Absorption and Emission of Light
High-Yield Summary
- Bohr model: electrons occupy quantized energy levels — they can't exist between levels, only jump via absorbing or emitting a photon.
- Absorption: electron jumps to a higher level (excited state) by absorbing a photon whose energy exactly matches the gap, E = hf.
- Emission: electron falls to a lower level, releasing a photon with energy equal to the gap — its wavelength sets its color/spectral position.
- Absorption spectroscopy (IR = bond vibrations, UV-Vis = electronic transitions) measures absorbed wavelengths; emission spectroscopy measures released wavelengths (e.g., flame-test spectral lines).
- Fluorescence: absorbs high-energy (e.g., UV) light, re-emits lower-energy visible light — some energy lost as heat, so emitted λ is longer than absorbed λ (Stokes shift).
Absorption vs. Emission
| Absorption | Emission |
|---|---|
| Energy absorbed (in) | Energy released (out) |
| Electron: lower level → higher level | Electron: higher level → lower level |
| Resulting state: excited | Resulting state: ground (or lower excited) state |
Photon Energy at a Transition
E = hf
- E = Energy gap between the two levels (photon energy)
- h = Planck's constant
- f = Frequency of the absorbed or emitted photon
- The photon's energy must exactly match the level gap — no partial absorption/emission.
Key Terms
- IR spectroscopy
- Absorption spectroscopy studying molecular vibrations — specific chemical bonds absorb specific IR wavelengths.
- UV-Vis spectroscopy
- Absorption spectroscopy examining UV/visible light absorption, revealing electronic transitions.
- Emission spectroscopy
- Measures wavelengths of light released as excited electrons fall to lower energy states (e.g., flame tests).
- Fluorescence
- Absorbing high-energy light and re-emitting lower-energy visible light after some energy is lost as heat.
- Stokes shift
- The shift toward longer wavelength between the light a fluorescent material absorbs and the light it emits.
Common MCAT Trap
- A photon can only be absorbed if its energy exactly matches an available level gap — don't assume any photon with 'enough' energy gets absorbed.
- Fluorescence emission is always lower energy / longer wavelength than the absorbed light, never the same or shorter — energy is lost as heat during relaxation.
Quick Recall
Why can't an electron sit 'between' two Bohr energy levels?
A material absorbs UV light and emits blue visible light. What is this phenomenon called, and why is the emitted light lower energy?