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Light, Energy, & Electrons

Light, Energy, & Electrons. Discrepant Events/Questions. Chapter 6 Part I. EM Spectrum Light as a wave l v=c Light as a particle E=hv Line spectra Rydberg Equation Bohr’s Hydrogen Model Hydrogen Equation Wave equation of matter. Light is a wave. Light is a particle.

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Light, Energy, & Electrons

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  1. Light, Energy, & Electrons

  2. Discrepant Events/Questions

  3. Chapter 6 Part I • EM Spectrum • Light as a wave • lv=c • Light as a particle • E=hv • Line spectra • Rydberg Equation • Bohr’s Hydrogen Model • Hydrogen Equation • Wave equation of matter

  4. Light is a wave Light is a particle Dual nature of light…

  5. Light as a wave…. • Light acts as a wave • Evidence: • Polarization

  6. Light acts as a wave… • More Evidence • Diffraction Grating (Prism)

  7. Light acts as a wave… • More Evidence • Laser • Laser with Colored Lens • Flashlight with 2 colored lenses • 3D

  8. Filtering • Colored flashlight on other colors

  9. Parts of a “wave” • Wavelength • The distance between two adjacent peaks (or troughs)

  10. Parts of a “wave” • Frequency • The number of waves that pass a given point per second • Frequency & Wavelength are related by: vl=c • V= Frequency • C=speed of light (2.998 x 108 m/s) • l= wavelength

  11. SAMPLE EXERCISE 6.2Calculating Frequency from Wavelength The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of 589 nm. What is the frequency of this radiation? Solution Analyze:We are given the wavelength,  of the radiation and asked to calculate its frequency, . Plan: The relationship between the wavelength (which is given) and the frequency (which is the unknown) is given by Equation 6.1. We can solve this equation for  and then use the values of  and c to obtain a numerical answer. (The speed of light, c, is a fundamental constant whose value is given in the text or in the table of fundamental constants on the back inside cover.) Solve:Solving Equation 6.1 for frequency gives  = c/ . When we insert the values for c and , we note that the units of length in these two quantities are different. We can convert the wavelength from nanometers to meters, so the units cancel: Check:The high frequency is reasonable because of the short wavelength. The units are proper because frequency has units of “per second,” or s–1.

  12. PRACTICE EXERCISE (a) A laser used in eye surgery to fuse detached retinas produces radiation with a wavelength of 640.0 nm. Calculate the frequency of this radiation. (b) An FM radio station broadcasts electromagnetic radiation at a frequency of 103.4 MHz (megahertz; MHz = 106 s–1). Calculate the wavelength of this radiation. Answers:(a) 4.688 1014 s–1, (b) 2.901 m

  13. EM Spectrum • Electromagnetic Radiation • All forms of energy that have “wave-like” behavior” • Electromagnetic Spectrum • A full scale of all the forms of EM radiation

  14. Looking at the spectrum… • Why does Tennent do such a good job of blocking some waves? • Why do microwave windows have a grid?

  15. Light as a Particle • Treating light as a wave accounts for a lot of behaviors, but not all • Examples: • Why heated objects act as they do (Their color changes) • Why metals eject electrons when certain lights shine on them (solar cells)

  16. Energy Relates to Frequency • Absorbing & Emitting Energy • Objects can only absorb/emit energy in certain amounts (packets, quantum) • Energy can be determined by: • E = H * frequency of light Energy Constant Light Emitted or Absorbed H = 6.626 x 10-34 Js

  17. Quantized Energy • What if energy in a car was “quantized”? • This would mean your car can only go at certain speeds (10, 20, 30mph). • Why doesn’t this happen? • Photoelectric Effect • Light of certain frequencies can force electrons out of metals • solar cells (Calculators etc.) • Light Intensity does not matter – only frequency • Photon: energy packet

  18. Photoelectric Effect

  19. SAMPLE EXERCISE 6.3continued PRACTICE EXERCISE (a) A laser emits light with a frequency of 4.69 x 1014 s–1. What is the energy of one photon of the radiation from this laser? (b) If the laser emits a pulse of energy containing 5.0 x 1017 photons of this radiation, what is the total energy of that pulse? (c) If the laser emits 1.3  10–2 J of energy during a pulse, how many photons are emitted during the pulse? Answers:(a) 3.11  10–19 J, (b) 0.16 J, (c) 4.2  1016 photons

  20. Flame Tests • What was responsible for the different colors? • What can we narrow it down to?

  21. Low-Pressure High Voltage Gas Tubes • What color do you “see”? • What color is given off? • Are there any other wavelengths given off?

  22. Continuous vs. Line Spectrum Continuous Spectrum Line Spectrum

  23. Continuous vs. Line Spectrum • Continuous: The rainbow of colors containing all wavelengths • Line Spectrum: Spectrum containing radiation of only specific wavelengths

  24. Balmer & Rydberg • Mid-1800’s • Johann Balmer showed how the wavelengths of the 4 visible lines fit a formula • Additional lines found • Ultraviolet & infrared regions • Rydberg Equation • Calculation of the spectral lines of Hydrogen

  25. Bohr’s Model • Bohr wanted to describe the hydrogen line spectrum more fully • “Planetary” model of electrons • 3 Main Points: • Only orbits of certain “radii”, corresponding to certain energies, are allowed for an electron • An electron in a “level” has a certain energy • Energy is emitted or absorbed only when the electron changes from one level to another

  26. Bohr’s Model Summarized Small orbit = low energy state Large orbit = high energy state • Atom has distinct energy levels, starting with n=1 then n=2, n=3… • Ground State: lowest energy level • When excited, it jumps to a higher state (excited state) • When it goes back down, it emits energy (light) • ‘Step ladder’

  27. Bohr Model ft. Rydberg • Rydberg’s equation showed wavelength • Bohr derived energyfrom this • E=hv and lv=c

  28. Figure 4.16 – Prentice Hall Chemistry Bohr’s Line Spectra • Energy of light given off is due to how far the electron is ‘falling’ through levels • Not all of it is visible • Different jumps give different wavelengths • Grouped in “series” • Lyman series: Emits light in the UV region • Balmer series: Emits light in the visible spectrum • Paschen series: Emits light in the IR region

  29. It neither emits nor absorbs energy. • It both emits and absorbs energy simultaneously. • It emits energy. • It absorbs energy.

  30. It neither emits nor absorbs energy. • It both emits and absorbs energy simultaneously. • It emits energy. • It absorbs energy.

  31. Predict which of the following electronic transitions will produce the longest wavelength spectral line. n = 4 to n = 2 n = 5 to n = 2 n = 5 to n = 3 n = 6 to n = 4

  32. Correct Answer: The wavelength increases as frequency decreases. The lowest frequency (longest wavelength) is associated with the lowest energy, and the smallest energy difference here is between n = 6 and n = 4. n = 4 to n = 2 n = 5 to n = 2 n = 5 to n = 3 n = 6 to n = 4

  33. Practice Exercise 6.4 • Indicate whether each of the following electronic transitions emits energy or requires the absorption of energy: (a)n = 3 to n = 1; (b)n = 2 to n = 4 . Answers:(a) emits energy, (b) requires absorption of energy

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