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

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Presentation transcript:

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 SpectraThreshold energy

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…

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

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

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

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.

Bohr (1913) – proposed that spectral lines are light from excited electrons. Restricting electrons to fixed orbits (n) of different quantized energy levels Energy n = x 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

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

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

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

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

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

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

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

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