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Wave-Particle Duality : The Beginnings of Quantum Mechanics

Wave-Particle Duality : The Beginnings of Quantum Mechanics. Explain the basics of wave-particle duality . Define the relationship between quantum, photon and electron. Describe how a produced line spectra relates to the Bohr diagram for a specific element. Additional KEY Terms

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Wave-Particle Duality : The Beginnings of Quantum Mechanics

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

  2. Explain the basics of wave-particle duality. • Define the relationship between quantum, photon and electron. • Describe how a produced line spectra relates to the Bohr diagram for a specific element. Additional KEY Terms Absorption Spectra Threshold energy

  3. PHOTOELECTRIC EFFECT Under certain conditions, shininglight on a metal surface will eject electrons. Electrons given enough energy (threshold energy) can escape the attraction of the nucleus. *Light is acting like a “particle” in this experiment – collision.

  4. Only high frequency light (> 1.14 x 10 15 Hz) will eject electrons - acting as particle. Can only explain it if you think of it using photons in a collision.

  5. Only more intense light (higher amplitude) will eject more electrons - acting as wave. Can only explain it if you think of it as changing the size of the wave.

  6. 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.

  7. Compton (1922) – first experiment to show particle and wave properties of EMR simultaneously. Incoming x-rays lost energy and scattered in a way that can be explained with physics of collisions.

  8. Quantum Mechanical Model of the Atom

  9. Bohr (1922) – restrictingelectronsto fixed orbits (n) with different quantized energy levels. Created a math equation for energy of each orbit. Equations correctlypredicted the line colours of hydrogen spectra. Energyn = -(2.18 x 10-18 J)/n2

  10. 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. EMR e− Free Atom e− e− Ground State Excited State Absorption Ionization EMR nucleus > Threshold Energy < Threshold Energy

  11. ΔE = E higher-energy orbit - E lower-energy orbit = Ephoton emitted = hf

  12. Levels are discrete like quanta – no in between. • Each jump/drop is associated with a specific frequency photon - same transition, same photon.

  13. The size of nucleus will affect electron position around the atom – and the size of “jump” energy. Na: 11 p+ 11 e- Cl: 17 e- 17 p+

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

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

  16. CAN YOU / HAVE YOU? • Explain the basics of wave-particle duality. • Define the relationship between quantum, photon and electron. • Describe how a produced line spectra relates to the Bohr diagram for a specific element. Additional KEY Terms Absorption Spectra Threshold energy

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