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Electrons

Electrons. Model of atoms. The Wave Nature of Light. Light was first recognized as manifestation of electromagnetic energy and it was called electromagnetic radiation. http://college.hmco.com/chemistry/shared/media/animations/electromagneticwave.html.

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Electrons

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  1. Electrons Model of atoms

  2. The Wave Nature of Light Light was first recognized as manifestation of electromagnetic energy and it was called electromagnetic radiation

  3. http://college.hmco.com/chemistry/shared/media/animations/electromagneticwave.htmlhttp://college.hmco.com/chemistry/shared/media/animations/electromagneticwave.html • Produced by motion of electrically charged particles • Move through vacuum (at 3.00 x 108 m/s or 186,282 mi/hr), air and other substances • Have characteristic wavelengths/frequencies • Visible radiation has wavelengths between 400 nm (violet) and 750 nm (red) As electromagnetic wave, light has some characteristics in common with all forms of electromagnetic energy

  4. Wave: disturbance of medium which transports energy without permanently transporting matter • Medium • Substance or material that carries wave • Merely carries wave from source to other location

  5. Light is a repeating waveform in motion Rest Position: no energy present (amt. energy found in wave)

  6. Frequency, wavelength, and velocity are inversely proportional to each other • If frequency ↑, wavelength ↓ • Ex. Purple light has a frequency of 7.42 x 1014 Hz.  What is its wavelength? • c =  • 3.00 x 108 m/s = 7.42 x 1014 Hz () •  = 4.04 x 10-7

  7. Photons: packets of energy that make up light • Each carries specific energy related to its wavelength • Photons of short wavelength (blue light) carry more energy than long wavelength (red light) photons Einstein successful explained photo-electric effect within context of quantum physics

  8. Light is series of energy packets passing through space • Size of energy packets vary/change color of light • Quantized: electron limited to specific quantities of energy, not random value of energy • Distance between energy packets = wavelength • # photons passing point in period of time = frequency Particulate Theory of Light

  9. Suggested H’s electron moves around nucleus in only certain allowed orbits • Smaller orbit, lower energy level • Larger orbit, higher energy level • Electron can have different energy levels • Ground state: lowest level • Excited state: atom gains energy One of simplest working models of atom developed by NielsBohr

  10. Assigned quantum number, n, to each orbit Electron in ground state (1st energy level, n = 1) Does not radiate energy Quantum jump: electron moves from one energy level to another by gaining energy (excited state) or losing energy (ground state) in continuously changing amounts • Electron drops from higher to lower energy orbit • Photon with specific energy emitted as light • Shown as different colored line spectrums (atomic spectrum) • Every element has its own

  11. http://college.hmco.com/chemistry/shared/media/animations/h2linespectrum.htmlhttp://college.hmco.com/chemistry/shared/media/animations/h2linespectrum.html • Atomic emission spectrum (amount of electromagnetic radiation of each frequency gas emits when heated/excited) • Photon hits metal, is absorbed as electron takes up energy • Einstein deduced each photon possesses energy • Different metals require different minimum frequencies for electrons to exhibit photoelectric effect • Above threshold frequency, # electrons ejected depend on intensity of light • If photon’s frequency below minimum, electron remains bound to metal surface

  12. Fe

  13. http://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=chapter_03&folder=emission_absorptionhttp://www2.wwnorton.com/college/chemistry/gilbert/tutorials/interface.swf?chapter=chapter_03&folder=emission_absorption • Ephoton = nhν • n = # photons • h = Planck's constant, 6.626 x 10-34 J·s • ν= frequency of radiation Convert wavelength from nanometers to meters: 1 x 10-9 meters = 1 nm

  14. Groups of lines observed in emission spectrum of hydrogen atoms UV visibleIR IR IR

  15. Calculate the energy of a photon of yellow light with a frequency of 5.09 x 1014 s-1. E= nhn =(1)(6.626 x 10-34 J.s)(5.09 x 1014 s-1)= 3.37 x 10 -19 J Calculate the energy of a photon of wavelength 5.00 x 104 nm (infrared). E = nhn = nhc/l =(1)(6.626 x 10-34 J·s)(3.00 x 108 m/s) (5.00 x 10-5 m) = 3.98 x 10-21 J Calculate energy of mole of photons of yellow light with a frequency of 5.09 x 1014 s-1. E = nhn = (6.022 x 1023)(6.626 x 10-34 J·s) (5.09 x 1014 s-1) = 2.03 x 105 J

  16. What is the frequency in hertz of blue light having a wavelength of 425 nm? • 7.06 X 1014 Hz • A certain substance strongly absorbs infrared light having a wavelength of 6,500 nm.  What is the frequency in hertz of this light? • 4.62 X 1013 Hz • Yellow light has a wavelength of 600 nm. What is its frequency in hertz? • 5.00 X 1014 Hz • Green light has a wavelength of 550 nm. What is its frequency in hertz? • 5.45 X 1014 Hz

  17. Intense microwaves have a frequency of 9.5 X 1011 Hz. What is the wavelength of these particular microwaves? • 3.16 X 10-4 m  =  0.316 mm = 316 micrometers = 3.16 X 105 nm • Infrared waves can be seen if you look down the railroad tracks or a road on a hot day. They heat the air as they go past causing the air to refract or bend the light.  If infrared rays of 9.75 X 1013 Hz are being reflected off the tracks or road what will be the size of  the wavelengths in micrometers? • 3.08 X 10-6 m  = 3.08 micrometers • A sunbather forgot their sunblock. On the beach they get a unheathy dose of UV radiation of 5.66 X 1016 Hz.  What is the wavelength of these particular UV waves? • 5.30 X 10-9 m  = 5.00 nm

  18. Sodium vapor lamps are used to sometimes light streets. If the frequency of the light coming from them is 5.09 X 1014 Hz what is the energy in each photon? • 3.37 X 10-19 J/photon • What is the energy of each photon of red light that has a frequency of 4.0 X 1014 Hz? • 2.65 X 10-19 J/photon • Calculate the energy in joules/photon for green light having a wavelength of 550 nm. • 3.62 X 10-19 J/photon

  19. Microwaves are used to heat food in microwave ovens. The microwave radiation is absorbed by moisture in the food. This heats the water, and as water becomes hot, so does the food.   How many photons having a wavelength of 3.00 mm would have to be absorbed by 1.00 g of water to raise its temperature by 1oC? • 6.63 X 10-19 J/photon; 6.31 X 1022 photons • The wavelengths of X-rays are much shorter than those of ultraviolet or visible light. Show quantitatively why continued exposure to X-rays is more damaging than exposure to sunlight. • X-rays: 6.63 X 10-17 J/photon, UV rays: 6.63 X 10-19 J/photon, X-rays are 100 times more powerful than UV rays.

  20. Prepare 0.5 M solutions of barium/calcium/potassium lithium/sodium/ and strontium chloride (nitrates can be used). • Fold the end of a nichrome or platinum wire into a ball and tap the straight end to a wooden stick. • Dip the end into dilute hydrochloric acid, hold it in the burner until no color shows. • Dip the end into a test tube of one of the solutions, place it in the flame, record color on chart. Flame Tests

  21. Read 5.1, pp. 116-126 • Q pg. 126, #8-10 • Q pp. 146-147, #33, 36, 37, 65, 66, 70, 71, 74, 76 Homework:

  22. By the mid-1920s, scientists convinced Bohr atomic model was incorrect, formulated new explanations of how electrons arranged in atoms • de Broglie (de-broy-lee) 1924 • If light could act as both particles and waves, so could electrons • Since energy E of photon equals Planck’s constant times frequency f, or E = hf, momentum p of electron would equal Planck’s constant divided by wavelength

  23. Impossible to determine with perfect accuracy both position and momentum of particle simultaneously • Making measurements on object alters location/ momentum enough to disturb accuracy of reading location/momentum • More certain we are about particle's position, less certain we are about its velocity, and vice versa • Bohr ran into trouble because he tried to predict electron’s movement too precisely • Restricting electron to certain locations and having it move in orbits violated Heisenberg Uncertainty Principle Heisenberg’s Uncertainty Principleapplied de Broglie’s hypothesis

  24. Electromagnetic radiation has dual "personality“ • Acts like waves/photons with no mass • Displays behaviors characteristic of any wave (reflection, refraction, diffraction, interference, exhibits Doppler effect) that would be difficult to explain with pure particle-view Is light a wave or a particle?

  25. Equation contains both wave and particle terms • Electrons do not have planetary orbit • Location of electron is probability, not certain position Principles of Quantum Mechanics (Schrödinger)

  26. Describes mathematically wave properties of electrons and other very small particles • Cloud shapes now called orbitals • 3-D region around nucleus that indicates probable location of electron (“probability regions”) • Electrons not confined to fixed circular path Quantum Theory

  27. Specify properties of atomic orbitals/electrons in orbitals • 1st 3 from Schrödinger equation (main energy level, shape, and orientation of orbital) • 4th is spin quantum number • Electrons have specific energy levels (1st, 2nd) • Different energy levels associated w/different orbits • Those nearer nucleus have lower energy than those farther away • Electrons cannot exist between energy levels Quantum numbers

  28. Place a ball at the top of the stairs and roll it gently toward the flight of stairs. • Observe the motion and intermittent resting points of the ball as it moves down the stairs. • What is its final resting place on one step analogous to? • Toss a small ball toward the top of the stairs with as little spin as possible. • Where does it come to rest? What happens if you throw it harder (use more energy)? • What is the amount of energy you use analagous to? Quantization of energy

  29. Main energy level occupied by electron/ size of orbital • As n becomes larger, atom becomes larger and electron is further away from nucleus Principal Quantum Number, n

  30. Cartesian coordinate system(x, y, and z axes) as frame of reference; nucleus located at origin Boundary surface diagrams: volume of space that encloses 90% probability of finding electron within orbital’s boundary surfaces

  31. Shape of cloud • Divides shells into subshells (sublevels) (l) in each principal energy level (l= n-1) Azimuthal Quantum Number, l(angular momentum) • n = 1, 1 sublevel (s) • n = 2, 2 sublevels (p) • n = 3, 3 sublevels (d) • n = 4, 4 sublevels (f)

  32. Divides subshell into orbitals which hold electrons • Specifies 3-D orientation of each orbital around nucleus Magnetic Quantum Number, ml(effect of different orientations of orbitals 1st observed in presence of magnetic field)

  33. Each orbital has specific# sublevels s has 1 sublevel p has 3 sublevels (px, py, pz) d has 5 sublevels (dxy, dyz, dxz, dx2-y2, dz2) f has 7 sublevels • http://college.hmco.com/chemistry/shared/media/animations/1sorbital.html • http://college.hmco.com/chemistry/shared/media/animations/2pxorbital.html • http://college.hmco.com/chemistry/shared/media/animations/2pyorbital.html • http://college.hmco.com/chemistry/shared/media/animations/2pzorbital.html • http://college.hmco.com/chemistry/shared/media/animations/3dxy_orbital.html • http://college.hmco.com/chemistry/shared/media/animations/3dxz_orbital.html • http://college.hmco.com/chemistry/shared/media/animations/3dz2orbital.html

  34. Specifies orientation of spin axis of electron • Creates magnetic field because it spins, oriented in one of two directions • Pairs (diamagnetic) not attracted to magnets • Unpaired (paramagnetic) weakly attracted to magnets Magnetic Quantum Number, ms(spin quantum number)

  35. Each sublevel can contain maximum of two electrons • s has lowest energy (max 2 electrons) • p (max 6) • d (max 10) • f has highest energy (max 14) • Must have opposite spins • 1st electron to fill orbital has a ↑/+ spin • 2nd electron to fill the orbital has a ↓/-spin • You can use /, N/S, +/-

  36. Read 5.2, pp. 127-134 • Q pg. 134, #13, 15, 16 • Q pg. 146, #42, 45, 49, 52, 56, Homework:

  37. All orbitals related to energy sublevel are of equal energy (All three 2p orbitals are of equal energy) • Sublevels w/in principal energy level have diff. energies (Three 2p orbitals are of higher energy than 2s orbital) • In order of increasing energy, sequence of energy sublevels within principal energy level is s, p, d, and f • Orbitals within one principal energy level can overlap orbitals related to energy sublevels within another principal level (Orbital related to atom’s 4s sublevel has lower energy than five orbitals related to 3d sublevel) Aufbau (“building up” in German) principle: each electron occupies lowest energy orbital available

  38. No more than 2 electrons, each with opposing spin (↾⇂), can be located in energy level • No two electrons can have the same set of quantum numbers • If 1 energy level is available, then 2 electrons can be accommodated Pauli exclusion principle-atomic orbital has at most 2 electrons

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