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 SpectraThreshold energy

3. PHOTOELECTRIC EFFECT Under certain conditions, shining light 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 quantized energy 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.

9. Bohr (1913) – proposed that spectral lines are light from excited electrons. Restricting electrons to fixed orbits (n) of different quantized energy levels Energy n = -2.18 x 10 -18 J x Z 2 /n 2 His equations correctly predicted the structured spectral lines of Hydrogen… Created an equation for energy of an electron at each orbit

10. 1.Electron absorbs a photon of energy and jumps from ground state (its resting state) to a higher unstable energy level (excited state). Free Atom e−e− EMR e−e− Ground State e−e− Excited State Ionization Absorption EMR nucleus > Threshold Energy < Threshold Energy 2. Electron falls back to ground state – releasing the same photon of energy. “unstable” is the KEY - electrons are attracted to the nucleus and can’t stay away for long

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

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

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

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

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

16. 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 SpectraThreshold energy