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Chapter 5 “Electrons in Atoms”

Chapter 5 “Electrons in Atoms”. Ernest Rutherford’s Model. Discovered dense + piece at nucleus Electrons move around it, like planets around sun Mostly empty space Didn’t explain chemical properties of elements. Niels Bohr’s Model. Why don’t electrons fall into nucleus?

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Chapter 5 “Electrons in Atoms”

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  1. Chapter 5“Electrons in Atoms”

  2. Ernest Rutherford’s Model • Discovered dense + piece at nucleus • Electrons move around it, like planets around sun • Mostly empty space • Didn’t explain chemical properties of elements

  3. Niels Bohr’s Model • Why don’t electrons fall into nucleus? • Move like planets around sun. • specific circular orbits at different levels. • An amount of fixed energy separates one level from another.

  4. The Bohr Model of the Atom I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another. Niels Bohr

  5. Bohr’s model • Energy level measure of fixed energy e- can have • Like rungs of ladder • e- can’t exist between energy levels • can’t stand between rungs on ladder • Unlike ladder, energy levels not evenly spaced • Higher energy levels closer together – less energy needed for jump

  6. The Quantum Mechanical Model • Energy is “quantized” - in chunks. • A quantum is exact energy needed to move e- one energy level to another. • Since energy of atom is never “in between” quantum leap in energy must exist. • Erwin Schrodinger (1926) derived equation described energy and position of e- in an atom Schrodinger

  7. The Quantum Mechanical Model • Very small things behave differently from big things • quantum mechanical modelis a mathematical solution • not like anything you can see (like plum pudding!)

  8. The Quantum Mechanical Model • Has energy levels for e- • Orbits not circular. • only tells us probabilityof finding e- a certain distance from nucleus.

  9. The Quantum Mechanical Model • e- found inside blurry “electron cloud” (area where there’s chance of finding e- • Like fan blades or bicycle wheel

  10. Atomic Orbitals • Principal Quantum Number (n) = the energy level of e- (1, 2, 3, etc.) • Within each energy level, Schrodinger’s equation describes several shapes. • called atomic orbitals - regions of space w/high probability of finding e- • Sublevels like rooms in a hotel • letters s, p, d, and f

  11. Principal Quantum Number (n) denotes the energy level (shell) e- is located. Maximum number of e- that can fit in energy level is: 2n2 How many e- in level 2? level 3?

  12. Individual orbitals First possible energy level Maximum electrons # of orbitals s spherical 2 1 1st p dumbell 6 3 2nd 10 d clover leaf 5 3rd 14 f complicated 7 4th

  13. By Energy Level • Energy Level 1 (n=1) • Has only s sublevel • 1s (1 orbital) • w/ only 2 e- • 2 total e- • Energy Level 2 (n=2) • Has s and p sublevels available • 2s (1 orbital) • w/ 2 e- • 2p (3 orbitals) • w/ 6 e- • 2s22p6 • 8 total e- 2 2 6

  14. By Energy Level 2 • Energy level 3 (n=3) • s, p, and d sublevels • 3s (1 orbital) • w/ 2 e- • 3p (3 orbitals) • w/ 6 e- • 3d (5 orbitals) • w/ 10 e- • 3s23p63d10 • 18 total e- • Energy level 4 (n=4) • s, p, d, and f sublevels • 4s (1 orbital) • w/ 2 e- • 4p (3 orbitals) • w/ 6 e- • 4d (5 orbitals) • w/ 10 e- • 4f (7 orbitals) • w/ 14 e- • 4s24p64d104f14 • 32 total e- 2 6 6 10 14 10

  15. By Energy Level • Any more than the fourth and not all orbitals fill up. • You simply run out of e-’s • orbitals do not fill up in neat order. • energy levels overlap • Lowest energy fill first. s and p orbitals 1:20 d orbitals 3:40

  16. Electron distribution in an Atom Energy Level 3 3d 3p 3s ENERGY 2 2p 2s 1 1s NUCLEUS

  17. Section 5.2Electron Arrangement in Atoms • OBJECTIVES: • Describe how to write the electron configuration for an atom.

  18. Section 5.2Electron Arrangement in Atoms • OBJECTIVES: • Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

  19. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s aufbau diagram - page 133 Aufbau is German for “building up”

  20. Electron Configurations… • …the way electrons arranged in various orbitals around nuclei of atoms. Three rules tell us how: • Aufbau principle- electrons enter the lowest energy first. • causes difficulties b/c overlap of orbitals of different energies – follow diagram! • Pauli Exclusion Principle- 2 electrons max per orbital (hotel room) - different spins

  21. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. To show the different direction of spin, a pair in the same orbital is written as: Wolfgang Pauli

  22. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. 1) Principal quantum number 2) Angular momentum quantum number 3) Magnetic quantum number 4) Spin quantum number

  23. Electron Configurations • Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. • write e- configuration for Phosphorus • We need to account for all 15 e-’s in phosphorus

  24. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first two electrons go into the 1s orbital Notice the opposite direction of the spins • only 13 more to go...

  25. 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

  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 2p orbital • only 5 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 3s orbital • only 3 more...

  28. 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 (Hund’s) • 3 unpaired electrons • = 1s22s22p63s23p3 Orbital notation

  29. Stairstep pattern

  30. An internet program about electron configurations is: Electron Configurations (Just click on the above link) I electron config 3:24

  31. Orbitals fill in an order • Lowest energy  higher energy. • Adding e-’s changes energy of orbital. Full orbitals are best situation. • half filled orbitals  lower energy, next best • more stable. • Changes filling order

  32. Write the electron configurations for these elements: • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons • 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 (expected) • But this is not what happens!!

  33. Chromium is actually: • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals(the others are all still full) • Half full is slightly lower in energy. • The same principal applies to copper.

  34. Copper’s electron configuration • Copper has 29 electrons so we expect: 1s22s22p63s23p63d94s2 • But the actual configuration is: • 1s22s22p63s23p63d104s1 • This change gives one more filled orbital and one that is half filled. • Remember these exceptions: d4, d9

  35. Irregular configurations of Cr and Cu Chromium steals a 4s electron tomakeits 3d sublevel HALF FULL Copper steals a 4s electron toFILLits 3d sublevel

  36. Section 5.3Physics and the Quantum Mechanical Model • OBJECTIVES: • Describe the relationship between the wavelength and frequency of light.

  37. Section 5.3Physics and the Quantum Mechanical Model • OBJECTIVES: • Identify the source of atomic emission spectra.

  38. Section 5.3Physics and the Quantum Mechanical Model • OBJECTIVES: • Explain how the frequencies of emitted light are related to changes in electron energies.

  39. Section 5.3Physics and the Quantum Mechanical Model • OBJECTIVES: • Distinguish between quantum mechanics and classical mechanics.

  40. Light • Study of light led to development of quantum mechanical model. • Light is electromagnetic radiation. • EM radiation: gamma rays, x-rays, radio waves, microwaves • Speed of light = 2.998 x 108 m/s • abbreviated “c” • “celeritas“ Latin for speed • All EM radiation travels same in vacuum

  41. - Page 139 “R O Y G B I V” Frequency Increases Wavelength Longer

  42. Crest Wavelength Amplitude Trough Parts of a wave Origin

  43. Electromagnetic radiation propagates through space as a wave moving at the speed of light. Equation: c = c = speed of light, a constant (2.998 x 108 m/s)  (lambda) = wavelength, in meters (nu) = frequency, in units of hertz (hz or sec-1)

  44. Wavelength and Frequency • inversely related • one gets bigger, other smaller • Different frequencies of light different colors • wide variety of frequencies (spectrum)

  45. - Page 140 Use Equation: c =

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

  47. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY

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