We wouldn't need rose-colored glasses.
We would be nearer to understand how did the care bears felt when riding the rainbow, XD.
Apollinaris would look pink and would be a favorite drink in certain non-standard communities. One could maybe do some photochemistry with carbon dioxide and the sky would not be blue,if that (green/blue) absorption of carbon dioxide was strong enough. In everyday life the air we exhale would be colored, chimneys would release beautiful clouds of pink tinted water vapors. If that absorption was really strong, the plants would not use chlorophyll, animals eyes would make use of a different band of wavelengths to work by, and then maybe Apollinaris might look colorless again....RegardsGeorg
There would be more fatal house fires due to the empty fire extinguishers used at parties to make pink foam.
The oceans, which carry a good deal of dissolved carbon dioxide, would be of a darker color and less transparent. Therefore, there would be less life, and less light-dependent life. Therefore, fish would be less likely to have evolved vision.
If held to it's "elemental nature" in expression, it might "sound" a whole lot different then it natural "colourless" self? :)
Dear all,thanks a lot for the great comments! I especially like the idea with the fire extinguisher :-)I didn't think much about the consequences yet, especially not about the colour of carbonated water!The idea came to when when listening on the radio a few days ago a report about plans to cap the fuel consumption of cars, which nowadays is measured in "gram carbon dioxide per kilometre". A cap of 130 gram per kilometre is roughly 70 litre per kilometre, but I thought we do not have a good intuition for this amount, as we cannot see it. With pink, there would be a coloured kind of mist in the streets. And a room full of people without fresh air would slightly change in colour with time... Though, I have no idea how strong this colouring effect would be.Best, Stefan
Fires would be even more fascinating to watch.If CO2 was pink it might not be a greenhouse gas; hey you changed its spectral characteristics, so can I :) We could happily load the atmosphere with CO2 and coal soot.
“What if carbon dioxide was pink?”It would mean that Al Gore could have subtitled his “Inconvenient Truth” as “The Air Apparent”.
Hi Arun, Phil,yeah, it probably wouldn't be a greenhouse gas then. But there would be pink shades over cities instead... Though, I am not sure about the typical opacity of coloured gases.Cheers, Stefan
There would be less public flatulence.
Hi Low Math, Meekly Interacting,I believe you are confusing one greenhouse gas for another, which is methane rather the carbon dioxide. Besides there is already an olfactory give away that is meant as to serve to limit this.Best,Phil
"..., I am not sure about the typical opacity of coloured gases."Hello Stefan,there are rather few coloured gases. Chlorine, Fluorine, Nitrogen dioxide are those I remember. Colour and reactivity/instability are not the same, but closely related nevertheless. If You want to have a material which is coloured and stable, You need some effective mechanism for the deactivation of the upper state. This is possible in rather big molecules of dyes, not so in gaseous (small) molecules. The three gases mentioned above are very reactive and do not "survive" in atmosphere for considerable time. For that reasons coloured carbon dioxide is really fiction. RegardsGeorg
It wouldn't suprise me if it was, it's such a liberal, girly compound. The H2SO4 in acid rain, now, that's a dark blue macho kick-ass molecule.So are bosons and fermions male and female or the other way around?
Hi Georg,thanks for your comment! I had chlorine and fluorine in mind as examples coloured geses, but didn't think of nitrogen dioxide.Colour and reactivity/instability are not the same, but closely related nevertheless. If You want to have a material which is coloured and stable, You need some effective mechanism for the deactivation of the upper state.I don't get this - could you explain that a bit more? I'm not that much in molecular physics... What is the relation between absorption in the optical, say at 2.5 eV, and reactivity? OK, I guess at that energy, this has to be an electronic excitation. But why does this result in high chemical reactivity in small molecules, but not so in large ones?Thanks, Stefan
Actually, intestinal gas is mostly N2, and depending on your meal and the amount of air you typically swallow, the balance can be mostly CO2, and even O2, in varying proportions. The smelly thiols, sulfides, indoles (e.g. the aptly-named skatole) are only trace ingredients, despite their impact on perception. Methane (which is odorless) may be present or absent, depending on one's repertoire of intestinal flora.Anyone who has frequently ridden in crowded subway cars knows that anonymity can be maintained despite the stench of indiscretion. But if there were also the risk of a visible indicator emanating from the perpetrator, my guess is many more would then opt to get out of visual, as well as olfactory, range, before seeking relief.
Hello Stephan,some more explanations regarding color and gases. - Gases are common in the first period,only few in the second.- Electronic exitation of such species is at wavelengths around 200 nm, lower energys need some konjugated structure, this in turn needs more atoms and thus the molecules will form liquids. - Electronic transitions obey the "Franck-Condon principle", maybe You remember those "vertical" transitions between the electronic states. - Same principle rules the radiationless deactivation processes and intersystem crossing (singlett to triplett).- Thus the likelyhood of transitions of all kindsis dependent on the density of vibrational energy levels. If there are many (eg due to some "heavy" atoms in the molecule or due to the size of the molecule, deactivation is likely to happen. - Solvents in solutions or an adjacent dye molecule in solid state are another means to enhance deactivation. In this case broadening of vibrational levels by collision helps. But this mechanisms do not work in a gas. (light atoms, distances between vibrational levels are great, collisions are seldom) All taken together, only fluorine and chlorine are colored, they have weak single bonds only thus they are easily cleaved, chlorine is rather "heavy" in this context. Nitrogen dioxide is very special, it is a "stable radical", living in equilibrium with its dimer, dinitrogen tetroxide, O2N-NO2 (the latter is colorless).RegardsGeorg
Hi Georg,thanks again for the explanations. So, if I understand that correctly, small molecules that are coloured are weakly bound (otherwise, absorption would be further away in the UV, and the gas colourless), which means that they are quite reactive. Larger molecules, when absorbing visible (or UV) light, don't just swap from the bonding to an antibonding state and split, but can distribute the energy over lots of vibrations etc, hence, dye molecules can be stable and are less reactive?I have just seen that Condon in his paper discusses molecular chlorine - I'll see if I can have a closer look sometime...Anyway, thanks for your interesting comment!Best, Stefan
I would assume if CO2 were pink, you have changed the chemistry of the molecule itself (mostly bond distance, which translates to reactivity in my mind).
a What of then....? Something we'd had always taken for granted, and in colour, a much more "recognize value in living," then one we had ever recognized before? The table then comes to mean something much more, more then we had always known?:)Best,
Hi Shawn,if CO2 were pink, you have changed the chemistry of the molecule itself (mostly bond distance, which translates to reactivity in my mind)that's similar to what Georg has commented a few lines above. I still find this relation between absorption in the optical and chemical reacitvity surprising and unexpected. Do you know a simple and intuitive argument to understand this?Best, Stefan
Stefan,Probably a more intuitive case would deal with population statistics, which is exactly what Georg was hinting at. That is, if we assume this scenario was due to a shift in the electronic excitation of the C=O bond, at lower energies we expect more CO2 molecules to be in excited molecular states. This changes the partition function, which we know as an influence on the total equilibrium constant of any given chemical reaction.We could take this scenario into a more interesting regime. Instead of the 'pink' coming from a shift in molecular excitation levels, what if it was a shift in the way our eyes (the cones and rods, iirc) interpretted specific wavelengths? This would shift the boundaries of the visible spectrum itself, maintaining all wavelengths and all chemistries. I find this scenario equally fascinating, since instead of changing the world itself we have only changed our perception of it.
"...what if it was a shift in the way our eyes (the cones and rods, iirc) interpretted specific wavelengths?"Hello Shawn, there are reasons against that. (of course a small shift is possible, think of the bees seeing the near ultraviolet, but not red. (What about Bee in this respect, Stefan?)Bigger shifts are not practical, and evolution is always "practical".If You keep chemistry, You have the fact that below 350 nm the domain of organic photochemistry starts. You have to protect all tissue from such radiation, including eyes and the receptors. On the other hand, constructing dyes sensitive to infrared is not easy. Absolute lower limit are the shortest vibrational transitions at about 3 µm. The dyes used for infrared sensitizing of photographic emulsions are chemically rather instable and not easy to synthesize.(Dont know whether this is by chance) The fact that plants do not make use of the near infrared radiation, gives another hint to the "not uselful" nature of infrared (in a chemical sense).Chlorophyll strongly reflects IR, a rather uncommon effect among pigments. So, as long as chemistry is maintained, vision has to operate somewhere with the wavelengths we know. The vision band of bees is the lower limit, animals living in the shadow zone of the oceans may have a band without blue extending to the near infrared. RegardsGeorg
Hi Shawn, Georg,very interesting comments, again - thank you. Now, I see, inert small molecules such as N2 absorb in the UV, so when this absorption is to shift into the visible spectrum, electronic excitation has to be "more easy", which at the same time means higher chance to break up the molecule, hence higher reactivity. I hope that is, very roughly, the relation.BTW, the "what if" a wider range in the perception of light had been already scheduled before, so I should delete your last comments to maintain suspense ;-)Cheers, Stefan
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