Bohr model and electron configuration

1. Bohr model and electron configuration

2. 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.

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

4. 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)

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

6. High energy Low energy Low Frequency High Frequency X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

7. 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)

8. Line spectrum of various elements Helium Carbon

9. 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

10. 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

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

12. 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

13. 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

14. 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

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

16. The p Sublevel has 3 p orbitals

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

18. 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

19. Summary Starts at energy level

20. 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

21. Electron Configurations First Energy Level • only s sublevel (1 s orbital) • only 2 electrons • 1s2 Second Energy Level • s and p sublevels (s and p orbitals are available) • 2 in s, 6 in p • 2s22p6 • 8 total electrons

22. Third energy level • s, p, and d orbitals • 2 in s, 6 in p, and 10 in d • 3s23p63d10 • 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 • 4s24p64d104f14 • 32 total electrons

23. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

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

25. 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

26. 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

27. 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

28. 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

29. 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

30. 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

31. Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 is expected • But this is wrong!!

32. Chromium is actually • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principal applies to copper.

33. Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p64s23d9 • But the actual configuration is • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions

34. Great site to practice and instantly see results for electron configuration.

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