Why is there a difference when calculating pressure temperature and volume using the combined gas law, or when using adiabatic relations? As an idiot I am very confused. Why must I use a very similar equation to calculate the final temperature of an ideal gas but resulting in very different answers. Is the combined gas law only useful in non-adiabatic situations? I will have a follow up question to this most likely, as I am trying to determine the work done on a gas being compressed in this situation: Air being pumped (amazingly adiabatically) under water into a diving bell. I need to find the work done to accomplish this. As far as I can tell (keeping in mind that I am an idiot), Compression Work=(5/2)nR(T2-T1) I'm probably going about this all wrong, but I assumed the bell to be at a depth of about 30.5 meters, so roughly 4 atm. absolute pressure. Then I used Charles law and Boyles law (probably my problems started there?), to calculate the final temperature using that final pressure and an initial pressure of 1 atm, and I got crazy wrong answers. But using the adiabatic formula t2=t1(p2/p1)^(.283) gave more reasonable answers. Which is all good and fine, except I dont understand why, and so I dont know if I'm still wrong, or if I was ever even close. From what I can gather, it has something to do with specific heats and degrees of freedom, but I thought ideal gas relations could be calculated very proportionately without even dealing with those. This is just confusing me. I've tried researching this online, but I get lost in the abstract thermodynamics. Anybody help? Am I on the right track, and why?