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BELLWORK 10/18/2016

Explore nuclear decay problems and learn about the Bohr Model of the atom, the dual-wave nature of light, and the different energy levels in atoms. Understand the relationship between wavelength, frequency, and energy in the electromagnetic spectrum.

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BELLWORK 10/18/2016

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  1. BELLWORK 10/18/2016 • SOLVE THE FOLLOWING NUCLEAR DECAY PROBLEMS: • 1. 231 Pa 4 He + __________ 912 • 2. 24 Na 0 e  ___________ 11 -1

  2. ANSWERS • 1. 227Ac 89 • 24 Mg 12

  3. -The Bohr Model – Arrangement of Electrons in Atoms Chapter 4.1

  4. Rutherford’s Shortcoming Rutherford’s model of the atom failed to explain what prevents the electrons (that have a negative charge) from being pulled toward the nucleus of the atom (that has a positive charge)?

  5. In the early 20th century, a new atomic model evolved as a result of investigations into the properties of light.

  6. Light Behaves as a Wave Before 1900, scientists thought light behaved solely as a wave. Light was believed to travel through space and matter in repeated patterns, just like waves. Let’s quickly review important characteristics of waves

  7. Wavelength vs. Frequency • Wavelength (λ) – the distance between corresponding points on adjacent waves. Measured in m or nm) • Frequency (v)- The number of waves that pass a given point in one second. Measured in Hz • What happens to the frequency as the wavelength gets shorter? It increases • What happens to the wavelength as the frequency increases? They get shorter • Wavelength and frequency have an inverse relationship

  8. Visible Light is one type of Electromagnetic Radiation (ER). • Electromagnetic radiation is a form of energy that exhibits wave-like behavior • Other forms of ER include X-rays, ultraviolet and infrared light, microwaves, and radio waves. • Together, all these forms of ER form the electromagnetic spectrum.

  9. Electromagnetic Spectrum • What are the colors of the visible light in increasing frequency ?ROY G. BIV • All forms of electromagnetic radiation travel at a speed(c) of 3.00 x 108m/s in a vacuum. • c = λν • What do you notice about the frequency and energy of waves? Waves traveling with higher frequencies carry more energy. • The relationship between energy and frequency is a direct relationship.

  10. Determine the frequency of light whose wavelength is 4.257 x 10-7 cm c = λν c = 3.00 x 108 m/s; λ = 4.257 x 10-7cm *** You need to convert λ to meters by dividing by 100. λ = 4.257 x 10-9m 3.00 x 108 m/s = 4.257 x 10-9 m (v) 3.00 x 108 m/s = 4.257 x 10-9 m (v) 4.257 x 108 m/s c = 7.05 x 1016 Hz ( 3 Sig Figs)

  11. Dual-Wave Nature of Light • Double-slit experiment proved that light traveled as a wave, showing interference.

  12. Photoelectric Effect • The photoelectric effect proved that light exhibited particle like behavior.

  13. Photoelectric Effect • Light is composed of particles called photons that carry energy at certain wavelengths. • If the frequency of light is high enough, electrons can be removed from an atom.

  14. Bohr’s Model of the Atom • Explained why the electrons don’t’ collapse into the nucleus • Explained the relationship between the movement of an atom’s electrons and the production of light.

  15. The further away from the nucleus an electron orbit is, the more energy it has The orbits are a fixed distance from the nucleus Every orbit has a set amount of energy. Photons of light at different wavelengths are released as the electron falls back down to its ground state. An electron has to absorb the exact amount of energy to get from one energy level to another Electrons cannot exist in between orbits

  16. Ground State vs. Excited State Ground State – The lowest possible energy level. Excited State- The energy level the electron is located on once it jumps.

  17. What causes an electron to become “Excited?” • An electron becomes excited when it receives particles of energy (photons) transferred through heat, light, or electricity.

  18. What happens when an atom gets hit with photons? It absorbs energy from the photon (Absorption) and jumps to a higher energy level. An excited electron is unstable, so it will eventually fall back down. When it does, it emits radiation (light) at a certain wavelength.

  19. The energy levels are like the rungs of a ladder. The electrons can only be on one energy level or another, but never in between.

  20. Hydrogen Emission Spectra

  21. PRACTICE An electron absorbs a photon and jumps to the 3rd energy level, as it becomes unstable and falls back down to its ground state, it releases a photon with a wavelength of 700 nm. What color of light will be produced?

  22. Energy State Diagram for Hydrogen • Characterize each of the following as absorption or emission: • E2 to E1 (emission) • E1 to E3(absorption) • E6 to E3 (emission) • Which energy level change above emits the highest energy? • E1 to E3; Lyman series which represents ultraviolet light. We now that more energy is required for ultraviolet as compared to visible (Bahmer Series), or infrared (Paschen series) . • Which energy level change emits the lowest energy? E6 to E3

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