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