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Author Topic: How Does UV Break Down Chlorine (Hypochlorite ions)?  (Read 361 times)

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bryceking33

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How Does UV Break Down Chlorine (Hypochlorite ions)?
« on: May 18, 2017, 05:27:02 PM »

I haven't really got much to say; I'm doing an EEI (Extended Experimental Investigation) on the oxidation of pool chlorine by UV rays. The only trouble is that there is almost nothing online or anywhere else (that I have been able to find) as to how, specifically, the UV photon breaks the bond. This is only necessary because I would personally like a deeper understanding as well as the fact that a scientific report is required.

I understand that the wavelength/frequency of the hypochlorite ion's (OCL-) bond and the UV ray is the same (more or less) and that this is a factor. Other than that, it seems rather random - like the photon just passes through the bond and then the bond is broken. That seems far too simple and reliant on probability.

I would just like to understand how the UV photon actually breaks the hypochlorite ion as much as is possible.

Thanks for taking the time to read and much appreciated if you attempted to help :)
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Enthalpy

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Re: How Does UV Break Down Chlorine (Hypochlorite ions)?
« Reply #1 on: May 19, 2017, 07:08:35 AM »

Welcome, Bryceking33!

My understanding of this process is a bit crude, and other people here have a more refined one, but maybe it can serve as an introduction.

"Reliant on probability" is, in fine, how our world works. No correct microscopic theory is deterministic. But the splitting process can be described more in details, yes.

And I'm not sure hypochlorite gets "oxidized" in the operation. Broken somehow.

Now, the absorption of a photon is, on Earth, almost always done by an electron in a bound state or orbital, where it pertains to a molecule and has a defined binding energy as compared with a free electron. This energy can be written equivalently as a frequency, but beware that energies are relative to some reference, hence electron frequencies too are relative. "Free electron" is such a reference, but even this one isn't absolute.

The photon absorption puts this electron in a state of higher energy. This state can be free, as in a photocathode emitting electrons to vacuum, or it can be bound, for instance in the semiconductor of a camera CCD. In every case, the photon energy (or frequency) must match the energy (or frequency) difference between the electron's old and new states. This difference doesn't depend on the energy reference.

It would seem that the transition can absorb only one very narrow light frequency, but quite often, a bunch of molecules offers many initial states and many final states, hence many possible transitions. A rigid molecule can have states with energies close to an other; molecular vibrations due to the temperature create more possible states for the electrons; and the speed of each molecule shifts the absorbed frequency by Döppler effect. Also, if the absorption is quick enough, it accepts a small mismatch with the photon's frequency. So usually, the process can accept a set of frequencies, or even a continuum.

Most people are satisfied with "a photon of fitting frequency can be absorbed by an electron transition", or they feel uncomfortable thinking farther, but this process itself can be detailed. As orbitals, electrons are immobile, or with the more accurate word and idea, "stationary". That's why they don't radiate in a molecule at rest. But during a transition, when the electron's wave function is a linear combination of stationary wave functions (like the old and new orbitals), the wave function isn't stationary, the electron isn't immobile, it moves and accelerates, so it can emit or absorb light whose frequency matches the electron's movement, that is, the frequency or energy difference between the old and new states.

Photolysis with usual wavelengths involves a transition from one orbital to an excited one: the electron isn't ripped off. And since the electron uses to absorb the photon faster than the atoms get apart, the excited state of the electron pertains to the same initial molecule, only excited. So the photon energy needed for photolysis is not the energy needed to break the bond, it's significantly bigger - beware many authors are wrong. Quantum mechanics computations (can) succeed in telling what electron energies are possible and what photon frequencies can be absorbed.

Once one electron is excited, varied fates can result. The other electrons can still hold the atoms firmly together, as in ethylene, until the electron de-excites or until the excited molecule makes new bonds. Or the excitation energy can suffice to make the atom's proximity unfavourable, and push the atoms away until the bond breaks.

An "anti-bonding" orbital for the excited electron isn't always enough to separate the atoms. For instance, banal O2 molecules have two electrons on anti-bonding orbitals, but the other electrons are bonding and keep the atoms together. Since a bond often counts two electrons or more, the anti-binding state must be more unfavourable than the binding one(s) is favourable.

More bizarre situations exist, for instance in lamps and lasers using excimers and exciplexes, where molecules of Kr2, XeCl... exist for some time only because they're excited. After emitting a photon to reach the fundamental state, they break apart. So "excited" doesn't mean "anti-bonding".
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