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Wave-Particle Duality 1 : The Beginnings of Quantum Mechanics

Wave-Particle Duality 1 : The Beginnings of Quantum Mechanics. Define the relationship between quantum and photon . Describe how a produced line spectra relates to the Bohr diagram for a specific element. Additional KEY Terms Absorption Spectra Threshold energy . PHOTOELECTRIC EFFECT.

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Wave-Particle Duality 1 : The Beginnings of Quantum Mechanics

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

  2. Define the relationship between quantum and photon. • 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 Building on Planck’s quantum idea, Einstein tried to explain this phenomenon…

  4. Problem 1: Only high frequency light (high energy) will eject electrons - acting as particle. Only explained if thought of as particles in a collision

  5. Problem 2: Only more intense light (higher amplitude) will eject more electrons - acting as wave. Only explained if thought of as changing the “size” – amplitude of the wave

  6. Einstein (1905) – EMR is a stream of tiny “packets” of quantizedenergy carried in particles called - photons. A photon have no mass but carries a quantum of energy Light is an electromagnetic WAVE, made of PARTICLE-like photons of 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. Bohr (1913) – proposed that spectral lines are light from excited electrons. • Restrictingelectronsto fixed orbits (n) of different quantized energy levels • Created an equation for energy of an electron at each orbit Energyn = -2.18 x 10-18 J x Z2/n2 His equations correctly predicted the structured spectral lines of Hydrogen…

  9. EMR e− Free Atom e− e− Ground State Excited State • Electron absorbs a photon of energy and jumps from ground state (its resting state) to a higherunstable energy level (excited state). • Electron falls back to ground state • – releasingthe same photonof energy. “unstable” is the KEY - electrons are attracted to the nucleus and can’t stay away for long Absorption Ionization EMR nucleus > Threshold Energy < Threshold Energy

  10. ΔE = E higher-energy orbit - E lower-energy orbit = Ephoton emitted = hf The difference in energy requirements between orbits determines the “colour” of photon absorbed/released by the electron

  11. 3. Levels are discrete (like quanta) – No in-between. 4. Every jump/drop has a specific energy requirement - same transition, same photon.

  12. The size of the nucleus will affect electron position around the atom – and the energy requirements Na: 11 p+ 11 e- Cl: 17 e- Each element has a unique line spectrum as each element has a unique atomic configuration 17 p+

  13. We only “see” those excited electrons that require and releasing energy in the visible spectrum

  14. Notice energy absorbed is the same as energy released Absorption spectrum – portion of visible light absorbedby an element – heating up. Emission spectrum – portion of visible light emitted by that element – cooling down.

  15. CAN YOU / HAVE YOU? • 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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