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Chemistry ELECTRONIC STRUCTURE THE nature of light & The Bohr Model

Chemistry ELECTRONIC STRUCTURE THE nature of light & The Bohr Model. DO NOW: TAKE out homework to check Calculators REQUIRED Periodic tables not required. Waves. Light is a form of radiation . Light behaves as a wave , which have certain properties.

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Chemistry ELECTRONIC STRUCTURE THE nature of light & The Bohr Model

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  1. ChemistryELECTRONIC STRUCTURETHE nature of light & The Bohr Model DO NOW: TAKE out homework to check Calculators REQUIRED Periodic tables not required

  2. Waves • Light is a form of radiation. Light behaves as a wave, which have certain properties. • The distance between corresponding points on adjacent waves is the wavelength (λ). Wavelength determines the colorof the light. • The height of the wave is the amplitude. This determines the brightness of the light.

  3. Waves • The number of waves passing a given point per unit of time is the frequency (). • Wavelength and frequency are inversely proportional. • As one increases, the other decreases. • Wavelength: meters • Frequency: cycles per second (s-1) • Also called Hertz (Hz)

  4. Waves • The speed of any wave is given by the wavelength times the frequency of the wave: • Example: what is the speed of a wave that has a wavelength of1.0 x 10-3 m and a frequency of 2.24 x 105 Hz? • The speed of light, c, is 3.0 x 108 m/s.

  5. Electromagnetic Spectrum

  6. Electromagnetic Spectrum • Visible light – 400-750 nm wavelength • ROYGBIV • Ultraviolet: < 400 nm • Shorter wavelength • UV > X-ray > Gamma ray (shortest) • Higher frequency • Higher energy • Infrared: > 750 nm • Longer wavelength • IR < Microwaves < Radio (longest) • Lower frequency • Lower energy

  7. Practice • Which wave has the higher frequency? • If one wave represents visible light and the other represents infrared radiation, which wave is which?

  8. Electromagnetic Radiation • All electromagnetic radiation travels at the same velocity: the speed of light (c). • c = 3.00 x 108m/s • For light, we use a wave equation that related this speed, wavelength, and frequency:

  9. Practice • 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?

  10. IS LIGHT A PARTICLE OR WAVE? • Until the 1900s, light was treated as a wave only. • It has all the properties of a wave – frequency, wavelength, and energy. • The photoelectric effect seemed to be at odds with this wave theory. • Albert Einstein, in 1905, demonstrated that it is not the brightness of light that controls its energy, but its color. • This is at odds with the wave theory – if you treat light as only a wave, higher amplitudes, or brighter light,should mean more energy, if light were only a wave. • This discovery earned Einstein the Nobel Prize in Physics in 1921.

  11. PHOTOELECTriC EFFECT

  12. Photoelectric Effect • The only way to rationalize that the color matters more than intensity was to think of light as a “packet,” or particle of light called a photon. • This is called the wave-particle duality of light, and it is a centerpiece of modern atomic theory. It will help us understand later observations about electrons in the atom, since color indicates energy.

  13. PHOTOELECTRIC EFFECT • Einstein used this conclusion to make a statement about the energy of light. • He concluded that energy is proportional to frequency. • h = Planck’s constant = 6.626 x 10-34 J-s • Using the speed of light equation… • h = 6.626 x 10-34J-s c = 3.00 x 108 m/s

  14. Practice • Calculate the energy of one photon of yellow light with a wavelength of 589 nm. h = 6.626 x 10-34 J-s c = 3.00 x 108m/s

  15. Limitations of Rutherford’s Model • When electricity or heat was applied to elements, strange color patterns, or spectra, emerged. • Heating iron horseshoes led to the black iron becoming red, then yellow, then white. • Rutherford’s model had no way to explain this at the atomic level.

  16. Line Spectra • When we excite the electrons in hydrogen gas and view the emitted light through a spectroscope/prism, the following pattern emerges:

  17. Line Spectra • Instead of a full spectrum of light, only four distinct lines are shown! • This means that electrons in the hydrogen atom have only certain places that they can be, and they are not free to move throughout the hydrogen atom, because otherwise we would see the full spectrum.

  18. The Bohr Model • Danish physicist Niels Bohr proposed a model of the atom that fit the line spectra. • Bohr postulated that electrons could only be in certain energy levels, and that when an electron fell from a higher to a lower energy level, the difference in energy came out as light.

  19. Energy of Emitted Photons • Based on what happens to the light, we can tell the direction of the electron. • If light is absorbed, electrons are excited from a lower to higher energy level. • If light is emitted, electrons relax from a higher to lower energy level. • The energy of the light tells us the difference in energy between the energy levels. • If ΔE is positive, light is absorbed. • If ΔE is negative, light is emitted.

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