Whether or not a reaction happens in the gas phase does not change the fact that many reactions are calculated that way. It is easier to simulate reaction thermochemistry in dilue gas phase than in condensed aqueous phase. This is particularly the case in earlier decades, where taking into account intermolecular interactions was simply not possible given computation limitations at the time. If you go to the NIST chemistry webbook, you'll see many reactions are/were calculated this way, even if it isn't particularly realistic. That said, there is some potential relevance to having thermochemical values in the gas phase. Photochemical reactions like I
- + I
2 could occur in the gas phase. (I'm not saying that they do, just that they might.) Anyway, important thing is that you know what phase you're in when you're throwing around thermochemical information, because it does make a difference.
I'm not sure I follow - you are forming a new I-I bond, and the bond dissociation energy of I-I is +150 kJ/mol, so wouldn't you get 150 kJ/mol out of the formation of this bond, so that ΔH° = -150 kJ/mol for this reaction?
Yeah, you're right, my apologies, I accidentally had it reversed. It would be easier to discuss this quantitatively if I could find enough actual data for the standard molar entropies and standard heats of formation for the three species involved using a cursory search of the internet, but I couldn't. Either way, though, it's a bad approximation, because I
3 is ionic and the two bonds in I
3- aren't equivalent to the I-I bond in I
2. The paper I quoted earlier gives an overall reaction enthalpy of 136 kJ/mol but I can't access that paper at the moment to check it out and see the details. From just consideration of # of moles of gas consumed/producedm, it is pretty clear that the entropy change will be negative, which may by itself be enough to make the reaction unfavorable.
At the end of the day, it's just speculation on my part that this is why wikipedia calls the reaction endergonic. It's possible that whoever wrote the article made a simple error and the reaction is exergonic in both aqueous and gas phases at room temperature. Easy enough to do (as you saw above). What is definitely clear is that the reaction is exergonic in aqueous phase a room temperature.