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CHAPTER 5

CHAPTER 5. THE ELECTRON. Newtonian Physics. Describes normal particles moving at normal speed and the laws that pertain to them Example: car moving at 55 mph. QUANTUM PHYSICS. Describes extremely small particles moving at speeds close to the speed of light

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CHAPTER 5

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  1. CHAPTER 5 THE ELECTRON

  2. Newtonian Physics • Describes normal particles moving at normal speed and the laws that pertain to them • Example: car moving at 55 mph

  3. QUANTUM PHYSICS • Describes extremely small particles moving at speeds close to the speed of light • Example: electron moving in an atom

  4. Increasing energy nucleus BOHR MODEL • Said e- can be found at certain energy levels (like a ladder) • e- can move up and down that ladder by absorbing quantums of energy

  5. QUANTUM- Max Planck • Quantum= amount of energy an e- needs to absorb to move from lower to higher energy level • Photon = quanta of radiant energy (little packet of light energy)

  6. ELECTROMAGNETIC ENERGY • A radiant energy- types include visible light, radio, infrared, UV and X-Ray • Travels at the speed of light; has variations in electric and magnetic fields

  7. If we plot the strength of these variations vs. time, we get waves • Wavelength (l)= the physical distance between crests l

  8. FREQUENCY (f )= the # of wave breaks (crests) that occur in a unit of time • Measured in Hertz (Hz)- 1 peak/second

  9. ELECTROMAGNETIC SPECTRUM

  10. MOVIE

  11. Electromagnetic Spectrum Visible Light- 700 nm (red) – 400 nm (violet) Low energy high energy

  12. SPECTROSCOPY • The study of the interaction between matter or substances and radiant energy • Radiant energy = continuous excited energy

  13. Explanation of Atomic Spectra • Spectrum = occurs when you look at the light given off thru a diffraction grating (separates light into its diff. wavelengths- visible light) • Uses: can ID unknown gases; determine composition of the stars

  14. How it works: • Atoms can lose or gain energy • An atom gets excited but can’t hold onto that energy so it loses the energy in the form of light

  15. Excited electron- absorbed a quantum of energy- moved to higher energy level Electron can’t hold onto the Energy- falls back down to ground state- release photon (packet of light energy) e- e- Energy levels Ground state (lowest energy level an electron Occupies)

  16. Types of Spectrums Chemistry of Fireworks video Fireworks

  17. CONTINUOUS SPECTRUM • White light, gives you an uninterrupted spectrum of light • Get the colors of the rainbow blended together (ROYGBIV)

  18. Emission spectrum of hydrogen (bright line spectrum) EMISSION SPECTRUM • Created by the light emitted when the e- drop from their excited state to lower energy states • Get a series of lines in different colors on a blacked out background • Each line corresponds to a particular l of light emitted by an atom

  19. Absorption spectrum of hydrogen (dark-line spectrum) ABSORPTION SPECTRUM • Spectrum created by determining the energy absorbed by the atom as it moves up to higher energy states • See a rainbow background w/dark lines in the spot where the l of light have been absorbed • These lines correspond exactly to the colored lines of the emission spectrum

  20. PHOTOELECTRIC EFFECT • Albert Einstein (1921) • Found that when light shines on a metal surface at a very specific frequency, the metal emits electrons • Uses: automatic doors; solar calculator • Proved light has both wavelike and particle like properties

  21. WAVE PARTICLE DUALITY OF NATURE • Louis de Broglie • Used Einstein and Planck’s formulas • Described the behavior of the electron in orbit • Derived this equation: h l = mv h = Planck’s constant 6.626 x 10-34 J·s V = velocity M = mass l = wavelength

  22. It predicts that all moving particles have wave characteristics (even a moving car- just that the l is so small we can’t see it) • WAVE PARTICLE DUALITY OF NATURE = all particles have properties of waves and waves have properties of particles

  23. HEISENBERG UNCERTAINTY PRINCIPLE (1927) • Found that it is impossible to make any measurement on an object without disturbing the object • Electrons were detected by their interactions with photons (packets of light that have the same energy as an electron)

  24. photon Change of Position and v e- collision photon e- Uncertainty cont. • Interaction between the photon and electron alters the position and the velocity of the electron • HEISENBERG UNCERTAINTY PRINCIPLE= it is impossible to know precisely both the location and the velocity of a particle at the same time

  25. Schrödinger Wave Equation (1926) • Erwin Schrödinger developed a wave equation that laid the foundation for the quantum theory and quantum mechanical view of the atom • QUANTUM THEORY= describes the mathematical wave properties of electrons and other small particles • He used an equation to predict the most probable location of an electron around a nucleus (found that electron did not travel in neat paths)

  26. ATOMIC ORBITAL • A 3D region around the nucleus that indicated the probable location of electron (the electron cloud) H- 53 pm from nucleus (90% probablity)

  27. QUANTUM NUMBERS • From Schrodinger’s equation we get a set of variables that specify the properties of the electron in orbit • Each electron can be described by a unique set of 4 quantum numbers describing the electron’s behavior • Principle, angular momentum, Magnetic and spin Video- Electron Behavior

  28. PRINCIPLE QUANTUM NUMBER (n) • Indicates the main energy level occupied by the electron • As the energy level (n) increases the electron’s energy increases (more energy further from the nucleus) • Can have more than one electron in the energy level

  29. (n) cont. • Maximum number of electron in each level represented with the following formula: • 2n2 where n = energy level n = 1 n = 2 2 (1)2 = 2e- 2(2)2 = 8e- n = 3 n = 4 2 (3)2 = 18e- 2 (4)2 = 32 e- n = 5-7 same as n = 4; only 32 e-

  30. ANGULAR MOMENTUM QUANTUM NUMBER (L) • The orbitals or clouds in each energy level occur in different shapes • This quantum number indicates the shape of the orbital or the name of the sublevel

  31. (L) cont. • Types of Sublevels • s = sphere shape • p = peanut shape • d = double peanut • f = flower shape • The value of L = the value of n • n = 1 , has 1 sublevel (s) • n = 2, has 2 sublevels ( s & p) • n = 3, has 3 sublevels (s, p, &d) • n = 4 has four sublevels (s, p, d, & f)

  32. MAGNETIC QUANTUM NUMBER (m) • An orbital can have the same shape but be found in different orientations around the nucleus • This quantum number represents the number of positions found in each sublevel

  33. (m) cont. • s = 1 position max of 2e- • p = 3 positions max of 6e- • d = 5 positions max of 10 e- • f = 7 positions max of 14 e- • Each position can contain a pair of electrons

  34. SHREK Example? Okay, um, ogres are like onions. (he holds out his onion) DONKEY (sniffs the onion) They stink? SHREK Yes - - No! DONKEY They make you cry? SHREK No! DONKEY You leave them in the sun, they get all brown, start sproutin' little white hairs. SHREK No! Layers! Onions have layers. Ogres have layers! Onions have layers. You get it? We both have layers. (he heaves a sigh and then walks off) DONKEY (trailing after Shrek) Oh, you both have layers. Oh. {Sniffs} You know, not everybody likes onions. Cake! Everybody loves cakes! Cakes have layers. SHREK I don't care... what everyone likes. Ogres are not like cakes. DONKEY You know what else everybody likes? Parfaits. Have you ever met a person, you say, "Let's get some parfait," they say, "Hell no, I don't like no parfait"? Parfaits are delicious. SHREK No! You dense, irritating, miniature beast of burden! Ogres are like onions! And of story. Bye-bye. See ya later. DONKEY Parfaits may be the most delicious thing on the whole damn planet. SHREK You know, I think I preferred your humming. DONKEY Do you have a tissue or something? I'm making a mess. Just the word parfait make me start slobbering.

  35. 1s SPIN QUANTUM NUMBER (s) • Like the earth, the electron in an orbital can be thought of as spinning on its axis • This quantum number represents that spinning action • The electron can spin clockwise or counterclockwise • We represent the spin with up or down arrows as the electron

  36. ELECTRON CONFIGURATION • DEF.=AN ARRANGEMENT OF THE ELECTRONS IN AN ATOM • RULES: • 1. Aufbau principle- electrons occupy the lowest energy level that can receive it (fill the lower level up first)

  37. 2. Pauli Exclusion Principle- no two electrons in the same atom can have the same set of quantum #’s • 3. Hund’s Rule- orbitals of equal energy are each occupied by 1 electron before any orbital is occupied by a second electron • all electrons in single occupied orbitals must have the same spin

  38. TYPES OF CONFIGURATION

  39. 1s 2s 2p 3s 3p 1. ORBITAL NOTATION • Most detailed • Uses all four quantum numbers • Orbital – a space occupied by a pair of electrons • Example: aluminum; atomic #13 so 13 electrons

  40. DIAGONAL RULE (sublevel diagram) • Overlapping of sublevels = when you get to the larger energy levels the d and f sublevels are so close together they find greater stability if they overlap some of the lower orbitals

  41. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p 8s

  42. 1s2s 2p 3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f6s 6p 6d7s 7p

  43. Diagonal rule = use to help figure out when overlapping occurs; follow arrows and list the orbitals in the order listed

  44. 1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p

  45. Example: iron electrons = 26

  46. 1s 2s 2p 3s 3p Example: iron electrons = 26 4s 3d

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