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Quantum theory (2)

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Quantum theory (2)

  1. 1. Electrons in Atoms Chapter 4
  2. 2. Rutherford Model
  3. 3. A New Atomic Model Rutherford The ______________ model of the atom was a great improvement, but it was incomplete. It did electrons not explain why negative __________ were not drawn into the positive __________. protons In the early 20th century, a new model was absorption evolving based on the _____________ and ______________ of light by matter. emission The studies revealed a very intimate relationship electrons light between _____________ and ______________.
  4. 4. • Chemists found Rutherford’s nuclear model lacking because it did not begin to account for the differences in __________ chemical behavior among the various __________. elements• In the early 1900s, scientists began to unravel the puzzle of chemical behavior.• They had observed that certain elements visible light emitted ________ _______ when heated in a flame.
  5. 5. Properties of Light Before 1900, light was thought to behave solely as a ________. wave Electromagnetic Radiation ______________ ___________ is a form of energy that exhibits wave-like behavior as it travels through space – ALL FORMS travel at 3.0 x 108 m/s the speed of light: ______________ (c), including: ________________________ Light (visible spectrum), ________________________________ Microwaves, x-rays, ultraviolet, infrared, and gamma. ________________________________.
  6. 6. Electromagnetic Spectrum
  7. 7. Visible Light The visible spectrum includes the colors: ______, _________, __________, RED ORANGE YELLOW ________, ________, _________, and GREEN BLUE INDIGO __________. (ROY G BIV) VIOLET
  8. 8.  The significant feature about waves is repetitive their ___________ nature.
  9. 9. Characteristics of waves Wavelength _______________ (λ) is the distance between corresponding points on adjacent waves. It is measured in _____________ nanometers (nm).
  10. 10.  Frequency _______________ is the number of waves that pass a given point in a specific time, usually “per second”. It is measured Hertz in ________ (Hz).
  11. 11.  The relationship between wavelength and frequency is ____________, and inverse demonstrated in the following equation: c = λνWhere: c = 3.0 x 108 m/s λ = nm ν = Hz (waves/second)
  12. 12. Inverse Relationship
  13. 13. The Photoelectric Effect and the Particle Theory of Light The photoelectric effect refers to the emission of ________ from a metal when electrons light shines on that metal. The mystery is that this would only happen if the light’s frequency was above a certain ________ energy ________, regardless of how long the light level shone on it. According to the __________ Wave Theory, light should have been able to knock electrons loose no matter what frequency.
  14. 14. • In 1900, the Light photons German Electrons physicist ejected from _____________ Max Planck the surface (1858–1947) began searching for an explanation as he studied the light emitted from heated objects. Sodium metal
  15. 15.  Max Planck suggested that electromagnetic energy was NOT __________ like a wave, but was in small, continuous specific ________. bursts A __________ is the minimum amount of energy QUANTUM that can be lost or gained by an electron. He demonstrated mathematically that the energy of a QUANTUM is related to the frequency of the emitted radiation by the equation where E is energy, h is Planck’s constant, and v is frequency. h = 6.626 x 10–34 J · s (where J is the symbol for the joule)
  16. 16.  Albert Einstein furthered the new theory suggesting that light had a _____ wave-particle dual ________ _____ nature He used the term ______ as a photon particle of electromagnetic radiation having zero mass and carrying a quantum of energy
  17. 17. Dual Wave/Particle Theory
  18. 18. The Hydrogen-Atom Line-Emission Spectrum The lowest energy state of an atom is the ________________. ground state The state in which an atom has a higher potential energy is called its _______________. excited state Neils Bohr (1913) experimented with _________. Hydrogen It was known that element had it its own characteristic _____________________. Line-Emission Spectrum He thought that if hydrogen was excited, it would emit a ___________ emission line spectrum, continuous and that it did not matter what amount of energy was added to them, just as long as it was above the _____________. This did NOT happen. minimum
  19. 19.  Excited hydrogen Pinkish atoms emit a hydrogen ________ glow. pinkish When the visible portion of the emitted light is passed through a _______, it was prism separated into __________ specific wavelengths that 1. Not Continuous spectrum made up the hydrogen’s line- 2. Missing colors!! emission spectrum
  20. 20.  Bohr suggested that hydrogen atoms themselves are _________, and that they quantized exist only in certain definite energy states which are called _______________. energy levels This lead to a new atomic theory:
  21. 21.  The characteristic light spectrum of an atom is produced when electrons light emit _______, fall from a higher energy level _________________ to a ________________ lower energy level The energy levels were ladder compared to a _______, that electrons do not emit energy between the rungs, that they emit or absorb energy at specific energy levels ___________________.
  22. 22. Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state.
  23. 23.  It was determined that light was not given off by atoms during _________ of photons absorption of energy (when electrons go from a lower energy level to a higher energy level) but, rather when electrons fell from higher energy level to a lower energy level; this emission of photons was called ________________ (light).
  24. 24.  Bohr’s error was that he thought electrons traveled in ______ with a orbits specific radius around the nucleus. This would later be disproved. As hydrogen only has ____ electron, one his model also failed to demonstrate any atom that had more than ____ electron. one

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