This is a short post in support of the “Can CO2 Form a “Blanket?” post.
Let do a few simple thought experiments, and see what would happen if we mixed two gases under different circumstances. Let’s imagine we have a sealed container with a barrier in the center, dividing the container into two equal spaces, upper and lower. The barrier can be removed (without unsealing the container) to create one continuous space.
For our first experiment, lets pump CO2 into the bottom chamber until it displaces any previous gas and likewise fill the top chamber with O2. Let’s assume that both gases are at the same temperature and the same pressure. Now let’s quickly slide the barrier away and see what happens. The instant the barrier is remover, some CO2 molecules from the bottom half of the container would move into the O2 space and vice versa. They would continue until they hit another molecule and bounced off of it. At the interface, a little game of atomic billiards would have started. CO2 molecules, bouncing around randomly would enter the top chamber and, at first, mostly be bounced back by O2 molecules heading down into the bottom space. As time progressed, the space near the initial interface would become progressively more mixed and some CO2 molecules would, simply by the luck of how they were bouncing around, creep closer and closer to the top of the O2 chamber. And, the reverse would be happening with the O2 molecules; as time went on greater numbers of them would be headed towards the bottom of the container.
As the experiment unfolded, the concentration of each of the gases would form a gradient — almost all CO2 near the bottom of the container, almost no CO2 near the top, with the reverse holding for the O2. Finally, the gradients would disappear as the random jumbling of the molecules thoroughly mixed the gases in the chamber.
At the end of the experiment, the ideal gas laws would describe the container quite well. However, at the beginning of the experiment, they would not. That’s because the ideal gas laws describe the equilibrium condition of a system. The ideal gas laws have bothing to say about systems that are not at equilibrium or how fast they will approach equilibrium — which again is the key to understanding the idea of a CO2 blanket.
Now let’s modify our experiment and run it again. We’ll do everything the same except that — as the experiment progresses — we’ll keep pumping CO2 into the bottom chamber (from the bottom) and we’ll release any excess pressure through a valve at the top of the chamber. In this case, when the barrier is removed and the atomic billiards begin, extra CO2 molecules are entering the chamber from the bottom, slowing the progress of O2 molecules, compared to the previous experiment. Over time, since CO2 is being added, and O2 can only be lost, the chamber will eventually contain pure CO2.
Now, let’s modify the experiment one last time and rerun it. In this let’s start with the conditions of the first experiment except for one thing, let’s make the CO2 colder than the O2. If we run this experiment, the results are similar to what we got in the first experiment, only it took longer for the system to reach equilibrium. The game of atomic billiards was slowed because the CO2 molecules were (on average) moving slower than the O2 molecules and it took longer for the whole jumble to get evenly mixed.