1. Bohr model and electron configuration Sandy Bohr’s Model

2. 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.

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

4. Bohr postulated that: Each orbt corresponds to fixed energy(stationary orbits or shells) Electrons cannot exist between orbits(move in a specified paths) 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(becomes inert)

5. 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

6. Drawbacks  Could not explain multi electron atoms like He, Li, Be etc…  Could not account for the Zeeman effect.  Could not justify the quantization of angular momentum.  Could not account for the formation of chemical bonds.

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

8. 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

9. 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

10. 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

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

12. The p Sublevel has 3 p orbitals

13. 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

14. 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

15. Summary Starts at energy level

16. 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

17. 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

18. 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

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

20. 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

21. 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!!

22. 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. Lower energy higher stability The same principal applies to copper.

23. 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

24. Practice Time to practice on your own filling up electron configurations. Do electron configurations for the first 20 elements on the periodic table.