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E = hf

E = hf. E – energy of a quantum (Joules) h – Plank’s constant (6.626 x 10 -34 J  s) f – frequency of absorbed or emitted EMR. Wave-Particle Duality : The Beginnings of Quantum Mechanics. Describe the photoelectric effect. Understand the basics of wave-particle duality.

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E = hf

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  1. E = hf E – energy of a quantum (Joules) h – Plank’s constant (6.626 x 10-34 J s) f – frequency of absorbed or emitted EMR

  2. Wave-Particle Duality: The Beginnings of Quantum Mechanics

  3. Describe the photoelectric effect. • Understand the basics of wave-particle duality. • Be able to explain how the Bohr model fits with knowledge of line spectra. • Understand the difference between quantum, photon and electron.

  4. PHOTOELECTRIC EFFECT Shining light on a metal surface will immediately eject electrons. Electrons given enough energy (ionization) can escape the attraction of the nucleus. *Light is acting like a “particle” in this experiment – collision.

  5. Einstein (1905) - electromagnetic radiation is a stream of tiny bundles of energy called photons. Photons have no mass but carry a quantum of energy. One photon can remove one electron. Light is an electromagnetic wave, yet it contains particle-likephotonsof energy.

  6. Only high frequency light (> 1.14 x 10 15 Hz) will eject electrons - acting as particle. The higher the frequency (more energy), the faster the electrons move.

  7. Only more intense light (higher amplitude) will eject more electrons - acting as wave.

  8. Compton (1922) – first experiment to show particle and wave properties of EMR simultaneously. Incoming x-rays lost energy (lower frequency) and scattered after the collision with an electron.

  9. Quantum Mechanical Model of the Atom

  10. Recent Developments of Atomic Structure Thomson(1897) - "plum pudding” model. Large positive charge with very small electrons stuck randomly in.

  11. Rutherford(1911) - “Gold foil ” experiment. Helium nuclei – alpha (α) particles - fired at thin gold foil reflected strongly. Discovered the nucleus– electrons just fly around.

  12. Bohr (1922) – explains emission (line) spectrum of elements by restricting electrons to fixed orbits with different quantized energy levels.

  13. Electron absorbs radiation and jumps from • ground state (its resting state) to a higherunstable energy level (excited state). • Electron soon loses energy and drops back down to a lower energy level – emitting the absorbed EMR..

  14. Energy levels are discrete – no in between. • Each jump/drop is associated with a specific frequency photon - same transition, same photon.

  15. *Each element has a unique line spectrum as each element has a unique atomic configuration.

  16. Emission spectrum – portion of visible light emitted by that element – cooling down. Absorption spectrum – portion of visible light absorbed by an element – heating up.

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