Wave-Particle Duality 1 : The Beginnings of Quantum Mechanics
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Transcript of Wave-Particle Duality 1 : The Beginnings of Quantum Mechanics
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Wave-Particle Duality 1:The Beginnings of Quantum Mechanics
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• 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
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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…
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Problem 1: Only high frequency light (high energy) will eject electrons - acting as particle.
Only explained if thought of as particles in a collision
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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
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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
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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.
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Bohr (1913) – proposed that spectral lines are light from excited electrons.• Restricting electrons to fixed orbits (n) of
different quantized energy levels
Energyn = -2.18 x 10-18 J x Z2/n2
His equations correctly predicted the structured spectral lines of Hydrogen…
• Created an equation for energy of an electron at each orbit
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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−EMR
e−
Ground State
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
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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
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3. Levels are discrete (like quanta) – No in-between.
4. Every jump/drop has a specific energy requirement - same transition, same photon.
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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
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We only “see” those excited electrons that require and releasing energy in the visible spectrum
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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
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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