This sample was evacuated in preparation for vaporizing the fluid for mass spectrometry, decreasing the pressure and, consequently, the boiling point temperature. It turned out that this pressure/temperature combo was near the substance's triple point, the particular temperature and pressure at which the three states of matter (solid, liquid, and gas) exist in thermodynamic equilibrium. As the liquid boils, high energy molecules leave the liquid as gas, lowering the temperature of the liquid left behind and causing it to freeze. This process of boiling and freezing continues while the substance remains at this pressure and temperature.

Edit: I was mistaken when I referred to this as the triple point, as we would not be seeing phase changes at the triple point. Instead, we're seeing phase changes near the triple point; at this particular pressure, the temperatures required for freezing and vaporization are very close. At the true triple point, the three states are in perfect equilibrium, which can be tricky to achieve, as you can see below:

Update: reddit user BantamBasher135 adds that we may also be seeing a solvent at play:

Chemist here. You were indeed freezing it at the same time you were vaporizing it. As /u/m1ld pointed out, you found the triple point of this particular sample. The reason the bubbles were freeing first... well, there are a couple options. First I suspect there might have been some solvent left over. When that was evacuated the pure compound froze instantly. Second, the bubbling at the surface allows for easy re-orientation of the molecules, which allows them to form a crystal lattice without any additional energy expenditure. Additionally, the growth of the crystals on top of the liquid provided surface area—and especially sharp points— for the liquid to then boil.

Typically when we evacuate a sample, we gently heat the flask to avoid such things. A hand holding the flask is usually sufficient.

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