CO2 condensation in Antarctica at -113F?

Below the freezing point of a gas, it can exist as a solid. In the winter, we see the transition from gas to solid as frost. Frost occurs even at extremely low humidities.
There are other things going on as well. The ice may be getting hit by photons or other energetic particles which cause them to convert back to gas. This can be seen when the ice is exposed to sun or wind, even on a very cold day. The higher the humidity, the greater the tendency for water molecules to form ice. This phenomenon is not covered by phase diagrams, which simply show the preferred state at a given ambient temperature and pressure.
At -113F and 1atm, that state would be solid for CO2.

@TonyS… Here another article on clouds on Mars:
http://www.sciencemag.org/cgi/content/full/sci;282/5396/2053?maxtoshow=&HITS=10&hits=10&RESULTFORMAT=&fulltext=clouds+on+mars&searchid=1&FIRSTINDEX=20&resourcetype=HWCIT
Fragment of the abstract:
“Clouds observed over the polar cap are likely composed of CO2 that condensed out of the atmosphere during northern hemisphere winter.“

There is nothing magical about -140F, just as there is nothing magical about Hansen’s Armageddon value of 350 ppm CO2 . The freezing point of CO2 at 1 atmosphere is -109F.
Once again, the phase diagram for CO2 does not consider the partial pressure of CO2. It considers only the atmospheric pressure, which is much higher. The equilibrium between sublimation and deposition is a complex function of wind, sun, chemistry, surface roughness, etc. and can not be inferred from a phase diagram.
The case we are all familiar with is frost forming on a dirty windshield, but not on a clean windshield – with everything else identical.

The phase diagram shows unambiguously that the equilibrium state of CO2 at one atmosphere at 113F is solid. The freezing point of CO2 is -109F at 1 atmosphere.
http://www.chemicalogic.com/download/co2_phase_diagram.pdf
I see Phil has created a new magic number of -130F.
Personally, I think the world was doomed at 300ppm CO2. I’ve got work to do. l8r

The atmospheric pressure of Mars is less than 1% of earth.

So basically everyone is correct, the argument is one of at what temperature it happens? By the way −128.6°F is not that far off -140F, just 8’K or 4.5%?
“At this point, after doing research of my own, including a telephone call to a local expert on compressed gases (since I am no expert in that field) and reading comments, particularly that of George E. Smith which answered the question in my mind of why we could have solid to gas phase transition of CO2 (sublimation) but not the other way around (at least at -113F) I’m ready to call this myth of CO2 “snow” condensing from the free atmosphere at the south pole at -113F as transitioning from “plausible” to “busted”.
Though, it will happen in the free atmosphere at -140F, of that there is no doubt, but since the coldest natural temperature ever recorded on Earth was −89.2°C (−128.6°F) at the Russian Vostok Station in Antarctica July 21, 1983, I don’t think we’ll see this effect in the free atmosphere today, and I sincerely hope I never witness it firsthand ;-)”

“”” George E. Smith (13:52:09) :
The triple point of CO2 is -56.6 deg C at a pressure of 5.11 atmospheres; that’s 3883.6 mm Hg pressure.
Below that temperature, liquid CO2 canot exist. At zero degree C you can liquify CO2 at a pressure of about 15-20 atmospheres, and solidify it at about 1000 atmospheres.
In the case of the Antarctic plateau, the CO2 vapor pressure stays about constant at around 0.447 mm Hg, and that is not enough to stop continued sublimation till the temperature is about -140 deg C.
So any remaining amount of CO2 ice would be very hard to locate at -113 deg C.
REPLY: Well said, and without labeling anyone, insults, or snark being tossed. Phil. could learn a few things from you. BTW I think you meant -113 and -140 deg F – Anthony “””
Well thank you Anthony, and yes that was supposed to be F rather than C.
When we had this discussion before; I was pretty darn sure I was right; particularly since I ran into a chap who had been to the south pole and he claimed he had been walking around on dry ice at the south pole, and it wasn’t anywhere near as cold as Vostok. And I thought I understood it; because I sure know what happens with the vapor pressure of Arsenic over hot gallium Arsenide in a diffusion sealed ampoule.
But Phil; and I believe it was Ric Werme kept insisting I was wrong; and I couldn’t see it, till I had the phase diagram that Phil posted; and I started thinking about the fact that it is a dynamic equilibrium if both species coexist like water vapor over a dish of water under glass, and then I could see that no reasonable amount of dry ice could survive even at -78.5 C because it would all evaporate without raising the population of CO2 molecules in the vapor phase to where the rate of returning molecules to the solid surface increased enough to balance the loss rate.
And once I “got it”, I got mightily embarrassed that I hadn’t seen it on my own without Phil’s and Ric’s prodding. Now I can’t figure out why I didn’t figure it out by myself.
But on Mars since most of the atmosphere is CO2 (90% I believe), then most of the atmospheric presssure; even though it is low is CO2 vapor partial pressure, and it is enough to equilibrate with solid CO2 at the -120 to -125 deg C polar temperatures on Mars. So I have no doubt that CO2 ice does exist on Mars, and pretty much doesn’t on earth; I suppose one could calculate how many molecules of CO2 clumped together, would be small enough dry ice sample to remain in equilibrum with 385 ppm of CO2 in the atmosphere; and at Vostok or South pole altitude where the atmospheric pressure is diminished.
But it was Phil’s nudging that convinced me that it was the partial pressure of the species that had to be in equilibrium with the solid because those are the two parties exchanging molecules. It is true that the total pressure affects the boiling point; because the boiling point is simply that temperature at which bubbles of the vapor species can appear in the bulk of the liquid, because the vapor pressure at that temperature is equal to the ambient pressure. (remembering in the case of at least water with a sizeable surface tension, that there must be an internal excess pressure of 2T/r in the bubble. You can prove this by using the principle of virtual work. Imagine a miniscule increase in radius doing a miniscule amount of work equal to the excess pressure times the surface area times that miniscule radius increase. that has to be equal to the work done against surface tension, which is simply the surface tension (T) timews the increase in surface area due to the radius increase. Put those two equal, and you get p = 2T/r
In the case of a soap bubble you have two surfaces so the internal excess pressure is 4T/r.
George