Bohr model and electron configuration
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Bohr's model proposed that electrons orbit the nucleus in fixed, quantized energy levels. Later models, like the quantum mechanical model, provided a more accurate description of electron behavior. Electrons fill atomic orbitals according to specific rules, building up the electron configuration of elements in a predictable order based on energy levels. Exceptions occur when half or completely filled orbitals result in a lower energy configuration. Practice filling electron configurations to understand these patterns.
Bohr model and electron configuration
- 1. Bohr model and electron configuration Mrs. A. Kay Chem 11
- 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 Electron Orbit 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/chemistr y/essentialchemistry/flash/linesp16.swf
- 6. Low High energy energy Radio Micro Infrared Ultra- X- Gamma waves waves . violet Rays Rays Low High Frequency Frequency Long Short Wavelength 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
- 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 Fifth from the nucleus Fourth Increasing energy means more energy. Third There is no “in between” Second energy Energy Levels 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 # of shapes Max # of energy Sublevel (orbitals) electrons level s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 4
- 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 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p 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 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p The first to electrons go 3s into the 1s orbital 2p 2s Notice the opposite spins only 13 more 1s
- 26. 7p 6d 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p The next electrons go 3s into the 2s orbital 2p 2s only 11 more 1s
- 27. 7p 6d 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p • The next electrons go 3s into the 2p orbital 2p 2s • only 5 more 1s
- 28. 7p 6d 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p • The next electrons go 3s into the 3s orbital 2p 2s • only 3 more 1s
- 29. 7p 6d 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p • The last three electrons 3s go into the 3p orbitals. 2p • They each go into 2s separate shapes • 3 unpaired electrons 1s • 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.