Remember these exceptions

40. Electron Configuration and the Periodic Table Groups 1 and 2 represent the s orbital Groups 13-18 represent the p orbital Groups 3-12 represent the d orbital Lanthanides and Actinides represent f orbital

42. Practice 1. Time to practice: Draw the following energy level diagrams on your own filling up electron configurations: H, He, Be, N, Na, Ni, Br, 2. Do electron configurations for the elements listed in #1.