1 / 34

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.

ratana
Télécharger la présentation

Light, Energy, & Electrons

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related