Quantum Fluorescence.

Understanding the universe’s most popular particle

Image by FLY:D on Unsplash

What is light? Where does it come from? Why does it seem to interact differently with different things? And what about fireflies, or things that glow in the dark, and everything else that seems to defy any intuitive understanding of light?

Welcome to Quantum City

Some orbital solutions to Schrödinger’s wavefunction. The red contours represent regions of high electron probability (take them with a grain of salt since they’re not to scale). In reality, the orbitals shown here should keep getting bigger and further from the nucleus as you go to the right.
A so-called Jabłoński diagram of electrons at different shells and orbitals next to our neighbourhood analogy of electron configurations. Any electrons within the same house (or orbital) have the same energy, while electrons in different houses. Why? That’s up next.

Energy and Luminescence

The main idea here is that neither elastics nor subatomic particles will stay in a place that’s energetically unfavourable to them. You need energy to hold them there.

Similarly, any electron held within the attractive grasp of a nucleus has a default position termed its ground state. This is an orbital that the electron can exist in without any need for external pushing or pulling, and corresponds to a state of relatively low electric potential energy.

Fₑ represents the electrostatic force, while k is a constant, and q1 and q2 represent the two charges involved. It may seem contradictory that the orbital with the bigger radius has a higher PE, but keep in mind that Coulomb’s law solves for force. To even get an electron this far away from its default state (near a proton), you’d first have to continually apply this force to get to the same radius as the smaller orbital — and then some.
E=ℏv relates the frequency of a photon (γ) to the amount of energy it carries. represents the Planck constant, which is pretty big deal in quantum mechanics. When a photon meets an electron and has enough energy to help it bridge a shell (e.g. n1-n2), excitation can occur. Then, the electron re-emits all that energy. When E is lesser than the energy required to bridge a shell, electrons don’t absorb the photon or go anywhere.

Stairwells and Fluorescence

It’s an odd-looking staircase in the background, but a staircase nonetheless…
In elastic scattering, no energy loss occurs between when the electron is excited and when it emits a photon. As a result, the photon that is released has all the energy that was in the original photon.
Because of non-radiative energy loss before an electron re-emits a photon, the excited electron will always emit light that is less energetic than the light it was illuminated with.

Applying Fluorescence Experimentally

This is the closest thing we have to an atomic barcode: the next time we observe a reflected photon with an identical energy and wavelength, we can trace it right back to the particular molecule that released it.

With filters, we can selectively hone in on light fluoresced by NAD+ at 341nm and 343nm, as shown by γ(fluorescent). All other wavelengths, which must have come from other molecules, can be discarded.
A relatively low concentration of NAD+ in our sample (illustrated on the left) would naturally make it less likely for our incident photons to collide with them. This would reduce the quantum yield, ϕ, which would be defined as the ratio of 341 and 343nm photons picked up by the sensor. With a higher [NAD+], more collisions would occur and more of the light that’s released would be able to cross the filter. This would eventually make ϕ approach 1. Note that wavelength is denoted by the lowercase Greek letter lambda (λ)

Can I have my Ph.D. Now?

But fundamentally, all this tech is the answer to one of our most wild and ancient pursuits as humans: How do you see without actually seeing?

TLDR (for real this time)



I write about things every week(ish).

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