6
$\begingroup$

Without appealing to reactions that require photons / light, how do chemists achieve a very high concentration of continually produced, though quite short lived, singlet oxygen in aqueous solution (i.e. $\ce{{}^1O_2}$ where dioxygen is typically in the triplet $\ce{{}^3O_2}$ state)? I suppose I'll also ask the same question if we have the freedom to select an organic solvent of choice for the buffer? By "high concentration" I mean at concentration from $\approx 10~\mathrm{\mu{}M} - 100(+)~\mathrm{\mu{}M}$.

My guess would be that should be a way to do this starting with a solution containing some amount of $\ce{H_2O_2}$, but I'm not quite sure what a reasonable process might be.

Here's the wikipedia link for singlet oxygen ($\ce{{}^3O_2}$)

This reaction should happen at or around room temperature.

The reason I'm specifying that I'd like to do this in the dark, and to have the singlet oxygen continually produced at room temperature (given its short lifetime, which is $\leq 100~\mathrm{\mu{}s}$ under any conditions I've heard of), is because I'm curious if I could easily set up some buffer that lets me dose a sample with singlet oxygen on a lab bench with a pipette. Given the short lifetime of singlet oxygen, having to use a continuous radiation source presents difficulties that, while surmountable, lead me to ask this question.

Provided the above, I'd like to avoid reactions that vent gases like $\ce{Cl_2}$ and require fume hoods in the sense of "... well, if you value your life ..." vs. "... OSHA says so ...".

$\endgroup$
3
$\begingroup$

Potassium ferricyanide (and other ion complexes, like hemoglobin) will decompose hydrogen peroxide into water and (presumably) singlet oxygen. This process is necessary for the luminol reaction.

$$\ce{2H2O2 \overset{Fe}{->} 2H2O + ^1O2}$$

$\endgroup$
  • $\begingroup$ Thanks, do you know what temperature is required for this reaction? Also, how do we know that singlet oxygen is being produced? $\endgroup$ – Bob Jan 28 '14 at 13:26
2
$\begingroup$

I'd always go by the old photochemical Schenck method, but since you insist:

The thermal decomposition of

  • Phosphonite ozonides or
  • 9,10-diphenylanthracene endoperoxide

yield $\ce{^1O2}$.

Technically, $\ce{^1O2}$, may also be generated at low oxygen pressure (i a vacuum line) by microwave radiation. (Source)

As far as aqueous solutions are concerned, one should keep in mind that the lifetime of singlet oxygen ($^1\Delta$) is around 2 µs.

For detection, one typically relies on the luminescence at $\lambda$ = 1270 nm.

$\endgroup$
  • $\begingroup$ I think this is getting close to what I'm looking for... what temperatures are required? I'm looking for something that happens at or around room temperature (e.g. $25C$)? $\endgroup$ – Bob Jan 28 '14 at 13:27
  • $\begingroup$ I would have to look that up. $\endgroup$ – Klaus-Dieter Warzecha Jan 28 '14 at 13:36
2
$\begingroup$

Unraveling a stoichiometric organic adduct in situ is the way to go. However,

$\ce{H2O2 + NaOCl → O2(a1Δg) + NaCl + H2O}$

It is one thing to mix streams of supermarket $\ce{H2O2}$ and laundry bleach to get a dimer nice red glow in a darkened room. It is quite another thing to mix higher concentrations of each. Concentrated $\ce{H2O2}$ is exceptionally dangerous on skin! It diffuses in until it hits living tissue with peroxidase. It then pops into subcutaneous gas bubbles attended with remarkable pain. Concentrated hypochlorite is evil, including conproportionation with chloride to chlorine gas.

Full safety equipment: face shield, gloves, apron - and remember to shield your neck. In a running fume hood with the door down. Get literature references before you even think about doing it. A hefty dual barrel syringe pump feeding a static mixer might barely pull it off. Keep the reacting volume small as you rocket through the juices. You might take singlet oxygen somewhere outside the reaction zone.

100 mph = 4.47 mm/100 microsecond half-life

$\endgroup$
  • $\begingroup$ This sounds awesome. Are you aware of any literature references for the kinetics of singlet oxygen production over time (e.g. as monitored by absorbance @ 1268 nm) as a function of $\ce{H2O2}$ and $\ce{NaOCl}$? $\endgroup$ – Bob Jan 28 '14 at 22:57
  • $\begingroup$ I'm slightly worried though, since it seems like this would also yield chlorine gas? Consider that $\ce{H2O2 + 2NaOCl -> 2NaOH + O2 + Cl2}$? $\endgroup$ – Bob Jan 28 '14 at 23:00
  • $\begingroup$ Hypochlorite is already ~pH 13, then Le Chatelier. Any generated chlorine will disproportionate in strong alkali. Do it with ventilation. $\endgroup$ – Uncle Al Jan 30 '14 at 22:44
  • $\begingroup$ In my experience, concentrated hydrogen peroxide is painful on skin, but leave no chemical burns, and whitened skin returns to normal quite soon. On the other hand, 30% hydrogen peroxide may decompose with high heat output if a proper catalyst (manganese or iron salt) is poured into it, and highly concentrated hydrogen peroxide (>50%) may be explosive. $\endgroup$ – permeakra Jul 30 '14 at 5:47
0
$\begingroup$

There is an excellent article on singlet oxygen on Wikipedia. Singlet oxygen can be generated by the action of light on Rose Bengal dye. The brighter the light, the more generated. It is reported to have a half-life of microseconds in solvents and 72 minutes in air at ambient temperature.

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.